Introduction to Chemistry

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 Course Objectives by Chapter .................................................................................................. 5
Chapter 1: Introduction to Chemistry & the Nature of Science............................................... 8
1.1: The Process of Science ................................................................................................. 8
1.2: Hypothesis, Law, & Theory ......................................................................................... 14
Table of Contents
1.3: Graphing ...................................................................................................................... 18
Chapter 2: The Structure of the Atom .................................................................................... 24
2.1: Early Ideas of Atoms ................................................................................................... 24
2.2: Further Understanding of the Atom ............................................................................ 28
2.3: Protons, Neutrons, and Electrons in Atoms ................................................................. 35
2.4: Atomic Mass ................................................................................................................ 41
2.5: The Nature of Light ..................................................................................................... 43
2.6: Electron Arrangement in Atoms .................................................................................. 50
Chapter 3: The Organization of the Elements ......................................................................... 55
3.1: Mendeleev’s Periodic Table ........................................................................................ 55
3.2: Metals, Nonmetals, and Metalloids ............................................................................. 59
3.3: Valence Electrons ........................................................................................................ 61
3.4: Families and Periods of the Periodic Table ................................................................. 62
3.5: Periodic Trends ............................................................................................................ 65
Chapter 4: Describing Compounds ......................................................................................... 71
4.1: Introduction to Compounds ......................................................................................... 71
4.2: Types of Compounds and Their Properties ................................................................. 74
4.3: Names and Charges of Ions ......................................................................................... 78
4.4: Writing Ionic Formulas ................................................................................................ 84
4.5: Naming Ionic Compounds ........................................................................................... 86
4.6: Covalent Compounds & Lewis Structures................................................................... 90
4.7: Molecular Geometry .................................................................................................... 94
4.8: Polarity & Hydrogen Bonding ..................................................................................... 97
Chapter 5: Problem Solving & the Mole .............................................................................. 104
5.1: Measurement Systems ............................................................................................... 104
5.2: Scientific Notation ..................................................................................................... 109
5.3: Math in Chemistry ..................................................................................................... 111
5.4: The Mole .................................................................................................................... 114
Chapter 6: Mixtures & Their Properties ............................................................................... 118
6.1: Solutions, Colloids, and Suspensions ........................................................................ 118
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6.2: Solution Formation .................................................................................................... 121

6.3: Concentration ............................................................................................................. 124
6.4: Colligative Properties ................................................................................................ 128
Chapter 7: Describing Chemical Reactions .......................................................................... 134
7.1: Chemical & Physical Change .................................................................................... 134
7.2: Reaction Rate ............................................................................................................. 137
7.3: Chemical Reactions and Equations ............................................................................ 145
7.4: Balancing Chemical Equations ................................................................................. 148
7.5: Types of Reactions..................................................................................................... 153
7.6: Stoichiometry ............................................................................................................. 159
7.7: Reversible reaction & Equilibrium ............................................................................ 165
7.8: Equilibrium Constant ................................................................................................. 168
7.9: The Effects of Applying Stress to Reactions at Equilibrium ..................................... 171
Chapter 8: Describing Acids & Bases .................................................................................. 177
8.1: Classifying Acids and Bases ...................................................................................... 177
8.2: pH............................................................................................................................... 180
8.3: Neutralization............................................................................................................. 184
8.4: Titration ..................................................................................................................... 186
Chapter 9: Energy of Chemical Changes .............................................................................. 190
9.1: Energy ........................................................................................................................ 190
9.2: Endothermic and Exothermic Changes...................................................................... 191
9.3: Oxidation – Reduction ............................................................................................... 194
Chapter 10: Nuclear Changes ............................................................................................... 201
10.1: Discovery of Radioactivity ...................................................................................... 201
10.2: Types of Radiation ................................................................................................... 203
10.3: Half-life & Rate of Radioactive Decay .................................................................... 209
10.4: Applications of Nuclear Changes ............................................................................ 213
10.5: Big Bang Theory ...................................................................................................... 219
Unit 3: Gases ......................................................................................................................... 222
11.1: Gases and Kinetic Theory ........................................................................................ 222
11.2: Gas Laws.................................................................................................................. 226
11.3: Ideal Gas Law .......................................................................................................... 231
Answers to Selected Problems .............................................................................................. 234
Glossary ................................................................................................................................ 246
 Course Objectives by Chapter
Unit 1: Introduction to Chemistry and the Nature of Science
Nature of Science Goal—Science is based on observations, data, analysis and conclusions.
1. I can distinguish between observable (qualitative) and numeric (quantitative) data.
2. I can construct and analyze data tables and graphs.
3. I can identify independent, dependant, and controlled variables in an experiment
description, data table or graph.
4. I can write a laboratory summary in a Claim-Evidence Format

Unit 2: The Structure of the Atom

Nature of Science Goal—Scientific understanding changes as new data is collected.
1. I can use atomic models to explain why theories may change over time.
2. I can identify the relative size, charge and position of protons, neutrons, and electrons
in the atom.
3. I can find the number of protons, neutrons and electrons in a given isotope of an
element if I am given a nuclear symbol or name of element and mass number.
4. I can describe the difference between atomic mass and mass number.
5. I can describe the relationship between wavelength, frequency, energy and color of
light (photons).
6. I can describe the process through which the electrons give off photons (energy) and
describe the evidence that electrons have specific amounts of energy.
7. I can identify an unknown element using a flame test or by comparison to an emission
spectra.
8. I can write electron configurations for elements in the ground state.

Unit 3: The Organization of the Elements

Nature of Science Goal—Classification systems lead to better scientific understanding.
1. I can describe the advantages of Mendeleev’s Periodic Table over other
organizations.
2. I can compare the properties of metals, nonmetals, and metalloids.
3. I can determine the number of valence electrons for elements in the main block.
4. I can explain the similarities between elements within a group or family.
5. I can identify patterns found on the periodic table such as reactivity, atomic radius,
ionization energy and electronegativity.

Unit 4: Describing Compounds

Nature of Science Goal—Vocabulary in science has specific meanings.
1. I can indicate the type of bond formed between two atoms and give properties of
ionic, covalent, metallic bonds and describe the properties of materials that are
bonded in each of those ways.
2. I can compare the physical and chemical properties of a compound to the elements
that form it.
3. I can predict the charge an atom will acquire when it forms an ion by gaining or
losing electrons using the octet rule.
4. I can write the names and formulas of ionic compounds.

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 5. I can indicate the shape and polarity of simple covalent compounds from a model or
drawing.
6. I can describe how hydrogen bonding in water affects physical, chemical, Nature of Science Goal--Nature is moving toward equilibrium
and biological phenomena.
1. I can describe properties of acids and bases and identify if a solution is acidic
Unit 5: Problem Solving and the Mole or basic.
2. I can calculate the pH of a solution.
Nature of Science Goal— Mathematics is a tool to increase scientific understanding.
3. I can write a neutralization reaction between an acid and base.
1. I can describe the common measurements of the SI system of measurements
4. I can calculate the concentration of an acid or base from data collected in a titration.
2. I can convert between standard notation and scientific notation.
3. I can convert between mass, moles, and atom or molecules using factor-label Unit 9: Energy of Chemical Changes
methods.
Nature of Science Goal—Science provides technology to improve lives.
Unit 6: Mixtures and Their Properties 1. I can classify evidence of energy transformation (temperature change) as
endothermic or exothermic.
Nature of Science Goal-- Science provides predictable results.
2. I can describe how electrical energy can be produced in a chemical reaction and
1. I can use the terms solute and solvent in describing a solution.
identify which element gained and which element lost electrons.
2. I can sketch a solution, colloid, and suspension at the particle level.
3. I can identify the parts of a battery, including anode, cathode, and salt bridge.
3. I can describe the relative amount a solute particles in concentrated and
dilute solutions. Unit 10: Nuclear Changes
4. I can calculate concentration in terms of molarity and molality.
Nature of Science Goal—Correct interpretation of data replaces fear and superstition.
5. I can describe the colligative properties of solutions. (Boiling point elevation,
1. I can compare the charge, mass, energy, and penetrating power of alpha, beta, and
Freezing point depression, Vapor pressure lowering) in terms of every day
gamma radiation and recognize that of the products of the decay of an unstable
applications.
nucleus include radioactive particles and wavelike radiation.
6. I can identify which solution of a set would have the lowest freezing point or highest
2. I can interpret graphical data of decay processes to determine half-life and the age
boiling point.
of a radioactive substance.
Unit 7: Describing Chemical Reactions 3. I can compare and contrast the amount of energy released in a nuclear reaction to the
amount of energy released in a chemical reaction.
Nature of Science Goal—Conservations laws are investigated to explore
4. I can describe the differences between fission and fusion.
science relationships.
5. I can describe scientific evidence that all matter in the universe has a common origin.
1. I can classify a change as chemical or physical and give evidence of chemical
changes reactions.
2. I can describe the principles of collision theory and relate frequency, energy
of collisions, and addition of a catalyst to reaction rate.
3. I can write a chemical equation to describe a simple chemical reaction.
4. I can balance chemical reactions and recognize that the number of atoms in
a chemical reaction does not change.
5. I can classify reactions as synthesis, decomposition, single replacement,
double replacement or combustion.
6. I can use molar relationships in a balanced chemical reaction to predict the mass
of product produced in a simple chemical reaction that goes to completion.
7. I can explain the concept of dynamic equilibrium as it relates to chemical reactions.
8. I can describe whether reactants or products are favored in equilibrium when
given the equilibrium constant.
9. I can predict the effect of adding or removing either a product or a reactant or the
effect of changing temperature to shift equilibrium.

Unit 8: Describing Acids and Bases

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 Chapter 1: Introduction to Chemistry & the Nature of Science
1.1: The Process of Science
Objectives Scientific Methods of Problem Solving
Explain the necessity for experimentation In the 16th and 17th centuries, innovative thinkers were developing a new way to
In an experiment, identify the independent, dependent, and controlled variables. discover the nature of the world around them. They were developing a method that relied
upon making observations of phenomena and insisting that their explanations of the nature of
the phenomena corresponded to the observations they made.
The scientific method is a method of investigation involving experimentation and
observation to acquire new knowledge, solve problems, and answer questions. Scientists
Introduction frequently list the scientific method as a series of steps. Other scientists oppose this listing of
Socrates (469 B.C. - 399 B.C.), Plato (427 steps because not all steps occur in every case, and sometimes the steps are out of order. The
B.C. - 347 B.C.), and Aristotle (384 B.C. - 322 B.C.) scientific method is listed in a series of steps here because it makes it easier to study. You
are among the most famous of the Greek should remember that not all steps occur in every case, nor do they always occur in order.
philosophers. Plato was a student of Socrates, and
Aristotle was a student of Plato. These three were The Steps in the Scientific Method
probably the greatest thinkers of their time. Step 1: Identify the problem or
Aristotle's views on physical science profoundly phenomenon that needs explaining. This
shaped medieval scholarship, and his influence is sometimes referred to as "defining the
extended into the Renaissance (14th century - 16th problem."
century). Aristotle's opinions were the authority on Step 2: Gather and organize data on the
nature until well into the 1300s. Unfortunately, many problem. This step is also known as
of Aristotle's opinions were wrong. It is not intended "making observations."
here to denigrate Aristotle's intelligence; he was Step 3: Suggest a possible solution or
without doubt a brilliant man. It was simply that he explanation. A suggested solution is
was using a method for determining the nature of the called a hypothesis.
physical world that is inadequate for that task. The Step 4: Test the hypothesis by making
philosopher's method was logical thinking, not new observations.
Image obtained from:
making observations on the natural world. This led to http://upload.wikimedia.org/wikipedia/c Step 5: If the new observations support
many errors in Aristotle's thinking on nature. Let's ommons/a/ae/Aristotle_Altemps_Inv857 the hypothesis, you accept the hypothesis
consider two of Aristotle's opinions as examples. 5.jpg for further testing. If the new
In Aristotle's opinion, men were bigger and observations do not agree with your
hypothesis, add the new observations to
stronger than women; therefore, it was logical to him that men would have more teeth than your observation list and return to Step 3.
women. Thus, Aristotle concluded it was a true fact that men had more teeth than women.
Apparently, it never entered his mind to actually look into the mouths of both genders and
count their teeth. Had he done so, he would have found that men and women have exactly Experimentation
the same number of teeth. Experimentation is the primary way through which science gathers evidence for
In terms of physical science, Aristotle thought about dropping two balls of exactly ideas. It is more successful for us to cause something to happen at a time and place of our
the same size and shape but of different masses to see which one would strike the ground choosing. When we arrange for the phenomenon to occur at our convenience, we can have
first. In his mind, it was clear that the heavier ball would fall faster than the lighter one and all our measuring instruments present and handy to help us make observations, and we can
he concluded that this was a law of nature. Once again, he did not consider doing an control other variables. Experimentation involves causing a phenomenon to occur when and
experiment to see which ball fell faster. It was logical to him, and in fact, it still seems where we want it and under the conditions we want. An experiment is a controlled method
logical. If someone told you that the heavier ball would fall faster, you would have no reason of testing an idea or to find patterns. When scientists conduct experiments, they are usually
to disbelieve it. In fact, it is not true and the best way to prove this is to try it. seeking new information or trying to verify someone else's data.
Eighteen centuries later, Galileo decided to actually get two balls of different masses, Experimentation involves changing and looking at many variables. The independent
but with the same size and shape, and drop them off a building (Legend says the Leaning variable is the part of the experiment that is being changed or manipulated. There can only be
Tower of Pisa), and actually see which one hit the ground first. When Galileo actually did one independent variable in any experiment. Consider, for example, that you were trying to
the experiment, he discovered, by observation, that the two balls hit the ground at exactly the determine the best fertilizer for your plants. It would be important for you to grow your plants
same time . . . Aristotle's opinion was, once again, wrong. with everything else about how they are grown being the same except for the fertilizer

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you were using. You would be changing the type
of fertilizer you gave the plants and this would be
the independent variable. If you also changed how much water the plants received, the type
of plants you were growing, and some of the plants were grown inside and others outside,
you could not determine whether or not it was actually the fertilizer that caused the plants to
grow better or if it was something else you had changed. This is why it is important that
The scientist found support for the hypothesis from this experiment; fresh water
there is only one independent variable.
freezes at a higher temperature than salt water. Much more support would be needed before
The dependent variable is what is observed or measured as a result of what
the scientist would be confident of this hypothesis. Perhaps she would ask other scientists
happened when the independent variable was changed. In the plant experiment described
to verify the work.
above, you might measure the height of the plant and record their appearance and color.
These would be the dependent variables. The dependent variable is also sometimes called In the scientist's experiment, it was necessary that she freeze the salt water and fresh
the resultant variable. water under exactly the same conditions. Why? The scientist was testing whether or not the
Controlled variables are conditions of the experiment that are kept the same for presence of salt in water would alter its freezing point. It is known that changing air pressure
various trials of the experiment. Once again, if we were testing how fertilizer affected how will alter the freezing point of water, so this and other variables must be kept the same, or
well our plants grew, we would want everything else about how the plants are grown to be they must be controlled variables.
kept the same. We would need to use the same type of plant (maybe green beans), give them
Example: In the experiment described above, identify the:
the same amount of water, plant them in the same location (all outside in the garden), give
them all the same pesticide treatment, etc. These would be controlled variables. a) independent variable(s)
Suppose a scientist, while walking along the beach on a very cold day following a b) dependent variable(s)
rainstorm, observed two pools of water in bowl shaped rocks near each other. One of the c) controlled variable(s)
pools was partially covered with ice, while the other pool had no ice on it. The unfrozen pool Solution:
seemed to be formed from seawater splashing up on the rock from the surf, but the other a) Remember, the independent variable is what the scientist changed in his/her
pool was too high for seawater to splash in, so it was more likely to have been formed from experiment. In this case, the scientist added salt to one container and not to another
rainwater. container. The independent variable is whether or not salt was added.
The scientist wondered why one pool was partially frozen and not the other, since b) Dependent variables are what we look for as a result of the change we made. The scientist
both pools were at the same temperature. By tasting the water (not a good idea), the scientist recorded the temperature and physical state (liquid or solid) over time. These are the
determined that the unfrozen pool tasted saltier than the partially frozen one. The scientist dependent variables.
thought perhaps salt water had a lower freezing point than fresh water, and she decided to c) Controlled variables are kept the same throughout all of the trials. The scientist selected
go home and try an experiment to see if this were true. So far, the scientist has identified a identical containers, put the same amount of water in the containers, and froze them in the
question, gathered a small amount of data, and suggested an explanation. In order to test this same conditions in the same freezer. These are all controlled variables.
hypothesis, the scientist will conduct an experiment during which she can make accurate
observations. Suppose you wish to determine which brand of microwave popcorn (independent
For the experiment, the scientist prepared two identical variable) leaves the fewest unpopped kernels (dependent variable). You will need a supply of
containers of fresh water and added some salt to one of them. various brands of microwave popcorn to test and you will need a microwave oven. If you
A thermometer was placed in each liquid and these were put used different brands of microwave ovens with different brands of popcorn, the percentage
in a freezer. The scientist then observed the conditions and of unpopped kernels could be caused by the different brands of popcorn, but it could also be
temperatures of the two liquids at regular intervals. caused by the different brands of ovens. Under such circumstances, the experimenter would
not be able to conclude confidently whether the popcorn or the oven caused the difference.
To eliminate this problem, you must use the same microwave oven for every test. By using
the same microwave oven, you control many of the variables in the experiment. What if you
allowed the different samples of popcorn to be cooked at different temperatures? What if you
allowed longer heating periods? In order to reasonably conclude that the change in one
The Temperature and Condition of Fresh The Temperature and Condition of Salt variable was caused by the change in another specific variable, there must be no other
Water in a Freezer Water in a Freezer
Time (min) Temp (°C) Condition Time (min) Temp (°C) Condition
variables in the experiment. All other variables must be kept constant or controlled.
0 25 Liquid 0 25 Liquid When stating the purpose of an experiment, it is important to clarify the independent
5 20 Liquid 5 20 Liquid and dependent variables. The purpose is frequently stated in a sentence such as:
10 15 Liquid 10 15 Liquid “To see how changing _____________ affects ____________.”
15 10 Liquid 15 10 Liquid in which the independent variable is listed in the first blank, and the dependent variable
20 5 Liquid 20 5 Liquid
25 0 Frozen
is listed in the second blank.
25 0 Liquid
30 -5 Frozen 30 -5 Frozen
In the popcorn experiment, we would state the purpose as: “To see how changing the
brand of popcorn affects the percentage of unpopped kernels”. The independent variable is

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the brand of popcorn and the dependent variable
is what percentage of the popcorn didn’t pop. In
the salt water experiment described earlier, we would state the purpose as “To see how
adding salt to water affects the temperature the water freezes.”

Lesson Summary b) Gary wanted to make sure the size of the container did not affect plant growth in his
experiment.
Scientists use experimentation to test their ideas.
In an experiment, it is important to include only one independent variable (to c) Gary wanted to control how much plant food his plants received.
change only one thing in the experiment) d) Gary wanted his garden to look organized.
The dependent variable is what is measured or observed as a result of how the e) There is no possible scientific reason for having the same size containers.
independent variable changed. 2) What scientific reason might Gary have for insisting that all plants receive the
Controlled variables are those which are kept the same throughout various trials in the same amount of water every day?
experiment. a) Gary wanted to test the effect of shade on plant growth and therefore, he wanted to
have no variables other than the amount of sunshine on the plants.
Vocabulary b) Gary wanted to test the effect of the amount of water on plant growth.
c) Gary's hypothesis was that water quality was affecting plant growth.
Experiment: A controlled method of testing a hypothesis.
d) Gary was conserving water.
Controlled experiment: An experiment that compares the results of an
e) There is no possible scientific reason for having the same amount of water for
experimental sample to a control sample.
each plant every day.
Further Reading / Supplemental Links 3) What was the variable being tested in Gary's experiment (what is the
independent variable)?
http://learner.org/resources/series61.html: The learner.org website allows users to
a) The amount of water
view streaming videos of the Annenberg series of chemistry videos. You are
b) The amount of plant food
required to register before you can watch the videos but there is no charge. The
c) The amount of soil
website has two videos that apply to this lesson. One is a video called The World of
d) The amount of sunshine
Chemistry that relates chemistry to other sciences and daily life. Another video
e) The type of soil
called Thinking Like Scientists relates to the scientific method. The audience on the
4) Which of the following factors may be varying in Gary's experimental setup that he
video is young children but the ideas are full grown.
did not control?
Website of the James Randi Foundation. James Randi is a staunch opponent of fake
a) Individual plant variation
science. http://www.randi.org/site/
b) Soil temperature due to different colors of containers
Websites dealing with the history of the scientific method.
c) Water loss due to evaporation from the soil
http://www.historyguide.org/earlymod/lecture10c.html
d) The effect of insects which may attack one set of plants but not the other
http://www.history.boisestate.edu/WESTCIV/science/
5) A student decides to set up an experiment to determine the relationship between the
1.1: Review Questions
growth rate of plants and the presence of detergent in the soil. He sets up 10 seed pots. In
Use the following paragraph to answer questions 1-4: five of the seed pots, he mixes a precise amount of detergent with the soil. The other five
Gary noticed that two plants which his mother planted on the same day that were the same seed pots have no detergent in the soil. The five seed pots with detergent are placed in
size when planted were different in size after three weeks. Since the larger plant was in the the sun and the five seed pots with no detergent are placed in the shade. All 10 seed pots
full sun all day and the smaller plant was in the shade of a tree most of the day, Gary receive the same amount of water and the same number and type of seeds. He grows the
believed the sunshine was responsible for the difference in the plant sizes. In order to test plants for two months and charts the growth every two days. What is wrong with his
this, Gary bought ten small plants of the same size and type. He made sure they had the experiment?
same size and type of pot. He also made sure they have the same amount and type of soil.
a) The student has too few pots.
Then Gary built a frame to hold a canvas roof over five of the plants while the other five
b) The student has two independent variables.
were nearby but out in the sun. Gary was careful to make sure that each plant received
c) The student has two dependent (resultant) variables.
exactly the same amount of water and plant food every day.
d) The student has no experimental control on the soil.
1) What scientific reason might Gary have for insisting that the container size for the
A scientist plants two rows of corn for experimentation. She puts fertilizer on row 1 but
all plants be the same?
does not put fertilizer on row 2. Both rows receive the same amount of sun and water. She
a) Gary wanted to determine if the size of the container would affect the plant growth.
checks the growth of the corn over the course of five months.
6) What is the independent variable in this experiment?
12 7) What is the dependent variable in this experiment?
8) What variables are controlled in this experiment?

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1.2: Hypothesis, Law, & Theory
Objectives
 Describe the difference between hypothesis and theory as scientific terms.
Describe the difference between a theory and scientific law.
Explain the concept of a model.
everything is made of atoms) or the germ theory of disease (which states that
Explain why scientists use models.
Explain the limitations of models as scientific representations of reality. many diseases are caused by germs). Our understanding of gravity is still a
work in progress. But the phenomenon of gravity, like evolution, is an
Introduction accepted fact. “
Although all of us have taken science classes throughout the course of our study,
Note some key features of theories that are important to understand from this
many people have incorrect or misleading ideas about some of the most important and basic
principles in science. We have all heard of hypotheses, theories, and laws, but what do they description:
really mean? Before you read this section, think about what you have learned about these Theories are explanations of natural phenomenon. They aren’t predictions (although
terms before. What do these terms mean to you? What do you read contradicts what you we may use theories to make predictions). They are explanations why we observe
thought? What do you read supports what you thought? something.
Theories aren’t likely to change. They have so much support and are able to explain
Hypotheses satisfactorily so many observations, that they are not likely to change. Theories can,
One of the most common terms used in science classes is a “hypothesis”. The word indeed, be facts. Theories can change, but it is a long and difficult process. In order
can have many different definitions, depending on the context in which it is being used: for a theory to change, there must be many observations or evidence that the theory
“An educated guess” – because it provides a suggested solution based on the cannot explain.
evidence. Note that it isn’t just a random guess. It has to be based on evidence to be Theories are not guesses. The phrase “just a theory” has no room in science. To be a
a scientific hypothesis. scientific theory carries a lot of weight; it is not just one person’s idea about
Prediction – if you have ever carried out a science experiment, you probably made something.
this type of hypothesis, in which you predicted the outcome of your experiment.
Tentative or Proposed explanation – hypotheses can be suggestions about why Laws
something is observed, but in order for it to be scientific, we must be able to test the Scientific laws are similar to scientific theories in that they are principles that can be
explanation to see if it works, if it is able to correctly predict what will happen in a used to predict the behavior of the natural world. Both scientific laws and scientific theories
situation, such as: if my hypothesis is correct, we should see ___ result when we are typically well-supported by observations and/or experimental evidence. Usually
perform ___ test. A hypothesis is very tentative; it can be easily changed. scientific laws refer to rules for how nature will behave under certain conditions, frequently
written as an equation. Scientific theories are more overarching explanations of how nature
Theories works and why it exhibits certain characteristics. As a comparison, theories explain why we
The United States National Academy of Sciences describes what a theory is as observe what we do and laws describe what happens.
follows: For example, around the year 1800, Jacques Charles and other scientists were
“Some scientific explanations are so well established that no new working with gases to, among other reasons, improve the design of the hot air balloon.
evidence is likely to alter them. The explanation becomes a scientific theory. These scientists found, after many, many tests, that certain patterns existed in the
In everyday language a theory means a hunch or speculation. Not so in observations on gas behavior. If the temperature of the gas increased, the volume of the gas
science. In science, the word theory refers to a comprehensive explanation of increased. This is known as a natural law. A law is a relationship that exists between
an important feature of nature supported by facts gathered over time. variables in a group of data. Laws describe the patterns we see in large amounts of data, but
Theories also allow scientists to make predictions about as yet unobserved do describe why the patterns exist.
phenomena.” A common misconception is that scientific theories are rudimentary ideas that will
“A scientific theory is a well-substantiated explanation of some eventually graduate into scientific laws when enough data and evidence has been
aspect of the natural world, based on a body of facts that have been accumulated. A theory does not change into a scientific law with the accumulation of new
repeatedly confirmed through observation and experimentation. Such fact- or better evidence. Remember, theories are explanations and laws are patterns we see in
supported theories are not "guesses" but reliable accounts of the real world. large amounts of data, frequently written as an equation. A theory will always remain a
The theory of biological evolution is more than "just a theory." It is as factual theory; a law will always remain a law.
an explanation of the universe as the atomic theory of matter (stating that A model is a description, graphic, or 3-D representation of theory used to help
enhance understanding. Scientists often use models when they need a way to communicate
14 their understanding of what might be very small (such as an atom or molecule) or very large
(such as the universe). A model is any simulation, substitute, or stand-in for what you are
actually studying. A good model contains the essential variables that you are concerned
with in the real system, explains all the observations on the real system, and is as simple as

15
possible. A model may be as uncomplicated as a sphere representing the earth or billiard
balls representing gaseous molecules, or as complex as mathematical equations representing
light.
http://en.wikipedia.org/wiki/Hypothesis Video on Demand –
Chemists rely on both careful observation and well-known physical laws. By putting
Modeling the Unseen
observations and laws together, chemists develop models. Models are really just ways of
(http://www.learner.org/resources/series61.html?pop=yes&pid=793#)
predicting what will happen given a certain set of circumstances. Sometimes these models
are mathematical, but other times, they are purely descriptive. 1.2: Review Questions
If you were asked to determine the contents of a box that cannot be opened, you
Multiple Choice
would do a variety of experiments in order to develop an idea (or a model) of what the box
1) A number of people became ill after eating oysters in a restaurant. Which of the
contains. You would probably shake the box, perhaps put magnets near it and/or determine
following statements is a hypothesis about this occurrence?
its mass. When you completed your experiments, you would develop an idea of what is
a) Everyone who ate oysters got sick.
inside; that is, you would make a model of what is inside a box that cannot be opened.
b) People got sick whether the oysters they ate were raw or cooked.
A good example of how a model is useful to scientists is how models were used to
c) Symptoms included nausea and dizziness.
explain the development of the atomic theory. As you will learn in a later chapter, the idea of
d) Bacteria in the oysters may have caused the illness.
the concept of an atom changed over many years. In order to understand each of the different
2) If the hypothesis is rejected (proved wrong) by the experiment, then:
theories of the atom according to the various scientists, models were drawn, and the
a) The experiment may have been a success.
concepts were more easily understood.
b) The experiment was a failure.
Chemists make up models about what happens when different chemicals are mixed
c) The experiment was poorly designed.
together, or heated up, or cooled down, or compressed. Chemists invent these models using
d) The experiment didn't follow the scientific method.
many observations from experiments in the past, and they use these models to predict what
3) A hypothesis is:
might happen during experiments in the future. Once chemists have models that predict the
a) A description of a consistent pattern in observations.
outcome of experiments reasonably well, those working models can help to tell them what
b) An observation that remains constant.
they need to do to achieve a certain desired result. That result might be the production of an
c) A theory that has been proven.
especially strong plastic, or it might be the detection of a toxin when it’s present in your
d) A tentative explanation for a phenomenon.
food.
4) A scientific law is:
Lesson Summary a) A description of a consistent pattern in observations.
b) An observation that remains constant.
A hypothesis is a tentative explanation that
CC – Tracy Poulsen c) A theory that has been proven.
can be tested by further investigation.
d) A tentative explanation for a phenomenon.
A theory is a well-supported explanation of observations.
5) A well-substantiated explanation of an aspect of the natural world is a:
A scientific law is a statement that summarizes the relationship between variables.
a) Theory.
An experiment is a controlled method of testing a hypothesis.
b) Law.
A model is a description, graphic, or 3-D representation of theory used to help
c) Hypothesis.
enhance understanding.
d) None of these.
Scientists often use models when they need a way to communicate their
6) Which of the following words is closest to the same meaning as hypothesis?
understanding of what might be very small (such as an atom or molecule) or very
a) Fact
large (such as the universe).
b) Law
Vocabulary c) Formula
d) Suggestion
Hypothesis: A tentative explanation that can be tested by further investigation.
e) Conclusion
Theory: A well-established explanation 7) Why do scientists sometimes discard theories?
Scientific law: A statement that summarizes the relationship between variables. a) The steps in the scientific method were not followed in order.
Model: A description, graphic, or 3-D representation of theory used to help b) Public opinion disagrees with the theory.
enhance understanding. c) The theory is opposed by the church.
d) Contradictory observations are found.
Further Reading / Supplemental Links 8) True/False: When a theory has been known for a long time, it becomes a law.
http://en.wikipedia.org/wiki/Scientific_theory

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1.3: Graphing
Objectives
 Correctly graph data utilizing dependent variable, independent variable, scale and
units of a graph, and best fit curve.
Recognize patterns in data from a graph.
When you draw a line graph, you
Solve for the slope of given line graphs.
should arrange the numbers on the axis to
use as much of the graph paper as you can.
If the lowest temperature in your data is
100 K and the highest temperature in your
data is 160 K, you should arrange for 100
K to be on the extreme left of your graph
Introduction and 160 K to be on the extreme right of
Volume Pressure
Scientists search for regularities and trends in your graph. The creator of the graph on the
(liters) (atm)
data. Two common methods of presenting data that aid 10.0 0.50
left did not take this advice and did not
in the search for regularities and trends are tables and 5.0 1.00 produce a very good graph. You should
graphs. The table below presents data about the pressure 3.33 1.50 also make sure that the axis on your graph
and volume of a sample of gas. You should note that all 2.50 2.00 are labeled and that your graph has a title.
tables have a title and include the units of the 2.00 2.50
measurements. 1.67 3.00 When constructing a graph, there are some
You may note a regularity that appears in this general principles to keep in mind:
table; as the volume of the gas decreases (gets smaller), Take up as much of the graph paper
its pressure increases (gets bigger). This regularity or as possible. The lowest x-value
trend becomes even more apparent in a graph of this should be on the far left of the paper
data. A graph is a pictorial representation of patterns and the highest x-value should be
using a coordinate system. When the data from the table on the far right side of the paper.
is plotted as a graph, the trend in the relationship Your lowest y-value should be near
between the pressure and volume of a gas sample the bottom of the graph and the
becomes more apparent. The graph gives the scientist highest y-value near the top.
information to aid in the search for the exact regularity Choose your scale to allow you to
that exists in these data. do this. You do not need to start CC – Tracy Poulsen
When scientists record their results in a data table, the independent variable is put in counting at zero.
the first column(s), the dependent variable is recorded in the last column(s) and the Count your x- and y-scales by consistent amounts. If you start counting your x-axis
controlled variables are typically not included at all. Note in the data table that the first where every box counts as 2-units, you must count that way the course of the entire
column is labeled “Volume (in liters)” and that the second column is labeled “Pressure (in axis. Your y-axis may count by a different scale (maybe every box counts as 5
atm). That indicates that the volume was being changed (the independent variable) to see instead), but you must count the entire y-axis by that scale.
how it affected the pressure (dependent variable). Both of your axis should be labeled, including units. What was measured along that
In a graph, the independent variable is recorded along the x-axis (horizontal axis) or axis and what unit was it measured in?
as part of a key for the graph, the dependent variable is recorded along the y-axis (vertical
axis), and the controlled variables are not included at all. Note in the data table that the X- For X-Y scatter plots, draw a best-fit-line or curve that fits your data, instead of
axis is labeled “Volume (liters)” and that the Y-axis is labeled “Pressure (atm). That connecting the dots. You want a line that shows the overall trend in the data, but
indicates that the volume was being changed (the independent variable) to see how it might not hit exactly all of your data points. What is the overall pattern in the data?
affected the pressure (dependent variable).
Reading Information from a Graph
Drawing Line Graphs When we draw a line graph from a set of data points, we are creating data points between
Reading information from a line graph is easier and more accurate as the size of the known data points. This process is called interpolation. Even though we may have four
actual data points that were measured, we assume the relationship that exists between the
graph increases. In the two graphs shown below, the first graph uses only a small fraction of the
quantities at the actual data points also exists at all the points on the line graph between the
space available on the graph paper. The second graph uses all the space available for the same
graph. If you were attempting to determine the pressure at a temperature of 260 K, using the actual data points. Consider the following set of data for the solubility of KClO3 in water.
graph on the left would give a less accurate result than using the graph on the right. The table shows that there are exactly six known data points. When the data is
graphed, however, the graph maker assumes that the relationship between the temperature

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and the solubility remains the same. The line is drawn by interpolating the data
points between the actual data points. line graph doesn't work. Additionally, each year represents a group that we are looking at,
and not a measured quantity. A bar graph is better suited for this type of data. From this bar
Solubility
Temperature (°C) graph, you could very quickly answer questions like, “Which year was most likely a
(g/100 mL H2O)
drought year for Trout Creek?”, and “Which year was Trout Creek most likely to have
0 3.3
20 7.3 suffered from a flood?”
40 13.9 Rainfall
60 23.8 Year
(inches)
80 37.5 1980 24.7
100 56.3 1981 21.2
1982 14.5
CC – Tracy Poulsen 1983 13.2
We can now reasonably certainly read data from the graph for points that were not 1984 21.1
actually measured. If we wish to determine the solubility of KClO3 at 70°C, we follow the 1985 16.8
vertical grid line for 70°C up to where it touches the graphed line and then follow the 1986 19.9
horizontal grid line to the axis to read the solubility. In this case, we would read the solubility 1987 29.2
1988 31.6
to be 30. g/100 mL of H2O at 70°C. 1989 21.0 CC – Tracy Poulsen
There are also occasions when scientists
wish to determine data points from a graph that are
not between actual data points but are beyond the Finding the Slope of a Graph
ends of the actual data points. Creating data points As you may recall from algebra, the
slope of the line may be determined from the Temperature vs. Volume for a Gas
beyond the end of the graph line, using the basic Volume of Gas
shape of the curve as a guide is called graph. The slope represents the rate at which Temperature (°C) (mL)
extrapolation. one variable is changing with respect to the 20 60
Suppose the graph for the solubility of other variable. For a straight-line graph, the 40 65
potassium chlorate has been made from just three slope is constant for the entire line but for a 60 70
non-linear graph, the slope is different at 80 75
actual data points. If the actual data points for the
CC – Tracy Poulsen different points along the line. For a straight- 100 80
curve were the solubility at 60°C, 80°C, and 120 85
100°C, the graph would be the solid line shown on line graph, the slope for all points along the
the graph above. If the solubility at 30°C was desired, we could extrapolate (the dotted line) line can be determined from any section of the
from the graph and suggest the solubility to be 5.0 g/100 mL of H2O. If we check on the graph. For a non-linear graph, the must be
more complete graph above, you can see that the solubility at 30°C is close to 10 g/100 mL determined for each point from data at that
point. Consider the given data table and the
of H2O. The reason the second graph produces such a poor answer is that the relationship
that appears in the less complete graph does not hold beyond the ends of the graph. For this linear graph that follows.
reason, extrapolation is only acceptable for graphs where there is evidence that the The relationship in this set of data is
relationship shown in the graph will be true beyond the ends of the graph. Extrapolation is linear, that is, it produces a straight-line graph.
more dangerous that interpolation in terms of possibly producing incorrect data. The slope of this line is constant at all points
In situations in which both the independent and dependent variables are measured or on the line. The slope of a line is defined as the
counted quantities, an X-Y scatter plot is the most useful and appropriate type of graph. A rise (change in vertical position) divided by the
line graph cannot be used for independent variables that are groups of data, or nonmeasured run (change in horizontal position).
data. In these situations in which groups of data, rather than exact measurements, were Frequently in science, all of our data points do CC – Tracy Poulsen
recorded as the independent variable, a bar graph can typically be used. Consider the data in not fall exactly on a line. In this situation, we
the following table. draw a best fit line, or a line that goes as close to all of our points as possible. When finding the
For this data, a bar graph is more appropriate because independent variable is a group, slope, it is important to use two points that are on the best fit line itself, instead of our measured
not a measurement (for example, everything that happened in 1980). The concept of the data points which may not be on our best fit line. For a pair of points on the line, the
average yearly rainfall halfway between the years 1980 and 1981 does not make sense, so a coordinates of the points are identified as (x 1, y1) and (x 2, y2). In this case, the points selected
are (260, 1.3) and (180, 0.9). The slope can then be calculated in the manner:
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 c) If the graph is a straight line, calculate the slope, including units.
d) What would you expect the mass of 2.5 mL of solution to have?
e) What volume would you expect 60 g of the solution to occupy?
Therefore, the slope of the line is 0.005 atm/K. The fact that the slope is positive indicates
that the line is rising as it moves from left to right and that the pressure increases by 0.005
atm for each 1 Kelvin increase in temperature. A negative slope would indicate that the line 4) Donna is completing an experiment to find the effect #4 data
was falling as it moves from left to right. of the concentration of ammonia on rate (or speed) of Concentration of
the reaction. She has collected the given data from her Time (s) ammonia (mol/L)
Lesson Summary time trials and is ready for the analysis. 0.71 2.40
Two common methods of presenting data that aid in the search for regularities a) Identify the independent and dependent variables in 1.07 2.21
and trends are tables and graphs. this experiment. 1.95 2.00
When we draw a line graph from a set of data points, we are creating data b) Draw a graph to represent the data, including a 5.86 1.53
points between known data points. This process is called interpolation. best-fit-line
10.84 1.30
Creating data points beyond the end of the graph line, using the basic shape of c) If the concentration of ammonia was 0.30 mol/L,
14.39 1.08
the curve as a guide is called extrapolation. how much time has passed?
d) After 8 seconds, what will be the 20.43 0.81
The slope of a graph represents the rate at which one variable is changing
approximate concentration of ammonia? 29.67 0.60
with respect to the other variable.
39.80 0.40
Vocabulary 49.92 0.20
Graph: a pictorial representation of patterns using a coordinate system
Interpolation: the process of estimating values between measured values
Extrapolation: the process of creating data points beyond the end of the graph
line, using the basic shape of the curve as a guide 5) Consider the data table for an experiment on #5 data
the behavior of gases. Temperature Pressure
Slope: the ratio of the change in one variable with respect to the other variable.
a) Identify the independent and dependent (°C) (mmHg)
Further Reading / Supplemental Links variables in this experiment. 10 726
Use the following link to create both x-y and bar graphs: b) Draw a graph to represent the data.
20 750
http://nces.ed.gov/nceskids/createagraph/default.aspx c) Calculate the slope, including units.
d) What would be the pressure at 55°C? 40 800
These websites offer more tips on graphing and interpreting data: 70 880
http://staff.tuhsd.k12.az.us/gfoster/standard/bgraph2.htm and e) What would be the pressure at 120°C?
100960
http://www.sciencebuddies.org/science-fair-projects/project_data_analysis.shtml

1.3: Review Questions

1) On a data table, where is the independent variable typically listed? What about the
dependent variable?
2) On a graph, how do you identify the #3 data
independent variable and dependent variable? Volume of Mass of
Solution (mL) Solution (g)
3) Andrew was completing his density lab for his 0.3 3.4
chemistry lab exam. He collected the given 0.6 6.8
data for volume and mass.
a) Identify the independent and dependent 0.9 10.2
variables in this experiment. 1.9 21.55
b) Draw a graph to represent the data, 2.9 32.89
including a best-fit-line. 3.9 44.23
4.9 55.57

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 Chapter 2: The Structure of the Atom So how could the Greek philosophers have known that
Democritus had a good idea with his theory of “atomos?" It
would have taken some careful observation and a few simple
2.1: Early Ideas of Atoms
experiments. Now you might wonder why Greek
Objectives philosophers didn’t perform any experiments to actually test
Give a short history of the concept of the atom. Democritus’ theory. The problem, of course, was that Greek
Describe the contributions of Democritus and Dalton to atomic theory. philosophers didn’t believe in experiments at all. Remember,
Summarize Dalton's atomic theory and explain its historical development. Greek philosophers didn’t trust their senses, they only trusted
the reasoning power of the mind.
Introduction
The early Greek philosophers tried to understand the
You learned earlier how all matter in the universe is made out of tiny building blocks
nature of the world through reason and logic, but not through
called atoms. All modern scientists accept the concept of the atom, but when the concept of
experiment and observation. As a result, they had some very
the atom was first proposed about 2,500 years ago, ancient philosophers laughed at the idea.
interesting ideas, but they felt no need to justify their ideas
It has always been difficult to convince people of the existence of things that are too small Greek philosophers tried to
understand the nature of the based on life experiences. In a lot of ways, you can think of the
to see. We will spend some time considering the evidence (observations) that convince
world through reason and Greek philosophers as being “all thought and no action.” It’s
scientists of the existence of atoms. logic but not through truly amazing how much they achieved using their minds, but
experiment and observation.
Democritus and the Greek Philosophers because they never performed any experiments, they missed or
rejected a lot of discoveries that they could have made otherwise. Greek philosophers
Before we discuss the experiments and evidence
dismissed Democritus’ theory entirely. Sadly, it took over two millennia before the theory
that have, over the years, convinced scientists that matter is
of atomos (or “atoms,” as they’re known today) was fully appreciated.
made up of atoms, it’s only fair to give credit to the man Unlike the Greek
who proposed “atoms” in the first place. About 2,500 years philosophers, John Dalton
believed in both logical Dalton's Atomic Theory
ago, early Greek philosophers believed the entire universe
was a single, huge, entity. In other words, “everything was
thinking and Although the concept of atoms is now widely accepted, this wasn’t always the case.
experimentation. Scientists didn’t always believe that everything was composed of small particles called
one.” They believed that all objects, all matter, and all
substances were connected as a single, big, unchangeable atoms. The work of several scientists and their experimental data gave evidence for what is
“thing.” now called the atomic theory.
One of the first people to propose “atoms” was a In the late 1700’s, Antoine Lavoisier, a French scientist, experimented with the
man known as Democritus. As an alternative to the reactions of many metals. He carefully measured the mass of a substance before reacting and
beliefs of the Greek philosophers, he suggested that again measured the mass after a reaction had occurred in a closed system (meaning that
atomos, or atomon – tiny, indivisible, solid objects - nothing could enter or leave the container). He found that no matter what reaction he looked
make up all matter in the universe. at, the mass of the starting materials was always equal to the mass of the ending materials.
Democritus then reasoned that changes occur when Democritus was known as “The This is now called the law of conservation of mass. This went contrary to what many
the many atomos in an object were reconnected or Laughing Philosopher.” It’s a good scientists at the time thought. For example, when a piece of iron rusts, it appears to gain
thing he liked to laugh, because most mass. When a log is burned, it appears to lose mass. In these examples, though, the reaction
recombined in different ways. Democritus even extended other philosophers were laughing at
his theory, suggesting that there were different varieties of his theories.
does not take place in a closed container and substances, such as the gases in the air, are able
atomos with different shapes, sizes, and masses. He to enter or leave. When iron rusts, it is combining with oxygen in the air, which is why it
thought, however, that shape, size and mass were the only properties differentiating the seems to gain mass. What Lavoisier found was that no mass was actually being gained or
different types of atomos. According to Democritus, other characteristics, like color and lost. It was coming from the air. This was a very important first step in giving evidence for
taste, did not reflect properties of the atomos themselves, but rather, resulted from the the idea that everything is made of atoms. The atoms (and mass) are not being created or
different ways in which the atomos were combined and connected to one another. destroyed. The atoms are simply reacting with other atoms that are already present.
Greek philosophers truly believed that, above all else, our understanding of the world In the late 1700s and early 1800s, scientists began noticing that when certain
should rely on “logic.” In fact, they argued that the world couldn’t be understood using our substances, like hydrogen and oxygen, were combined to produce a new substance, like
senses at all, because our senses could deceive us. Therefore, instead of relying on water, the reactants (hydrogen and oxygen) always reacted in the same proportions by mass.
observation, Greek philosophers tried to understand the world using their minds and, more In other words, if 1 gram of hydrogen reacted with 8 grams of oxygen, then 2 grams of
specifically, the power of reason. hydrogen would react with 16 grams of oxygen, and 3 grams of hydrogen would react with
24 grams of oxygen. Strangely, the observation that hydrogen and oxygen always reacted in

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the “same proportions by mass” wasn’t special. In fact, it turned out that the reactants in
every chemical reaction reacted in the same proportions by mass. This observation is
summarized in the law of definite proportions. Take, for example, nitrogen and hydrogen,
5. Atoms of one element can combine with atoms of another element to form
which react to produce ammonia. In chemical reactions, 1 gram of hydrogen will react with
“compounds” – new, complex particles. In a given compound, however, the
4.7 grams of nitrogen, and 2 grams of hydrogen will react with 9.4 grams of nitrogen. Can
you guess how much nitrogen would react with 3 grams of hydrogen? Scientists studied different types of atoms are always present in the same relative numbers.
reaction after reaction, but every time the result was the same. The reactants always reacted
Lesson Summary
in the same proportions.
2,500 years ago, Democritus suggested that all matter in the universe was made up of
At the same time that scientists were finding this
tiny, indivisible, solid objects he called “atomos.”
pattern out, a man named John Dalton was experimenting with
Other Greek philosophers disliked Democritus’ “atomos” theory because they felt
several reactions in which the reactant elements formed more
it was illogical.
than one type of product, depending on the experimental
Dalton used observations about the ratios in which elements will react to combine and
conditions he used. One common reaction that he studied was
The Law of Conservation of Mass to propose his Atomic Theory. Dalton’s Atomic
the reaction between carbon and oxygen. When carbon and
Theory states:
oxygen react, they produce two different substances – we’ll
1. Matter is made of tiny particles called atoms.
call these substances “A” and “B.” It turned out that, given the
2. Atoms are indivisible. During a chemical reaction, atoms are rearranged, but
same amount of carbon, forming B always required exactly
they do not break apart, nor are they created or destroyed.
twice as much oxygen as forming A. In other words, if you can
3. All atoms of a given element are identical in mass and other properties.
make A with 3 grams of carbon and 4 grams of oxygen, B can
4. The atoms of different elements differ in mass and other properties.
be made with the same 3 grams of carbon, but with 8 grams
5. Atoms of one element can combine with atoms of another element to form
oxygen. Dalton asked himself – why does B require 2 times as
“compounds” – new complex particles. In a given compound, however, the different
much oxygen as A? Why not 1.21 times as much oxygen, or
types of atoms are always present in the same relative numbers.
0.95 times as much oxygen? Why a whole number like 2?
The situation became even stranger when Dalton tried Vocabulary
similar experiments with different substances. For example,
when he reacted nitrogen and oxygen, Dalton discovered that he could make three different Atom: Democritus’ word for the tiny, indivisible, solid objects that he believed made
substances – we’ll call them “C,” “D,” and “E.” As it turned out, for the same amount of up all matter in the universe
nitrogen, D always required twice as much oxygen as C. Similarly, E always required exactly Dalton’s Atomic Theory: the first scientific theory to relate chemical changes to the
four times as much oxygen as C. Once again, Dalton noticed that J.J. Thomson conducted structure, properties, and behavior of the atom
small whole numbers (2 and 4) seemed to be the rule. This experiments that suggested
observation came to be known as the law of multiple that Dalton’s atomic theory Further Reading / Supplemental Links
wasn’t telling the entire
proportions. story. To see a video documenting the early history of the concept of the atom, go to
Dalton thought about his results and tried to find some http://www.uen.org/dms/. Go to the k-12 library. Search for “history of the atom”.
theory that would explain it, as well as a theory that would explain the Law of Conservation Watch part 01. (you can get the username and password from your teacher)
of Mass (mass is neither created nor destroyed, or the mass you have at the beginning is Vision Learning: From Democritus to Dalton:
equal to the mass at the end of a change). One way to explain the relationships that Dalton http://visionlearning.com/library/module_viewer.php?c3=&mid=49&l=
and others had observed was to suggest that materials like nitrogen, carbon and oxygen were
composed of small, indivisible quantities which Dalton called “atoms” (in reference to 2.1: Review Questions
Democritus’ original idea). Dalton used this idea to generate what is now known as Dalton’s 1) (Multiple choice) Which of the following is not part of Dalton’s Atomic Theory?
Atomic Theory which stated the following: a) matter is made of tiny particles called atoms.
1. Matter is made of tiny particles called atoms. b) during a chemical reaction, atoms are rearranged.
2. Atoms are indivisible (can’t be broken into smaller particles). During a chemical c) during a nuclear reaction, atoms are split apart.
reaction, atoms are rearranged, but they do not break apart, nor are they created d) all atoms of a specific element are the same.
or destroyed.
3. All atoms of a given element are identical in mass and other properties. 2) Democritus and Dalton both suggested that all matter was composed of small
4. The atoms of different elements differ in mass and other properties. particles, called atoms. What is the greatest advantage Dalton’s Atomic Theory had
over Democritus’?

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3) It turns out that a few of the ideas in Dalton’s Atomic Theory aren’t entirely correct.
Are inaccurate theories an indication that science is a waste of time?
separating the cathode and anode by a short distance, the cathode ray tube can generate
what are known as cathode rays – rays of electricity that flow from the cathode to the anode.
2.2: Further Understanding of the Atom J. J. Thomson wanted to know what cathode rays were, where cathode rays came from, and
Objectives whether cathode rays had any mass or charge. The techniques that J. J. Thomson used to
Explain the observations that led to Thomson's discovery of the electron. Describe answer these questions were very clever and earned him a Nobel Prize in physics. First, by
Thomson's "plum pudding" mode of the atom and the evidence for it Draw a cutting a small hole in the anode, J. J. Thomson found that he could get some of the cathode
diagram of Thomson's "plum pudding" model of the atom and explain why it has rays to flow through the hole in the anode and into the other end of the glass cathode ray
this name. tube. Next, J. J. Thomson figured out that if he painted a substance known as “phosphor”
onto the far end of the cathode ray tube, he could see exactly where the cathode rays hit
because the cathode rays made the phosphor glow.
Describe Rutherford's gold foil experiment and explain how this experiment J. J.
altered the "plum pudding" model. Thomson must
Draw a diagram of the Rutherford model of the atom and label the nucleus and have suspected that
the electron cloud. cathode rays were
charged, because
Introduction his next step was to
Dalton's Atomic Theory held up well to a lot of the place a positively
different chemical experiments that scientists performed to test it. charged metal plate
In fact, for almost 100 years, it seemed as if Dalton's Atomic on one side of the
Theory was the whole truth. However, in 1897, a scientist named cathode ray tube
J. J. Thomson conducted some research that suggested that and a negatively
Dalton’s Atomic Theory wasn’t the entire story. As it turns out, charged metal plate Thomson’s experiment with cathode rays found that the ray moved away
Dalton had a lot right. He was right in saying matter is made up on the other side of from negatively charged plates and toward positively charges plates. What
of atoms; he was right in saying there are different kinds of the cathode ray does this say about the charge of the ray? CC – Tracy Poulsen
atoms with different mass and other properties; he was “almost” tube, as shown in
right in saying atoms of a given Figure 3. The metal
element are identical; he was right in saying during a chemical plates didn’t actually touch the cathode ray tube, but they were close enough that a
reaction, atoms are merely rearranged; he was right in saying remarkable thing happened! The flow of the cathode rays passing through the hole in the
a given compound always has atoms present in the same anode was bent upwards towards the positive metal plate and away from the negative metal
relative numbers. But he was WRONG in saying atoms were plate. Using the “opposite charges attract, like charges repel” rule, J. J. Thomson argued
indivisible or indestructible. As it turns out, atoms are that if the cathode rays were attracted to the positively charged metal plate and repelled
divisible. In fact, atoms are composed of even smaller, more fundamental particles. These from the negatively charged metal plate, they themselves must have a negative charge!
particles, called subatomic particles, are particles that are smaller than the atom. We’ll talk J. J. Thomson then did some rather complex experiments with magnets, and used his
about the discoveries of these subatomic particles next. results to prove that cathode rays were not only negatively charged, but also had mass.
Remember that anything with mass is part of what we call matter. In other words, these
Thomson’s Plum Pudding Model cathode rays must be the result of negatively charged “matter” flowing from the cathode to
In the mid-1800s, scientists were beginning to realize that the study of chemistry and the anode. But there was a problem. According to J. J. Thomson’s measurements, either
the study of electricity were actually related. First, a man named Michael Faraday showed these cathode rays had a ridiculously high charge, or else had very, very little mass – much
how passing electricity through mixtures of different chemicals could cause chemical less mass than the smallest known atom. How was this possible? How could the matter
reactions. Shortly after that, scientists found that by forcing electricity through a tube filled making up cathode rays be smaller than an atom if atoms were indivisible? J. J. Thomson
with gas, the electricity made the gas glow! Scientists didn’t, however, understand the made a radical proposal: maybe atoms are divisible. J. J. Thomson suggested that the small,
relationship between chemicals and electricity until a British physicist named J. J. Thomson negatively charged particles making up the cathode ray were actually pieces of atoms. He
began experimenting with what is known as a cathode ray tube. called these pieces “corpuscles,” although today we know them as electrons. Thanks to his
The figure shows a basic diagram of a cathode ray tube like the one J. J. Thomson clever experiments and careful reasoning, J. J. Thomson is credited with the discovery of the
would have used. A cathode ray tube is a small glass tube with a cathode (a negatively electron.
charged metal plate) and an anode (a positively charged metal plate) at opposite ends. By

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 Now imagine what would happen if atoms were made entirely of electrons. First of J.J. Thomson had measured the charge to mass ratio of the electron, but had been
all, electrons are very, very small; in fact, electrons are about 2,000 times smaller than the unable to accurately measure the charge on the electron. With his oil drop experiment, Robert
smallest known atom, so every atom would have to contain a whole lot of electrons. But Millikan was able to accurately measure the charge of the electron. When combined with the
there’s another, even bigger problem: electrons are negatively charged. Therefore, if atoms charge to mass ratio, he was able to calculate the mass of the electron. What Millikan did was
were made entirely out of electrons, atoms would be negatively charged themselves… and to put a charge on tiny droplets of oil and measured their rate of descent. By varying the
that would mean all matter was negatively charged as well. Of course, matter isn’t negatively charge on different drops, he noticed that the electric charges on the drops were all multiples
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charged. In fact, most matter is what we call neutral – it has no charge at all. If matter is of 1.6x10 C, the charge on a single electron.
composed of atoms, and atoms are composed of negative electrons, how can matter be
neutral? The only possible explanation is that atoms consist of more than just electrons. Rutherford’s Nuclear Model
Atoms must also contain some type of positively charged material that balances the negative Everything about Thomson’s experiments suggested
charge on the electrons. Negative and positive charges of equal size cancel each other out, the “plum pudding” model was correct – but according to the
just like negative and positive numbers of equal size. What do you get if you add +1 and -1? scientific method, any new theory or model should be tested
You get 0, or nothing. That’s true of numbers, and that’s also true of charges. If an atom by further experimentation and observation. In the case of the
contains an electron with a -1 charge, but also some form of material with a +1 charge, “plum pudding” model, it would take a man named Ernest
overall the atom must have a (+1) + (-1) = 0 charge – in other words, the atom must be Rutherford to prove it inaccurate. Rutherford and his
neutral, or have no charge at all. experiments will be the topic of the next section.
Based on the fact that atoms are neutral, and based on J. J. Thomson’s discovery that Disproving Thomson’s “plum pudding” model began
atoms contain negative subatomic particles called “electrons,” scientists assumed that atoms with the discovery that an element known as uranium emits
must also contain a positive substance. It turned out that this positive substance was another positively charged particles called alpha particles as it
kind of subatomic particle, known as the proton. Although scientists knew that atoms had to undergoes radioactive decay. Radioactive decay occurs when
contain positive material, protons weren’t actually discovered, or understood, until quite a bit one element decomposes into another element. It only happens
later. with a few very unstable elements. Alpha particles themselves Ernest Rutherford
When Thomson discovered the negative electron, he realized that atoms had to contain didn’t prove anything about the structure of the atom, they
positive material as well – otherwise they wouldn’t be neutral overall. As a result, Thomson were, however, used to conduct some very interesting experiments.
formulated what’s known as the “plum pudding” model for the atom. According to Ernest Rutherford was fascinated by all aspects of alpha particles. For the most part,
the “plum pudding” model, the negative electrons were though, he seemed to view alpha particles as tiny bullets that he could use to fire at all
like pieces of fruit and the positive material was like the kinds of different materials. One experiment in particular, however, surprised Rutherford,
batter or the pudding. This made a lot of sense given and everyone else.
Thomson’s experiments and observations. Thomson had Rutherford found that
been able to isolate electrons using a cathode ray tube; when he fired alpha particles
however he had never managed to isolate positive at a very thin piece of gold
particles. As a result, Thomson theorized that the positive foil, an interesting thing
material in the atom must form something like the happened. Almost all of the
“batter” in a plum pudding, while the negative electrons alpha particles went straight
must be scattered through this “batter.” (If you’ve never Thomson’s plum pudding model through the foil as if they’d
was much like a chocolate chip
seen or tasted a plum pudding, you can think of a cookie. Notice how the chocolate hit nothing at all. This was
chocolate chip cookie instead. In that case, the positive chips are the negatively charged what he expected to happen.
material in the atom would be the “batter” in the electrons, while the positive charge If Thomson’s model was
chocolate chip cookie, while the negative electrons would is spread throughout the entire
accurate, there was nothing
be scattered through the batter like chocolate chips.) hard enough for these small
Notice how easy it would be to pick the pieces of fruit out of a plum pudding. On particles to hit that would
the other hand, it would be a lot harder to pick the batter out of the plum pudding, because cause any change in their Ernest Rutherford's Gold Foil Experiment in which alpha particles were
the batter is everywhere. If an atom were similar to a plum pudding in which the electrons motion. shot at a piece of gold foil. Most of the particles went straight through, but
are scattered throughout the “batter” of positive material, then you’d expect it would be Every so often, some bounced straight back, indicating they were hitting a very small, very
easy to pick out the electrons, but a lot harder to pick out the positive material. though, one of the alpha dense particle in the atom.
CC – Tracy Poulsen
particles would be deflected

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2) Electrons (cathode rays) can be repelled by a negatively charged metal plate.
3) J.J. Thomson is credited with the discovery of the electron.
4) The plum pudding model is the currently accepted model of the atom
2.3: Protons, Neutrons, and Electrons in Atoms
#5-11: Match each conclusion regarding subatomic particles and atoms with the Objectives
observation/data that supports it. Describe the locations, charges, and masses and the three main subatomic particles.
Conclusion Observations Define atomic number.
5) All atoms have electrons a. Most alpha particles shot at gold foil go straight Describe the size of the nucleus in relation to the size of the atom.
through, without any change in their direction. Define mass number.
6) Atoms are mostly empty b. A few alpha particles shot at gold foil bounce in the Explain what isotopes are and how isotopes affect an element’s atomic mass.
space. opposite direction. Determine the number of protons, neutrons, and electrons in an atom.
7) Electrons have a negative c. Some alpha particles (with positive charges) when
Introduction
charge shot through gold foil bend away from the gold.
d. No matter which element Thomson put in a cathode Dalton’s Atomic Theory explained a lot about matter, chemicals, and chemical reactions.
8) The nucleus is positively Nevertheless, it wasn’t entirely accurate, because contrary to what Dalton believed, atoms can, in
ray tube, the same negative particles with the same
charged fact, be broken apart into smaller subunits or subatomic particles. We have been talking about the
properties (such as charge & mass) were ejected.
e. The particles ejected in Thomson’s experiment bent electron in great detail, but there are two other particles of interest to use: protons and neutrons.
9) Atoms have a small, dense In this section, we’ll look at the atom a little more closely.
away from negatively charged plates, but toward
nucleus
positively charged plates. Protons, Electrons, and Neutrons
We already learned that J.J. Thomson discovered a negatively charged particle, called
10) What is the name given to the tiny clump of positive material at the center of an atom?
the electron. Rutherford proposed that these electrons orbit a positive nucleus. In subsequent
experiments, he found that there is a smaller positively charged particle in the nucleus which
11) Electrons are ______ negatively charged metals plates and ______ positively
is called a proton. There is a third subatomic particle, known as a neutron. Ernest Rutherford
charged metal plates. proposed the existence of a neutral particle, with the approximate mass of a proton. Years
later, James Chadwick proved that the nucleus of the atom contains this
Consider the following two paragraphs for #12-14 neutral particle that had been proposed by Ernest Rutherford. Chadwick observed that when
Scientist 1: Although atoms were once regarded as the smallest part of nature, they are beryllium is bombarded with alpha particles, it
composed of even smaller particles. All atoms contain negatively charged particles, emits an unknown radiation that has approximately
called electrons. However, the total charge of any atom is zero. Therefore, this means the same mass as a proton, but no electrical charge.
that there must also be positive charge in the atom. The electrons sit in a bed of Chadwick was able to prove that the beryllium
positively charged mass. emissions contained a neutral particle - Rutherford’s
Scientist 2: It is true that atoms contain smaller particles. However, the electrons are not neutron.
floating in a bed of positive charge. The positive charge is located in the central part of As you might have already guessed from its
the atom, in a very small, dense mass, called a nucleus. The electrons are found outside name, the neutron is neutral. In other words, it has Electrons are much smaller than
of the nucleus. no charge whatsoever, and is therefore neither protons or neutrons. If an electron was
12) What is the main dispute between the two scientists’ theories? attracted to nor repelled from other objects. the mass of a penny, a proton or a
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13) Another scientist was able to calculate the exact charge of an electron to be -1.6x10 Neutrons are in every atom (with one exception), neutron would have the mass of a large
C. What effect does this have on the claims of Scientist 1? (Pick one answer) and they’re bound together with other neutrons and bowling ball!
a) Goes against his claim protons in the atomic nucleus.
b) Supports his claim Before we move on, we must discuss how the different types of subatomic
c) Has no effect on his claim. particles interact with each other. When it comes to neutrons, the answer is obvious. Since
14) If a positively charged particle was shot at a thin sheet of gold foil, what would the neutrons are neither attracted to, nor repelled from objects, they don’t really interact with
second scientist predict to happen? protons or electrons (beyond being bound into the nucleus with the protons).
Even though electrons, protons, and neutrons are all types of subatomic particles, they
are not all the same size. When you compare the masses of electrons, protons and neutrons,
what you find is that electrons have an extremely small mass, compared to either protons or
neutrons. On the other hand, the masses of protons and neutrons are fairly similar, although
34 technically, the mass of a neutron is slightly larger than the mass of a proton. Because
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protons and neutrons are so much more massive than electrons, almost all of the mass of any
atom comes from the nucleus, which contains all of the neutrons and protons.
 The table shown gives the
properties and locations of electrons,
protons, and neutrons. The third
Sub-Atomic Particles, Properties and Location protons in its nucleus, scientists are always interested in this number, and how this number
column shows the masses of the three
subatomic particles in grams. The
differs between different elements. Therefore, scientists give this number a special name. An
Relative Electric element’s atomic number is equal to the number of protons in the nuclei of any of its atoms.
Particle Mass Location
Charge The periodic table gives the atomic number of each element. The atomic number is a whole
(amu) number usually written above the chemical symbol of each element. The atomic number for
outside the hydrogen is 1, because every hydrogen atom has 1 proton. The atomic number for helium is 2
electron -1
nucleus because every helium atom has 2 protons. What is the atomic number of carbon?
proton 1 +1 nucleus Of course, since neutral atoms have to have one electron for every proton, an
neutron 1 0 nucleus element’s atomic number also tells you how many electrons are in a neutral atom of that
element. For example, hydrogen has an atomic number of 1. This means that an atom of
second column, however, shows the masses of the three subatomic particles in “atomic hydrogen has one proton, and, if it’s neutral, one electron as well. Gold, on the other
mass units”. An atomic mass unit (amu) is defined as one-twelfth the mass of a carbon-12 hand, has an atomic number of 79, which means that an atom of gold has 79 protons, and,
atom. Atomic mass units (amu) are useful, because, as you can see, the mass of a proton and if it’s neutral, and 79 electrons as well.
the mass of a neutron are almost exactly 1.0 in this unit system. The mass number of an atom is the total number of protons and neutrons in its
In addition to mass, another important property of subatomic particles is their charge. nucleus. Why do you think that the “mass number” includes protons and neutrons, but not
You already know that neutrons are neutral, and thus have no charge at all. Therefore, we say electrons? You know that most of the mass of an atom is concentrated in its nucleus. The mass
that neutrons have a charge of zero. What about electrons and protons? You know that of an atom depends on the number of protons and neutrons. You have already learned that the
electrons are negatively charged and protons are positively charged, but what’s amazing is mass of an electron is very, very small compared to the mass of either a proton or a neutron
that the positive charge on a proton is exactly equal in magnitude (magnitude means (like the mass of a penny compared to the mass of a bowling ball). Counting the number of
“absolute value” or “size when you ignore positive and negative signs”) to the negative protons and neutrons tells scientists about the total mass of an atom.
charge on an electron. The third column in the table shows the charges of the three subatomic
particles. Notice that the charge on the proton and the charge on the electron have the same mass number A = (number of protons) + (number of neutrons)
magnitude.
Negative and positive charges of equal magnitude cancel each other out. This An atom’s mass number is a very easy to calculate provided you know the number of
means that the negative charge on an electron perfectly balances the positive charge on the protons and neutrons in an atom.
proton. In other words, a neutral atom must have exactly one electron for every proton. If a
neutral atom has 1 proton, it must have 1 electron. If a neutral atom has 2 protons, it must Example:
have 2 electrons. If a neutral atom has 10 protons, it must have 10 electrons. You get the What is the mass number of an atom of helium that contains 2 neutrons?
idea. In order to be neutral, an atom must have the same number of electrons and protons. Solution:
(number of protons) = 2 (Remember that an atom of helium always has 2 protons.)
Atomic Number and Mass Number (number of neutrons) = 2
Scientists can distinguish
between different elements by counting mass number = (number of protons) + (number of neutrons)
the number of protons. If an atom has mass number = 2 + 2 = 4
only one proton, we know it’s a
hydrogen atom. An atom with two There are two main ways in which scientists frequently show the mass number of
protons is always a helium atom. If an atom they are interested in. It is important to note that the mass number is not given on
scientists count four protons in an the periodic table. These two ways include writing a nuclear symbol or by giving the name
atom, they know it’s a beryllium atom. of the element with the mass number written.
An atom with three protons is a lithium To write a nuclear symbol, the mass number is placed at the upper left
atom, an atom with five protons is a (superscript) of the chemical symbol and the atomic number is placed at the lower left
boron atom, an atom with six protons is (subscript) of the symbol. The complete nuclear symbol for helium-4 is drawn below.
a carbon atom… the list goes on. It is difficult to find qualities that are different from each
Since an atom of one element element and distinguish on element from another. Each
can be distinguished from an atom of element, however, does have a unique number of protons.
Sulfur has 16 protons, silicon has 14 protons, and gold has
another element by the number of The following nuclear symbols are for a nickel nucleus with 31 neutrons and a uranium
79 protons.
nucleus with 146 neutrons.
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In the nickel nucleus represented above, the atomic number 28 indicates the nucleus
contains 28 protons, and therefore, it must contain 31neutrons in order to have a mass 4 neutrons). Moreover, it always contains the two in the same relative amounts (or “relative
number of 59. The uranium nucleus has 92 protons as do all uranium nuclei and this abundances”). In a chunk of lithium, 93% will always be lithium with 4 neutrons, while the
particular uranium nucleus has 146 neutrons. remaining 7% will always be lithium with 3 neutrons.
The other way of representing these nuclei would be Nickel-59 and Uranium-238, Dalton always experimented with large chunks of an element – chunks that contained
where 59 and 238 are the mass numbers of the two atoms, respectively. Note that the mass all of the naturally occurring isotopes of that element. As a result, when he performed his
numbers (not the number of neutrons) is given to the side of the name. measurements, he was actually observing the averaged properties of all the different isotopes
in the sample. For most of our purposes in chemistry, we will do the same thing and deal
Isotopes with the average mass of the atoms. Luckily, aside from having different masses, most other
Unlike the number of protons, which is always the same in atoms of the same properties of different isotopes are similar.
element, the number of neutrons can be different, even in atoms of the same element. Atoms We can use what we know about atomic number and mass number to find the
of the same element, containing the same number of protons, but different numbers of number of protons, neutrons, and electrons in any given atom or isotope. Consider the
neutrons are known as isotopes. Since the isotopes of any given element all contain the same following examples:
number of protons, they have the same atomic number (for example, the atomic number of
helium is always 2). However, since the isotopes of a given element contain different Example: How many protons, electrons, and neutrons are in an atom of ?
numbers of neutrons, different isotopes have different mass numbers. The following two
Solution:
examples should help to clarify this point. Finding the number of protons is
Write the nuclear symbol for each element simple. The atomic number, #
Example: described: 21) 32 neutrons in an atom with mass
number of 58 of protons, is listed in the
a) What is the atomic number and the mass number of an isotope of lithium containing bottom right corner. # protons =
3 neutrons. A lithium atom contains 3 protons in its nucleus. 22) An atom with 10 neutrons and 9 protons.
19.
b) What is the atomic number and the mass number of an isotope of lithium containing 4
neutrons. A lithium atom contains 3 protons in its nucleus. Indicate the number of protons, neutrons, and electrons in each of the following
Solution: atoms: 24)2 Sodium-23 2
3) 24 He For all atoms with no charge, the number of electrons
a) atomic number = (number of protons) = 3 5) 1 H is equal to the
1

(number of neutrons) = 3 26) 2 number of protons. # 28)

Iron-55 7) 1737Cl electrons =19. Boron-11
mass number = (number of protons) + (number of neutrons)
mass number = 3 + 3 = 6 2 30) The mass number, 40, is the sum of the
9) 23892U Uranium-235
protons and the neutrons. To find the #
b) atomic number = (number of protons) = 3 of neutron, subtract the number of protons from the mass number. # neutrons = 40 – 19
(number of neutrons) = 4 = 21.
mass number = (number of protons) + (number of neutrons) 2.4: Atomic Mass
mass number = 3 + 4 = 7 Example: How many protons, electrons, and neutrons in an atom of
Objectives:
zinc-65?
Explain what is meant by the atomic mass of an element.
Notice that because the lithium atom always has 3 protons, the atomic number for
Calculate the atomic mass of an element from the masses and relative
lithium is always 3. The mass number, however, is 6 in the isotope with 3 neutrons, and 7 in percentages of the isotopes of the element.
the isotope with 4 neutrons. In nature, only certain isotopes exist. For instance, lithium exists Solution:
as an isotope with 3 neutrons, and as an isotope with 4 neutrons, but it doesn’t exists as an Introduction Finding the number of protons is simple. The atomic number, # of protons,
isotope with 2 neutrons, or as an isotope with 5 neutrons.
In chemistry we very rarely deal with only one isotope of an element. We use a
This whole discussion of isotopes brings us back to Dalton’s Atomic Theory.
mixture of the isotopes of an element in chemical reactions and other aspects of
According to Dalton, atoms of a given element are identical. But if atoms of a given element chemistry, because all of the isotopes of an element react in the same manner. That
can have different numbers of neutrons, then they can have different masses as well! How means that we rarely need to worry about the mass of a specific isotope, but instead we
did Dalton miss this? It turns out that elements found in nature exist as constant uniform need to know the average mass of the atoms of an element. Using the masses of the
mixtures of their naturally occurring isotopes. In other words, a piece of lithium always different isotopes and how abundant each isotope is, we can find the average mass of the
contains both types of naturally occurring lithium (the type with 3 neutrons and the type with atoms of an element. The atomic mass of an element is the weighted average mass of the
atoms in a naturally occurring sample of the element. Atomic mass is typically reported in
38 atomic mass units.
www.ck12.org is found on the periodic table. All zinc atoms have # protons = 30.
For all atoms with no charge, the number of electrons is equal to the number of
www.ck1 protons. # electrons =30. 1
2.org
The mass number, 65, is the sum of the protons and the neutrons. To find the # of neutron, subtract the number of protons from the charge on an electron. As a result, a neutral atom must have an equal number of
mass number. # neutrons = 65 – 30 = 35. protons and electrons.
Each element has a unique number of protons. An element’s atomic number is
Lesson Summary equal to the number of protons in the nuclei of any of its atoms.
Electrons are a type of subatomic particle with a negative charge. The mass number of an atom is the sum of the protons and neutrons in the atom
Protons are a type of subatomic particle with a positive charge. Protons are bound together in an atom’s nucleus as a result
of the strong nuclear force. 39
Neutrons are a type of subatomic particle with no charge (they’re neutral). Like protons, neutrons are bound into the atom’s www.ck12.org
nucleus as a result of the strong nuclear force.
Protons and neutrons have approximately the same mass, but they are both much more massive than electrons
(approximately 2,000 times as massive as an electron). The positive charge on a proton is equal in magnitude to the negative

Use the periodic table to find the symbol for the element with:
13) 44 electrons in a neutral atom
Isotopes are atoms of the same element (same number of protons) that have different
14) 30 protons
numbers of neutrons in their atomic nuclei. 15) An atomic number of 36
Vocabulary
Neutron: a subatomic particle with no charge
Atomic mass unit (amu): a unit of mass equal to one-twelfth the mass of a carbon- Write the nuclear symbol for each element
described: 21) 32 neutrons in an atom with mass
twelve atom number of 58
Atomic number: the number of protons in the nucleus of an atom 40
22) An atom with 10 neutrons and 9 protons.
Mass number: the total number of protons and neutrons in the nucleus of an atom www.ck12.org
Isotopes: atoms of the same element that have the same number of protons
Indicate the number of protons, neutrons, and electrons in each of the following
but different numbers of neutrons atoms: 24)
2 Sodium-23 2
3) 24 He
Further Reading / Supplemental Material 5) 11H
26) 2 28)
Jeopardy Game: http://www.quia.com/cb/36842.html Iron-55 7) 1737Cl Boron-11
For a Bill Nye video on atoms, go to http://www.uen.org/dms/. Go to the k-12 2 30)
library. Search for “Bill Nye atoms”. (you can get the username and password from 9) 23892U Uranium-235
your teacher)

2.3: Review Questions 2.4: Atomic Mass

Label each of the following statements as true or false. Objectives:
1) The nucleus of an atom contains all of the protons in the atom. Explain what is meant by the atomic mass of an element.
2) The nucleus of an atom contains all of the electrons in the atom. Calculate the atomic mass of an element from the masses and relative
3) Neutral atoms must contain the same number of neutrons as protons. percentages of the isotopes of the element.
4) Neutral atoms must contain the same number of electrons as protons.
Introduction
Match the subatomic property with its description. In chemistry we very rarely deal with only one isotope of an element. We use a
Sub-Atomic Particle Characteristics mixture of the isotopes of an element in chemical reactions and other aspects of
5) electron a. has a charge of +1 chemistry, because all of the isotopes of an element react in the same manner. That
6) neutron b. has a mass of approximately 1/1840 amu means that we rarely need to worry about the mass of a specific isotope, but instead we
need to know the average mass of the atoms of an element. Using the masses of the
7) proton c. is neither attracted to, nor repelled from charged objects different isotopes and how abundant each isotope is, we can find the average mass of the
atoms of an element. The atomic mass of an element is the weighted average mass of the
Indicate whether each statement is true or false. atoms in a naturally occurring sample of the element. Atomic mass is typically reported in
8) An element’s atomic number is equal to the number of protons in the nuclei of any of atomic mass units.
its atoms.
9) A neutral atom with 4 protons must have 4 electrons.
10) An atom with 7 protons and 7 neutrons will have a mass number of 14. 1
www.ck1
11) An atom with 7 protons and 7 neutrons will have an atomic number of 14. 2.org
12) A neutral atom with 7 electrons and 7 neutrons will have an atomic number of 14.
In the table below, Column 1 contains data for 5 different elements. Column 2 contains
data for the same 5 elements, however different isotopes of those elements. Match the atom
in the first column to its isotope in the second column.
 Original element Isotope of the same element
16) an atom with 2 protons and 1 neutron
a. a C (carbon) atom with 6 neutrons
17) a Be (beryllium) atom with 5 neutrons b. an atom with 2 protons and 2
neutrons
18) an atom with an atomic number of 6 and mass c. an atom with an atomic number of 7
number of 13 and a mass number of 15
19) an atom with 1 proton and a mass number of 1 d. an atom with an atomic number of 1
and 1 neutron
20) an atom with an atomic number of 7 and 7 e. an atom with an atomic number of 4
neutrons and 6 neutrons
The periodic table gives the atomic mass of
each element. The atomic mass is a number that
usually appears below the element’s symbol in each
Calculating Atomic Mass
square. Notice that atomic mass of boron (symbol
You can calculate the atomic mass (or average mass) of an element provided you B) is 10.8, which is what we calculated in example
know the relative abundances (the fraction of an element that is a given isotope) the
5, and the atomic mass of neon (symbol Ne) is
element’s naturally occurring isotopes, and the masses of those different isotopes. We can
20.18, which is what we calculated in example 6.
calculate this by the following equation:
Take time to notice that not all periodic tables have
Atomic mass = (%1)(mass1) + (%2)(mass2) + …
the atomic number above the element’s symbol and
Look carefully to see how this equation is used in the following examples.
the mass number below it. If you are ever confused,
Example: Boron has two naturally occurring isotopes. In a sample of boron, 20% of the remember that the atomic number should
atoms are B-10, which is an isotope of boron with 5 neutrons and a mass of 10 amu. The
other 80% of the atoms are B-11, which is an isotope of boron with 6 neutrons and a mass
42
of 11 amu. What is the atomic mass of boron?
Solution: Boron has two isotopes. We will use the equation: www.ck12.org
Atomic mass = (%1)(mass1) + (%2)(mass2) + …
Isotope 1: %1=0.20 (write all percentages as decimals), mass1=10
Isotope 2: %2=0.80, mass2=11

Substitute these into the equation, and we get:

Atomic mass = (0.20)(10) + (0.80)(11)
Atomic mass = 10.8 amu
The mass of an average boron atom, and thus boron’s atomic mass, is 10.8 amu.

Example: Neon has three naturally occurring isotopes. In a sample of neon, 90.92% of the
atoms are Ne-20, which is an isotope of neon with 10 neutrons and a mass of 19.99 amu.
Another 0.3% of the atoms are Ne-21, which is an isotope of neon with 11 neutrons and a
mass of 20.99 amu. The final 8.85% of the atoms are Ne-22, which is an isotope of neon
with 12 neutrons and a mass of 21.99 amu. What is the atomic mass of neon?
Solution:
Neon has three isotopes. We will use the equation:
Atomic mass = (%1)(mass1) + (%2)(mass2) + …
Isotope 1: %1=0.9092 (write all percentages as decimals), mass1=19.99
Isotope 2: %2=0.003, mass2=20.99
Isotope 3: %3=0.0885, mass3=21.99

Substitute these into the equation, and we get:

Atomic mass = (0.9092)(19.99) + (0.003)(20.99) + (0.0885)(21.99)
Atomic mass = 20.17 amu
The mass of an average neon atom is 20.17 amu
 When given two comparative colors or areas in the electromagnetic spectrum,
identify which area has the higher wavelength, the higher frequency, and the
higher energy.
always be the smaller of the two and will be a whole number, while the atomic mass should always be the larger of the two and will
Describe the relationship between wavelength, frequency, and energy of light
be a decimal number. waves (EMR)
Lesson Summary Introduction
An element’s atomic mass is the average mass of one atom of that element. An element’s atomic mass can be calculated Most of us are familiar with waves, whether they are waves of water in the ocean,
provided the relative abundances of the element’s naturally occurring isotopes, and the masses of those isotopes are known. waves made by wiggling the end of a rope, or waves made when a guitar string is plucked.
The periodic table is a convenient way to summarize information about the different elements. In addition to the element’s Light, also called electromagnetic radiation, is a special type of energy that travels as a
symbol, most periodic tables will also contain the element’s atomic number, and element’s atomic mass. wave.
Vocabulary Light Energy
Atomic mass: the weighted average of the masses of the isotopes of an element Before we talk about the different
forms of light or electromagnetic radiation
2.4: Review Questions (EMR), it is important to understand some
1) Copper has two naturally occurring isotopes. 69.15% of copper atoms are Cu-63 and have a mass of 62.93amu. The other of the general characteristics that waves
30.85% of copper atoms are Cu-65and have a mass of 64.93amu. What is the atomic mass of copper? share.
The high point of a wave is called
2) Chlorine has two isotopes, Cl-35 and Cl-37. Their abundances are 75.53% and 24.47% respectively. Calculate the atomic mass the crest. The low point is called the
of chlorine. trough. The distance from one point on a
wave to the same point on the next wave is called the wavelength of the wave. You could
2.5: The Nature of Light
Objectives 43
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determine the wavelength by measuring the distance from one trough to the next or from the
top (crest) of one wave to the crest of the next wave. The symbol used for wavelength is the
Greek letter lambda, .
Another important characteristic of waves is called frequency. The frequency of a
wave is the number of waves that pass a given point each second. If we choose an exact
position along the path of the wave and count how many waves pass the position each
second, we would get a value for frequency. Frequency has the units of cycles/sec or
waves/sec, but scientists usually just use units of 1/sec or Hertz (Hz).
8
All types of light (EMR) travels at the same speed, 3.0010 ‫ ڄ‬m/s. Because of this, as
the wavelength increases (the waves get longer), the frequency decreases (fewer waves
pass). On the other hand, as the wavelength decreases (the waves get shorter), the frequency
increases (more waves pass). CC – Tracy Poulsen
Electromagnetic waves (light waves) have an extremely wide range of wavelengths, cancer. The tiny section next in the spectrum is the
frequencies, and energies. The electromagnetic spectrum is the range of all possible visible range of light. These are the frequencies
frequencies of electromagnetic radiation. The highest energy form of electromagnetic (energies) of the electromagnetic spectrum to
waves is gamma rays and the lowest energy form (that we have named) is radio waves.
which the human eye responds. The highest form
On the far left of the figure above are the electromagnetic waves with the highest
of visible light energy is violet light, with red light
energy. These waves are called gamma rays and can be quite dangerous in large numbers to
having the lowest energy of all visible light. Even
living systems. The next lowest energy form of electromagnetic waves is called x-rays. Most of
lower in the spectrum, too low in energy to see,
you are familiar with the penetration abilities of these waves. They can also be dangerous to
living systems. Next lower, in energy, are ultraviolet rays. These rays are part of sunshine and
are infrared rays and radio waves.
rays on the upper end of the ultraviolet range can cause sunburns and eventually skin
44
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 Vocabulary
Frequency of a wave: The number of waves passing a specific point each second.
The light energies that are in the visible range are electromagnetic waves that cause the human eye to respond when those
Wavelength: The distance between a point on one wave to the same point on the next
frequencies enter the eye. The eye sends signals to the brain and the individual “sees” various colors. The highest energy waves in the
wave (usually from crest to crest or trough to trough).
visible region cause the brain to see violet and as the energy of the waves decreases, the colors change to blue, green, then to yellow,
orange, and red. When the energy of the wave is above or below the visible range, the eye does not respond to them. When the eye Electromagnetic spectrum: A list of all the possible types of light in order of
receives several different frequencies at the same time, the colors are blended by the brain. If all frequencies of light strike the eye decreasing frequency, or increasing wavelength, or decreasing energy. The
together the brain sees white, and if there are no frequencies striking the eye the brain sees black. electromagnetic spectrum includes gamma rays, X-rays, UV rays, visible light, IR
All the objects that you see around you are light absorbers – that is, the chemicals on the surface of the objects absorb certain radiation, microwaves and radio waves.
frequencies and not others. Your eyes detect the frequencies that strike your eye. Therefore, if your friend is wearing a red shirt, it
2.5: Review Questions
means that the dye in that shirt absorbs every frequency except red and the red is reflected. When the red frequency from the shirt
strikes your eye, your visual system sees red and you say the shirt is red. If your only light source was one exact frequency of blue 1) Which color of visible light has the longer wavelength, red or blue?
light and you shined it on a shirt that absorbed every frequency of light except one exact frequency of red, then the shirt would look 2) What is the relationship between the energy of electromagnetic radiation and the
black to you because no light would be reflected to your eye. The light from many fluorescent types of light do not contain all the frequency of that radiation?
frequencies of sunlight and so clothes inside a store may appear to be a slightly different color than when you get them home. 3) Of the two waves drawn below, which one has the most energy? How do you know?

Lesson Summary
Wave form energy is characterized by velocity, wavelength, and frequency.
As the wavelength of a wave increases, its frequency decreases. Longer waves with lower frequencies have lower energy.
Shorter waves with higher frequencies have higher energy. 45
Electromagnetic radiation has a wide spectrum, including low energy radio waves and very high energy gamma rays. www.ck12.org
The different colors of light differ in their frequencies (or wavelengths).

spectrum is composed of four individual frequencies. The pink color of the tube is the result of
4) List the following parts of the our eyes blending the four colors. Every Over time, our understanding of the atom has evolved.
The light emitted by the sign containing neon gas (on the
electromagnetic spectrum in order of left) is different from the light emitted by the sign atom has its own characteristic spectrum; Dalton’s model (on the left) was altered when Thomson
containing argon gas (on the right). discovered the electron and proposed the plum pudding
INCREASING energy: radio, gamma, no two atomic spectra are alike. The model (in the middle). Rutherford discovered the nucleus
UV, visible light, and infrared image below shows the emission and altered the model to the one on the right. Since then,
spectrum of iron. Because each element Neils Bohr and other scientists discovered more about the
location and energy of the electrons.
5) List the visible colors of light in order of INCREASING energy. has a unique emission spectrum, elements
can be identified using them.

2.6: Atoms and Electromagnetic Spectra

Objectives
Describe the appearance of an atomic emission spectrum. Atomic spectrum of iron.
Explain that each element has a unique emission spectrum. You may have heard or read about scientists discussing what elements are present in
Explain how an atomic (or emission) spectrum can be used to identify elements the sun or some more distant star, and after hearing that, wondered how scientists could
Describe an electron cloud that contains Bohr's energy levels. know what elements were present in a place no one has ever been. Scientists determine what
Explain the process through which an atomic spectrum is emitted according to Bohr’s elements are present in distant stars by analyzing the light that comes from stars and finding
model of atoms. the atomic spectrum of elements in that light. If the exact four lines that compose
hydrogen’s atomic spectrum are present in the light emitted from the star, that element
contains hydrogen.
Introduction b t
Electric light bulbs contain a very e h
thin wire in them that emits light when c e
heated. The wire is called a filament. The a
particular wire used in light bulbs is u l
made of tungsten. A wire made of any s i
metal would emit light under these e g
circumstances but tungsten was chosen h
t it emits contains virtually every
Bohr’s Model of the Atom
frequency and therefore, the light
emitted by tungsten appears By 1913, the evolution of our
white. A wire made of some other element would emit light of some color that was not concept of the atom had proceeded
convenient for our uses. Every element emits light when energized by heating or passing from Dalton’s indivisible spheres
electric current through it. Elements in solid form begin to glow when they are heated
idea to J. J. Thomson’s plum
sufficiently and elements in gaseous form emit light when electricity passes through them.
pudding model and then to
Rutherford’s nuclear atom theory.
This is the source of light emitted by neon signs and is also the source of light in a fire.
Rutherford, in addition to
Each Element Has a Unique Spectrum carrying out the brilliant experiment
The light frequencies emitted by atoms are mixed together by our eyes so that we see that demonstrated the presence of the
a blended color. Several physicists, including Angstrom in 1868 and Balmer in 1875, passed atomic nucleus, also proposed that
the light from energized atoms through glass prisms in such a way that the light was spread the electrons circled the nucleus in a
out so they could see the individual frequencies that made up the light. planetary type motion. The solar
The emission spectrum (or atomic spectrum) of a chemical element is the system or planetary model of the
unique pattern of light obtained when the element is subjected to heat or electricity. atom was attractive to scientists because it was similar to
something with which they were already familiar, namely the
solar system.
Unfortunately, there was a serious flaw in the planetary
model. It was already known that when a charged particle (such
Atomic spectrum of hydrogen
Bohr proposed that electrons have specific
as an electron) moves in a curved path, it gives off some form of
When hydrogen gas is placed into a tube and locations (or energy levels) around the nucleus, light and loses energy in doing so. This is, after all, how we
electric current passed through it, the color of emitted much like there are specific steps on a ladder. produce TV signals. If the electron circling the nucleus in an Niels Bohr and Albert
CC – Tracy Poulsen Einstein in 1925. Bohr
light is pink. But when the color is spread out, we see atom loses energy, it would necessarily have to move closer to
the nucleus as it loses energy and would eventually crash into received the Nobel prize
that the hydrogen for physics in 1922.
the nucleus. Furthermore, Rutherford’s model was unable to
46 describe how electrons give off light forming each element’s
www.ck12.org unique atomic spectrum. These difficulties cast a shadow on the
planetary model and indicated that, eventually, it would have to
replaced.

47
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weakn
esses
of
In 1913, the Danish physicist Niels Bohr proposed a model of the electron cloud of Bohr’s
an atom in which electrons orbit the nucleus and were able to produce atomic spectrum. model
Understanding Bohr’s model requires some knowledge of electromagnetic radiation (or was
light). that he
could
Energy Levels not
Bohr’s key idea in his model of the atom offer a
is that electrons occupy definite orbits that reason
require the electron to have a specific amount of why
energy. In order for an electron to be in the only
electron cloud of an atom, it must be in one of certain
the allowable orbits and it must have the precise energy
energy required for that orbit. Orbits closer to levels
the nucleus would require smaller amounts of or
energy for an electron and orbits farther from orbits
the nucleus would require the electrons to have were
a greater amount of energy. The possible orbits allowe
are known as energy levels. One of the d.
 Bohr hypothesized that the only way electrons could gain or lose energy would be to
move from one energy level to another, thus gaining or losing precise amounts of energy.
The energy levels are quantized, meaning that only specific amounts are possible. It would
atom moving energy levels. The electrons typically have the lowest energy possible, called
be like a ladder that had rungs only at certain heights. The only way you can be on that ladder
ground state. If the electrons are given energy (through heat, electricity, light, etc) the
is to be on one of the rungs and the only way you could move up or down would be to move
electrons in an atom could absorb energy by jumping to a higher energy level or an excited
to one of the other rungs. Suppose we had such a ladder with 10 rungs. Other rules for the
state. The electrons then give off the energy in the form of a piece of light, called a photon,
ladder are that only one person can be on a rung and in normal state, the ladder occupants
they had absorbed to fall back to a lower energy level. The energy emitted by electrons
must be on the lowest rung available. If the ladder had five people on it, they would be on the
dropping back to lower energy levels would always be precise amounts of energy because the
lowest five rungs. In this situation, no person could
differences in energy levels were precise. This explains why you see specific lines of light
move down because all the lower rungs are full.
when looking at an atomic spectrum – each line of light matches a specific "step down" that
Bohr worked out rules for the maximum number of
electrons that could be in each energy level in his an electron can take in that atom. This also explains why each element produces a different
model and required that an atom is in its normal atomic spectrum. Because each element has different acceptable energy levels for their
state (ground state) had all electrons in the lowest electrons, the possible steps each element’s electrons can take differ from all other elements.
energy levels available. Under these circumstances,
Lesson Summary
no electron could lose energy because no electron
could move down to a lower energy level. In this Bohr's model suggests each atom has a set of unchangeable energy levels and
way, Bohr’s model explained why electrons circling electrons in the electron cloud of that atom must be in one of those energy levels.
the nucleus did not emit energy and spiral into the Bohr's model suggests that the atomic spectra of atoms is produced by electrons
gaining energy from some source, jumping up to a higher energy level, then
nucleus.
immediately dropping back to a lower energy level and emitting the energy
Bohr’s Model and Atomic Spectra In Bohr’s model of the atom, electrons difference between the two energy levels.
absorb energy to move to higher energy The existence of the atomic spectra is support for Bohr's model of the atom.
The evidence used to support Bohr’s model levels and release energy to move to
came from the atomic spectra. He suggested that an lower energy levels.
Bohr's model was only successful in calculating energy levels for the hydrogen atom.
atomic spectrum is made by the electrons in an Obtained from:
http://en.wikipedia.org/wiki/File:Bohr_a Vocabulary
tom_model_English.svg Emission spectrum (or atomic spectrum): The unique pattern of light given off by an
48 element when it is given energy
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Energy levels: Possible orbits that an electron can have in the electron cloud of an
atom.
Ground state: to be in the lowest energy level possible
Excited state: to be in a higher energy level
Photon: a piece of electromagnetic radiation, or light, with a specific amount
of energy
Quantized: having specific amounts

Further Reading / Supplemental Links

A short discussion of atomic spectra and some animation showing the spectra of
elements you chose and an animation of electrons changing orbits with the
absorbtion and emission of light can be viewed at Spectral Lines
(http://www.colorado.edu/physics/2000/quantumzone/index.html)
Tutorial:
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf
Tutorial: http://visionlearning.com/library/module_viewer.php?mid=51&l=&c3=
Video: http://www.youtube.com/watch%3Fv%3DQI50GBUJ48s
Fireworks & How Electrons Emit Photons video:
http://www.youtube.com/watch%3Fv%3DncdmqhlTmGA

49
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Review Questions
1) Bohr’s model of the atom is frequently referred to as the “quantum model”. Why? What
does it mean to be quantized? How are electrons in atoms quantized? The following table summarizes the possible energy levels and sublevels, including
2) Each element produces a unique pattern of light due to different energies within the atom. the number of orbitals that compose each sublevel and the number of electrons the
Why would this information be useful in analyzing a material?
sublevel can hold when completely filled.
3) It was known that an undiscovered element (later named helium) was on the sun before it
Energy Level Sublevel # of orbitals Maximum #
was ever discovered on earth by looking at the sun’s spectrum. How do scientists know
(related to the distance (related to the in each of electrons
that the sun contains helium atoms when no one has even taken a sample of material
from the sun? from the nucleus) shape) sublevel possible
4) According to Bohr's theory, how can an electron gain or lose energy? 1 1s 1 2
5) What happens when an electron in an excited atom returns to its ground level? 2s 1 2
2
6) Why do electrons of an element release only a specific pattern of light? Why don’t they 2p 3 6
produce all colors of light? 3s 1 2
7) Use the following terms to explain how an electron releases a photon of light: electron, 3 3p 3 6
energy level, excited state, ground state, photon. Draw a picture if it is helpful. 3d 5 10
4s 1 2
4p 3 6
4 4d 5 10
2.6: Electron Arrangement in Atoms
4f 7 14
Objectives
List the order in which electron energy levels/sublevels will fill … … … …
Write the electron configuration and abbreviated electron configuration for a given There are several patterns to notice when looking at the table of energy levels. Each energy
level has one more sublevel than the level before it. Also, each new sublevel has two more
atom. th
orbitals. Can you predict what the 5 energy level would look like?
Introduction When determining where the electrons in an atom are located, a couple of rules must
Chemists are particularly interested in the electrons in an atoms electron cloud. be followed:
This is because the electrons determine the chemical properties of elements, such as what 1. Each added electron enters the orbitals of the lowest energy available.
compounds the element will form and which reactions it will participate in. In this section, 2. No more than two electrons can be placed in any orbital.
we will learn where the electrons are in atoms.
The Electron Configuration
Electron Energy Levels It would be convenient if the sublevels filled in the order listed in the table, such as
Although Bohr’s model was particularly useful for hydrogen, it did not work well 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, etc. However, this is not the order the electrons fill the
for elements larger than hydrogen. However, other physicists built on his model to create sublevels. Remember, the electrons will always go to the lowest energy available. When
one that worked for all elements. It was found that the energy levels used for hydrogen were that is taken into account, the actual filling order is:
further composed of sublevels of different shapes. These sublevels were composed of 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p…
orbitals in which the electrons were located. Note that 4s has lower energy than 3d and, therefore, will fill first. The filling order gets
more overlapped the higher you go.
An electron configuration lists the number of electrons in each used sublevel for an
atom. For example, consider the element gallium, with 31 electrons. Its first two electrons
would fit in the lowest energy possible, 1s. The next two would occupy 2s. 2p, with three
orbitals, can hold its next 6 electrons. Gallium continues to fill up its orbitals, finally putting
1 electron in 4s. The electron configuration for gallium would be:
2 2 6 2 6 2 10 1

31Ga: 1s 2s 2p 3s 3p 4s 3d 4p
The shape of p-orbitals Although you can choose to memorize the list and how many electrons fit in each
The shape of d-orbitals sublevel for the purpose of writing electron configurations, there is a way for us to find
this order by simply using our periodic table.

50
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 Abbreviated Electron Configuration
As the electron configurations become longer and longer, it becomes tedious to write
them out. A shortcut has been devised so that writing the configurations is less tedious. The
2 2 6 2 6 1
electron configuration for potassium is 1s 2s 2p 3s 3p 4s . The electron configuration for
potassium is the same as the electron configuration for argon except that it has one more
2 2 6 2 6
electron. The electron configuration for argon is 1s 2s 2p 3s 3p and in order to write the
1
electron configuration for potassium, we need to add only 4s . It is acceptable to use [Ar] to
1
represent the electron configuration for argon and [Ar]4s to represent the electron configuration
for potassium. Using this shortcut, the abbreviated electron configuration for calcium would be
2 2 1
[Ar] 4s and the electron configuration for scandium would be [Ar]4s 3d .
Even though the periodic table was organized according to the chemical behavior of
the elements, you can now see that the shape and design of the table is a perfect reflection
of the electron configuration of the atoms. This is because the chemical behavior of the
http://en.wikipedia.org/wiki/File:Electron_Configuration_Table.jpg elements is also caused by the electron configuration of the atoms.
Look at the different sections of the periodic table. You may have noticed that there
are several natural sections of the periodic table. The first 2 columns on the left make up the Example: Write the abbreviated electron configurations for
first section; the six columns on the right make up the next section; the middle ten columns (a) potassium, K
make up another section; finally the bottom fourteen columns compose the last section. Note (b) arsenic, As
the significance of these numbers: 2 electrons fit in any s sublevel, 6 electrons fit in any p (c) phosphorus, P
sublevel, 10 electrons fit in any d sublevel, and 14 electrons fit in any f sublevel. The four Solution:
1
sections described previously are known as the s, p, d, and f blocks respectively. (a) 19K: [Ar] 4s
2 10 3
If you move across the rows starting at the top left of the periodic table and move (b) 33As: [Ar] 4s 3d 4p
2 3
across each successive row, you can generate the same order of filling orbitals that was (c) 15P: [Ne] 3s 3p
listed before and also how many electrons total fit in each orbital. Starting at the top left, you
are filling 1s. Moving onto the second row, 2s is filled followed by 2p. Continuing with the Lesson Summary
filling order you generate the list: Electrons are located in orbitals, in various sublevels and energy levels of atoms
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p… Electrons will occupy the lowest energy level possible.
An electron configuration lists the sublevels the electrons occupy and the number of It is possible to write the electron configuration of en element using a periodic table.
electrons in each of those sublevels, written as superscripts.
Example: Write the electron configurations for Vocabulary
(a) potassium, K Electron configuration: a list that represents the arrangement of electrons of an atom.
(b) arsenic, As
(c) phosphorus, P Further Reading / Supplemental Links
Solution: http://www.ethbib.ethz.ch/exhibit/pauli/elektronenspin_e.html
(a) Potassium atoms have 10 protons and, therefore, 19 electrons. Using our chart, we see http://www.lorentz.leidenuniv.nl/history/spin/goudsmit.html
that the first sublevel to fill is 1s, which can hold 2 of those 19 electrons. Next to fill is 2s, http://en.wikipedia.org/wiki
which also holds 2 electrons. Then comes 2p which holes 6. We keep filling up the sublevels
until all 19 of the electrons have been placed in the lowest energy level possible. 4s only has 2.7: Review Questions
1 electron in it, although it can hold up to 2 electrons, because there are only 19 1) Which principal energy level holds a maximum of eight electrons?
2 2 6 2 6 1
electrons total in potassium. Its electron configuration is written as: 1s 2s 2p 3s 3p 4s 2) Which sub-energy level holds a maximum of six electrons?
2 2 6 2 6 2 10 3
(b) 33As: 1s 2s 2p 3s 3p 4s 3d 4p 3) Which sub-energy level holds a maximum of ten electrons?
2 2 6 2 3 4) If all the orbitals in the first two principal energy levels are filled, how many electrons
(c) 15P: 1s 2s 2p 3s 3p
are required?
52 5) In which energy level and sub-level of the carbon atom is the outermost electron located?
www.ck12.org 6) How many electrons are in the 2p sub-energy level of a neutral nitrogen atom?

53
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 2 2 6 2 1
7) Which element’s neutral atoms will have the electron configuration: 1s 2s 2p 3s 3p ? Chapter 3: The Organization of the Elements
8) What energy level and sub-level immediately follow 5s in the filling order?
9) What is the outermost energy level and sub-level used in the electron configuration 3.1: Mendeleev’s Periodic Table
of potassium?
Objectives
Write electron configurations for each of the following neutral atoms: Describe the method Mendeleev used to make his periodic table.
List the advantages and disadvantages Mendeleev’s table had over other methods of
10) Magnesium 14) Krypton
organizing the elements.
11) Nitrogen 15) Cesium
Explain how our current periodic table differs from Mendeleev’s original table.
12) Yttrium 16) Uranium
13) Tin Introduction
Write the abbreviated electron configuration for each of the following neutral atoms: During the 1800s, when most of the elements were being discovered, many chemists
tried to classify the elements according to their similarities. In 1829, Johann Döbereiner noted
17) Fluorine 20) Arsenic
chemical similarities in several groups of three elements and placed these elements into what he
18) Aluminum 21) Rubidium
called triads. His groupings included the triads of 1) chlorine, bromine, and iodine,
19) Titanium 22) Carbon
2) sulfur, selenium, and tellurium, 3) calcium, strontium, and barium, and 4) lithium,
All images, unless otherwise stated, are created by the CK-12 Foundation and are under sodium, and potassium. In all of the triads, the atomic weight of the second element was
almost exactly the average of the atomic weights of the first and third element.
the Creative Commons license CC-BY-NC-SA.
In 1864, John Newlands saw a connection between the chemical properties of
elements and their atomic masses. He stated that if the known elements, beginning with
lithium are arranged in order of increasing mass, the eighth element will have
properties similar to the first, the ninth similar to
the second, the tenth similar to the
third, and so on. Newlands called his
relationship the law of octaves,
comparing the elements to the notes
in a musical scale. Newlands tried to
force all the known elements to fit into John Newlands' law of octaves suggested that, if elements
his octaves but many of the heavier are aligned in order of increasing mass, every eighth
elements, when discovered, did not fit element would have similar properties.
into his patterns.
Mendeleev Organized His Table According to Chemical Behavior
By 1869, a total of 63 elements had been
discovered. As the number of known elements grew,
scientists began to recognize patterns in the way
chemicals reacted and began to devise ways to
classify the elements. Dmitri Mendeleev,
a Siberian-born Russian chemist, was the first
scientist to make a periodic table much like the one
we use today.
Mendeleev’s table listed the elements in
order of increasing atomic mass. Then he placed
elements underneath other elements with similar
chemical behavior. For example, lithium is a shiny Dmitri Mendeleev created the first
metal, soft enough to be cut with a spoon. It reacts periodic table in 1869.
readily with oxygen and reacts violently with water.

54 55
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W
h
en it reacts with water, it
produces hydrogen gas and
lithium hydroxide. As we
proceed through the elements How was Mendeleev Mendeleev’s
with increasing mass, we will able to make such accurate Property prediction for Actual properties of Gallium
come to the element sodium. Sodium is a shiny metal, soft enough to be cut with a spoon. It predictions? He understood Eka-aluminium
reacts readily with oxygen and reacts violently with water. When it reacts with water, it the patterns that appeared atomic mass 68 69.72
produces hydrogen gas and sodium hydroxide. You should note that the description of the between elements within a
density (g/cm³) 6.0 5.904
chemical behavior of sodium is very similar to the chemical description of lithium. When family, as well as patterns
melting point
Mendeleev found an element whose chemistry was very similar to a previous element, he according to increasing mass, Low 29.78
(°C)
placed it below the similar element. that he was able to fill in the
Ea2O3 (density - 5.5 -3
Mendeleev avoided missing pieces of the g cm-3) Ga2O3 (density - 5.88 g cm )
patterns. The ability to make oxide's formula (soluble in both alkalis and
Newlands’ mistake of trying (soluble in both
to force elements into groups accurate predictions is was acids)
alkalis and acids)
put Mendeleev’s table apart chloride's
from other organization Ea2Cl6 (volatile) Ga2Cl6 (volatile)
formula
systems that were made at
where their chemistry did not the same time and is what led Mendeleev’s Actual
match, but still ran into a few to scientists accepting his table and Property predictions for Eka- properties of
problems as he constructed his periodic law. silicon Germanium
table. Periodically, the atomic atomic mass 72 72.61
mass of elements would not be Changes to our Modern Periodic density (g/cm³) 5.5 5.35
in the right order to put them Table
in the correct group. For melting point (°C) high 947
Mendeleev's 1869 periodic table
The periodic table we use today
example, look at iodine and is similar to the one developed by color grey grey
tellurium on your periodic table. Tellurium is heavier than iodine, but he put it before iodine Mendeleev, but is not exactly the same. refractory
oxide type refractory dioxide
in his table, because iodine has properties most similar to fluorine, chlorine, and bromine. There are some important distinctions: dioxide
Additionally, tellurium has properties more similar to the group with oxygen in it. On his Mendeleev’s table did not oxide density
4.7 4.7
table, he listed the element its place according to its properties and put a question mark (?) include any of the noble gases, which (g/cm³)
next to the symbol. The question mark indicated that he was unsure if the mass had been were discovered later. These were oxide activity feebly basic feebly basic
measured correctly. added by Sir William Ramsay as Group chloride boiling
Another problem Mendeleev encountered was that sometimes the next heaviest 0, without any disturbance to the basic point under 100°C 86°C (GeCl4)
element in his list did not fit the properties of the next available place on the table. He would concept of the periodic table. (These chloride density
skip places on the table, leaving holes, in order to put the element in a group with elements elements were later moved to form (g/cm³) 1.9 1.9
with similar properties. For example, at the time the elements Gallium and Germanium had group 18 or 8A.) Other elements were
not yet been discovered. After zinc, arsenic was the next heaviest element he knew about, also discovered and put into their places on the periodic table.
but arsenic had properties most similar to nitrogen and phosphorus, not boron. He left two As previously noted, Mendeleev organized elements in order of increasing atomic
holes in his table for what he claimed were undiscovered elements. Note the dashes (-) with mass, with some problems in the order of masses. In 1914 Henry Moseley found a
a mass listed after it in his original table. These indicate places in which he predicted relationship between an element's X-ray wavelength and its atomic number, and therefore
elements would later be discovered to fit and his predicted mass for these elements. organized the table by nuclear charge (or atomic number) rather than atomic weight. Thus
Mendeleev went further with his missing elements by predicting the properties of Moseley placed argon (atomic number 18) before potassium (atomic number 19) based on
elements in those spaces. In 1871 Mendeleev predicted the existence of a yet-undiscovered their X-ray wavelengths, despite the fact that argon has a greater atomic weight (39.9) than
element he called eka-aluminium (because its location was directly under aluminum’s on potassium (39.1). The new order agrees with the chemical properties of these elements,
the table). The table below compares the qualities of the element predicted by Mendeleev since argon is a noble gas and potassium an alkali metal. Similarly, Moseley placed cobalt
with actual characteristics of Gallium (discovered in 1875). before nickel, and was able to explain that tellurium should be placed before iodine, not
Mendeleev made similar predictions for an element to fit in the place next to because of an error in measuring the mass of the elements (as Mendeleev suggested), but
silicon. Germanium, isolated in 1882, provided the best confirmation of the theory up to because tellurium had a lower atomic number than iodine.
that time, due to its contrasting more clearly with its neighboring elements than the two Moseley's research also showed that there were gaps in his table at atomic numbers 43
previously confirmed predictions of Mendeleev do with theirs. and 61 which are now known to be Technetium and Promethium, respectively, both

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radioactive and not naturally occurring. Following in the footsteps of Dmitri
Mendeleev, Henry Moseley also predicted new elements.
You already saw that the elements in vertical columns are related to each other by their
7) List three ways in which our current periodic table differs from the one originally made
electron configuration, but remember that Mendeleev did not know anything about electron
by Mendeleev.
configuration. He placed the elements in their positions according to their chemical behavior.
Thus, the vertical columns in Mendeleev’s table were composed of elements with similar
chemistry. These vertical columns are called groups or families of elements.
3.2: Metals, Nonmetals, and Metalloids
Lesson Summary Objectives
The periodic table in its present form was organized by Dmitri Mendeleev. Describe the differences among metals, nonmetals, and metalloids.
Mendeleev organized the elements in order of increasing atomic mass and in groups Identify an element as a metal, nonmetal, or metalloid given a periodic table or
of similar chemical behavior. He also left holes for missing elements and used the its properties.
patterns of his table to make predictions of properties of these undiscovered
elements. The modern periodic table now arranges elements in order of increasing Introduction
atomic number. Additionally, more groups and elements have been added as they In the periodic table, the elements are arranged according to similarities in their
have been discovered. properties. The elements are listed in order of increasing atomic number as you read from
left to right across a period and from top to bottom down a group. In this section you will
Vocabulary learn the general behavior and trends within the periodic table that result from this
arrangement in order to predict the properties of the elements.
Periodic table: a tabular arrangement of the chemical elements according to atomic
number. Metals, Non-metals, and Metalloids
Mendeleev: the Russian chemist credited with organizing the periodic table in the There is a progression from metals to non-metals across each row of elements in the
form we use today. periodic table. The diagonal line at the right side of the table separates the elements into two
Moseley: the chemist credited with finding that each element has a unique atomic groups: the metals and the non-metals. The elements that are on the left of this line tend to be
number metals, while those to the right tend to be non-metals (with the exception of hydrogen which

Further Reading / Supplemental Links

Tutorial: Vision Learning: The Periodic Table
http://visionlearning.com/library/module_viewer.php?mid=52&l=&c3=
How the Periodic Table Was Organized (YouTube):
http://www.youtube.com/watch%3Fv%3DCdkpoQk2LDE
For several videos and video clips describing the periodic table, go to
http://www.uen.org/dms/. Go to the k-12 library. Search for “periodic table”. (you
can get the username and password from your teacher)

3.1: Review Questions

1) What general organization did Mendeleev use when he constructed his table?
2) How did Mendeleev’s system differ from Newlands’s system?
3) Did all elements discovered at the time of Mendeleev fit into this organization system?
How would the discovery of new elements have affecting Mendeleev’s arrangement
of the elements?
The division of the periodic table into metals and non-metals. The metalloids are
4) Look at Mendeleev’s predictions for Germanium (ekasilicon). How was Mendeleev able
most of the elements along the line drawn. Additionally, the element hydrogen is a
to make such accurate predictions? NONMETAL, even though it is on the left side of the periodic table.
5) What problems did Mendeleev have when arranging the elements according to his CC – Tracy Poulsen
criteria? What did he do to fix his problems? is a nonmetal). The elements that are directly on the diagonal line are metalloids, with some
6) What discovery did Henry Moseley make that changed how we currently recognize exceptions. Aluminum touches the line, but is considered a metal. Metallic character generally
the order of the elements on the periodic table? increases from top to bottom down a group and right to left across a period, meaning that
francium (Fr) has the most metallic character of all of the discovered elements.
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 Most of the chemical elements are metals. Most metals have the common properties
of being shiny, very dense, and having high melting points. Metals tend to be ductile (can
be drawn out into thin wires) and malleable (can be hammered into thin sheets). Metals are
good conductors of heat and electricity. All metals are solids at room temperature except for 3.3: Valence Electrons
mercury. In chemical reactions, metals easily lose electrons to form positive ions. Examples Objectives
of metals are silver, gold, and zinc. Define valence electrons.
Nonmetals are generally brittle, dull, have low melting points, and they are generally Indicate the number of valence electrons for selected atoms.
poor conductors of head heat and electricity. In chemical reactions, they ted to gain electrons
Introduction
to form negative ions. Examples of non-metals are hydrogen, carbon, and nitrogen.
Metalloids have properties of both metals and nonmetals. Metalloids can be shiny or The electrons in the outermost shell are the valence electrons these are the
dull. Electricity and heat can travel through metalloids, although not as easily as they can electrons on an atom that can be gained or lost in a chemical reaction. Since filled d or f
through metals. They are also called semimetals. They are typically semi-conductors, which subshells are seldom disturbed in a chemical reaction, the valence electrons include those
means that they are elements that conduct electricity better than insulators, but not as well electrons in the outermost s and p sublevels.
as conductors. They are valuable in the computer chip industry. Examples of metalloids are Gallium has the following electron configuration.
2 10 1
silicon and boron. Ga: [Ar] 4s 3d 4p
The electrons in the fourth energy level are further from the nucleus than the electrons in the
Lesson Summary third energy level. The 4s and 4p electrons can be lost in a chemical reaction, but not the
There is a progression from metals to non-metals across each period of elements electrons in the filled 3d subshell. Gallium therefore has three valence electrons: the two in
in the periodic table. 4s and one in 4p.
Metallic character generally increases from top to bottom down a group and right to
left across a period. Determining Valence Electrons
The number of valence electrons for an atom can be seen in the electron
2 2 6 2
Vocabulary configuration. The electron configuration for magnesium is 1s 2s 2p 3s . The outer
energy level for this atom is n=3 and it has two electrons in this energy level. Therefore,
periodic law: states that the properties of the elements recur periodically as their magnesium has two valence electrons.
atomic numbers increase 2 2 6 2 4
The electron configuration for sulfur is 1s 2s 2p 3s 3p . The outer energy level
ductile: can be drawn out into thin wires in this atom is n=3 and it holds six electrons, so sulfur has six valence electrons.
malleable: can be hammered into thin sheets 2 2 6 2 6 2 10 1
The electron configuration for gallium is 1s 2s 2p 3s 3p 4s 3d 4p . The outer
energy level for this atom is n=4 and it contains three electrons. You must recognize that
3.2: Review Questions
even though the 3d sub-level is mixed in among the 4s and 4p sub-levels, 3d is NOT in the
Label each of the following elements as a metal, nonmetal, or metalloid. outer energy level and therefore, the electrons in the 3d sub-level are NOT valence electrons.
1) Carbon 4) Plutonium Gallium has three
2) Bromine 5) Potassium electrons in the outer
3) Oxygen 6) Helium energy level and therefore,
it has three valence
Given each of the following properties, label the property of as that of a metal, nonmetal, electrons. The
or metalloid. identification of valence
7) Lustrous 11) Insulators electrons is vital because
8) Semiconductors 12) Conductors the chemical behavior of
9) Brittle 13) Along the staircase an element is determined
10) Malleable primarily by the
arrangement of the The number of valence electrons in an atom can be easily found by
14) The elements mercury and bromine are both liquids at room temperature, but mercury is electrons in the valence counting the s and p columns in the periodic table.
considered a metal and bromine is considered a nonmetal. How can that be? What shell. CC – Tracy Poulsen
properties do metals and nonmetals have? This pattern can be summarized very easily, by merely counting the s and p blocks of
the periodic table to find the total number of valence electrons. One system of numbering the
groups on the periodic table numbers the s and p block groups from 1A to 8A. The number
indicates the number of valence electrons.
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Lesson Summary
 Valence electrons are the electrons in the outermost principal quantum level of
an atom.
The number of valence electrons is important, because the chemical behavior of an The same pattern is true of other groups on the periodic table. Remember, Mendeleev
element depends primarily by the arrangement of the electrons in the valence shell. arranged the table so that elements with the most similar properties were in the same group
Vocabulary on the periodic table.
It is important to recognize a couple of other important groups on the periodic
Valence electrons: the electrons in the outermost energy level of an atom.
table by their group name. Group 7A (or 17) elements are also called halogens. This
3.3: Review Questions group contains very reactive nonmetals elements.
1) How many valence electrons are present in the following electron configuration: The noble gases are in group 8A. These elements also have similar properties to each
2 2 6 2 3 other, the most significant property being that they are extremely unreactive rarely forming
1s 2s 2p 3s 3p ?
compounds. We will learn the reason for this later, when we discuss how compounds form.
2) How many valence electrons are present in the following electron configuration:
2 2 6 2 6 2 10 1 The elements in this group are also gases at room temperature.
1s 2s 2p 3s 3p 4s 3d 4p ?

For each of the following atoms, indicate the total number of valence electrons in each atom:
3) Fluorine 7) Aluminum
4) Bromine 8) Gallium
5) Sodium 9) Argon
6) Cesium 10) Krypton

3.4: Families and Periods of the Periodic Table

Objectives
Give the name and location of specific groups on the periodic table, including
alkali metals, alkaline earth metals, noble gases, halogens, and transition metals.
Explain the relationship between the chemical behavior of families in the
periodic table and their electron configuration.
Identify elements that will have the most similar properties to a given element.
Families of the periodic table.
Introduction CC – Tracy Poulsen
The chemical behavior of atoms is controlled by their electron configuration. Since An alternate numbering system numbers all of the s, p, and d block elements from 1-
the families of elements were organized by their chemical behavior, it is predictable that 18. In this numbering system, group 1A is group 1; group 2A is group 2; the halogens (7A)
the individual members of each chemical family will have similar electron configurations. are group 17; and the noble gases (8A) are group 18. You will come across periodic table
with both numbering systems. It is important to recognize which numbering system is being
Families of the Periodic Table used and to be able to find the number of valence electrons in the main block elements
Remember that Mendeleev arranged the periodic table so that elements with the most regardless of which numbering systems is being used.
similar properties were placed in the same group. A group is a vertical column of the
periodic table. All of the 1A elements have one valence electron. This is what causes these Periods of the Periodic
elements to react in the same ways as the other members of the family. The elements in 1A Table
are all very reactive and form compounds in the same ratios with similar properties with If you can locate
other elements. Because of their similarities in their chemical properties, Mendeleev put an element on the
these elements into the same group. Group 1A is also known as the alkali metals. Although Periodic Table, you can
most metals tend to be very hard, these metals are actually soft and can be easily cut. use the element’s
Group 2A is also called the alkaline earth metals. Once again, because of their position to figure out the
similarities in electron configurations, these elements have similar properties to each other. energy level of the
element’s valence
electrons. A period is a
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table. For example, the elements sodium (Na) and magnesium (Mg) are both in period 3.
The elements astatine (At) and radon (Rn) are both in period 6.

Lesson Summary
Which family is characterized by each of the following descriptions?
The vertical columns on the periodic table are called groups or families because
8) A very reactive family of nonmetals
of their similar chemical behavior.
9) Have 7 valence electrons
All the members of a family of elements have the same number of valence electrons
10) A nonreactive family of nonmetals
and similar chemical properties
11) Forms colorful compounds
The horizontal rows on the periodic table are called periods.
12) Have 2 valence electrons
Vocabulary 13) A very reactive family of metals
Group (family): a vertical column in the periodic table
Alkali metals: group 1A of the periodic table
Alkali earth metals: group 2A of the periodic table 3.5: Periodic Trends
Halogens: group 7A of the periodic table Objectives
Noble gases: group 8A of the periodic table Explain what is meant by the term periodic law
Transition elements: groups 3 to 12 of the periodic table Describe the general trend in atomic size for groups and periods.
Describe the trends that exist in the periodic table for ionization energy.
Further Reading / Supplemental Links Describe the trends that exist in the periodic table for electronegativity.
http://www.wou.edu/las/physci/ch412/perhist.htm
Introduction
http://www.aip.org/history/curie/periodic.htm
http://web.buddyproject.org/web017/web017/history.html We have talked in great detail about how the periodic table was developed, but we
http://www.dayah.com/periodic have yet to talk about where the periodic table gets its name. To be periodic means to “have
http://www.chemtutor.com/perich.htm repeating cycles” or repeating patterns. In the periodic table, there are a number of physical
properties that are “trend-like”. This means is that as you move down a group or across a
3.4: Review Questions period, you will see the properties changing in a general direction.
Multiple Choice The periodic table is a powerful tool that provides a way for chemists to organize the
1) Which of the following elements is in the same family as fluorine? chemical elements. The word “periodic” means happening or recurring at regular intervals.
a) silicon The periodic law states that the properties of the elements recur periodically as their atomic
b) antimony numbers increase. The electron configurations of the atoms vary periodically with their
c) iodine atomic number. Because the physical and chemical properties of elements depend on their
d) arsenic electron configurations, many of the physical and chemical properties of the elements tend
e) None of these. to repeat in a pattern.
2) Elements in a ______________ have similar chemical properties. The actual repeating trends that are observed have to do with three factors. These
a) period factors are:
b) family (1) The number of protons in the nucleus (called the nuclear charge).
c) both A and B (2) The number of energy levels holding electrons (and the number of electrons in the
d) neither A nor B outer energy level).
3) Which of the following elements would you expect to be most similar to carbon? (3) The number of electrons held between the nucleus and its outermost electrons
a) Nitrogen (called the shielding effect). This affects the attraction between the valence electrons
b) Boron and the protons in the nucleus.
c) Silicon
Trends in Atomic Radius
Give the name of the family in which each of the following elements is located: The atomic radius is a way of measuring the size of an atom. Although this is
4) astatine 6) barium difficult to directly measure, we are, in essence, looking at the distance from the nucleus
5) krypton 7) francium to the outermost electrons.
Let’s look at the atomic radii or the size of the atom from the top of a family or group
to the bottom. Take, for example, the Group 1 metals. Each atom in this family (and all other
64 main group families) has the same number of electrons in the outer energy level as all the
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other atoms of that family. Each row (period) in the periodic table represents another added
energy level. When we first learned about principal energy levels, we learned that each new
energy level was larger than the one before. Energy level 2 is larger than energy level 1,
Periodic Trends in Ionization Energy
energy level 3 is larger than energy level 2, and so on. Therefore, as we move down the 2 1
periodic table from period to period, each successive period represents the addition of a Lithium has an electron configuration of 1s 2s . Lithium has one electron in its
larger energy level. outermost energy level. In order to remove this electron, energy must be added. Look at
the equation below:
You can imagine that with the increase in the number of energy levels, the size of
the atom must increase. The increase in the number of energy levels in the electron cloud Li(g) + energy Li+(g) + e-
takes up more space. Therefore the trend within a group or family on the periodic table is With the addition of energy, a lithium ion can be formed from the lithium atom by losing one
that the atomic size increases with increased number of energy levels. electron. This energy is known as the ionization energy. The ionization energy is the energy
required to remove the most loosely held electron from a gaseous atom. The higher the value
of the ionization energy, the harder it is to remove that electron.

In order to determine the trend for the periods, we need to look at the number of Ionization Energies for
Ionization Energies for some Period 2 Elements
protons (nuclear charge), the number of energy levels, and the shielding effect. For a row in
the periodic table, the atomic number still increases (as it did for the groups) and thus the Element Ionization Energy
Group 1 Elements
number of protons would increase. When we examine the energy levels for period 2, we find First Ionization Lithium, Li 520 kJ/mol
that the outermost energy level does not change as we increase the number of electrons. In Element Beryllium, Be 899 kJ/mol
Energy
period 2, each additional electron goes into the second energy level. So the number of energy Boron, B 801 kJ/mol
Lithium, Li 520 kJ/mol
levels does not go up. As we move from left to right across a period, the number of electrons Carbon, C 1086 kJ/mol
Sodium, Na 495.5 kJ/mol Nitrogen, N 1400 kJ/mol
in the outer energy level increases but it is the same outer energy level. Potassium, K 801 kJ/mol
Looking at the elements in period 2, the Oxygen, O 1314 kJ/mol
number of protons increases from lithium with Fluorine, F 1680 kJ/mol
three protons, to fluorine with nine protons.
Therefore, the nuclear charge increases across a Consider the ionization energies for the elements in group 1A of the periodic table,
period. Meanwhile, the number of energy levels the alkali metals. Comparing the electron configurations of lithium to potassium, we know
occupied by electrons remains the same. The that the electron to be removed is further away from the nucleus, as the energy level of the
numbers of electrons in the outermost energy valence electron increase. Because potassium’s valence electron is further from the nucleus,
level increases from left to right across a period, there is less attraction between this electron and the protons and it requires less energy to
but how will this affect the radius? With an remove this electron. As you move down a family (or group) on the periodic table, the
increase in nuclear charge, there is an increase in ionization energy decreases.
the pull between the protons and the outer We can see a trend when we look at the ionization energies for the elements in period
level, pulling the outer electrons toward the nucleus. The net result is that the atomic size 2. When we look closely at the data presented in the table above, we can see that as we move
decreases going across the row. across the period from left to right, in general, the ionization energy increases. As we move
Considering all the information about atomic size, you will recognize that the largest across the period, the atoms become smaller which causes the nucleus to have greater
atom on the periodic table is all the way to the left and all the way to the bottom, francium, attraction for the valence electrons. Therefore, as you move from left to right in a period on
#87, and the smallest atom is all the way to the right and all the way to the top, helium, #2. the periodic table, the ionization energy increases.
The fact that the atoms get larger as you move downward in a family is probably Example: Which of the following has a greater ionization energy?
exactly what you expected before you even read this section, but the fact that the atoms get (a) As or Sb
smaller as you move to the right across a period is most likely a big surprise. Make sure (b) Ca or K
you understand this trend and the reasons for it. (c) Polonium or Sulfur
Example: Which of the following has a greater radius? Solution:
(a) As or Sb (a) As because it is above Sb in Group 15.
(b) Ca or K (b) Ca because it is further to the right on the periodic table.
(c) Polonium or Sulfur (c) S because it is above Po in Group 16.
Solution:
(a) Sb because it is below As in Group 15. Periodic Trends in Electronegativity
(b) K because it is further to the left on the periodic table. Around 1935, the American chemist Linus Pauling developed a scale to describe the
(c) Polonium because it is below Sulfur in Group 16. attraction an element has for electrons in a chemical bond. This is the electronegativity.

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The values of electronegativity are higher for elements that more strongly attract
electrons. On this Pauling scale fluorine, with an electronegativity of 4.0 is the most
electronegative element, and cesium and francium, with electronegativities of 0.7, are the
The atomic radius increases from the top to the bottom in any group and decreases
least electronegative.
from left to right across a period.
The electronegativity of atoms increases as you move from left to right across a
period in the periodic table. This is because as you go from left to right across a period, the Ionization energy is the energy required to remove the most loosely held electron
atoms of each element have the same number of energy levels. However, the nucleus charge from a gaseous atom or ion.
increases, so the attraction that the atoms have for the valence electrons increases. Ionization energy generally increases across a period and decreases down a
The electronegativity of atoms decreases as you move from top to bottom down group. The higher the electronegativity of an atom, the greater its ability to attract
a group in the periodic table. This shared electrons.
is because as you go from top The electronegativity of atoms increases as you move from left to right across a
to bottom down a group, the period in the periodic table and decreases as you move from top to bottom down
atoms of each element have an a group in the periodic table.
increasing number of energy
Vocabulary
levels.
Atoms with low Nuclear charge: the number of protons in the nucleus
ionization energies have low Shielding effect: the inner electrons help “shield” the outer electrons and the
electronegativities because their nucleus from each other.
nuclei do not have a strong Ionization energy: the energy required to remove the most loosely held electron from
attraction for electrons. Atoms a gaseous atom or ion
with high ionization energies Electronegativity: the ability of an atom in a molecule to attract shared electrons
have high electronegativities
because the nucleus has a strong 3.5: Review Questions
attraction for electrons. Multiple choice
1) Why is the table of elements called “the periodic table”?
Lesson Summary a) it describes the periodic motion of celestial bodies.
b) it describes the periodic recurrence of chemical properties.
c) because the rows are called periods.
d) because the elements are grouped as metals, metalloids, and non-metals.
e) None of these.
2) Which of the following would have the largest ionization energy?
a) Na
b) Al
c) H
d) He
3) Which of the following would have the smallest ionization energy?
a) K
b) P
c) S
d) Ca

Short Answer
4) Which of the following would have a smaller radius: indium or gallium?
CC – Tracy Poulsen 5) Which of the following would have a smaller radius: potassium or cesium?
6) Which of the following would have a smaller radius: titanium or polonium?
Atomic size is the distance from the nucleus to the valence shell where the 7) Explain why iodine is larger than bromine.
valence electrons are located. 8) Arrange the following in order of increasing atomic radius: Tl, B, Ga, Al, In.
68 9) Arrange the following in order of increasing atomic radius; Ga, Sn, C.
10) Define ionization energy.
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11) Place the following elements in order of increasing ionization energy: Na, S, Mg, Ar
12) Define electronegativity.

For each pair of elements, choose the element that has the lower electronegativity. Chapter 4: Describing Compounds
13) Li or N
14) Cl or Na 4.1: Introduction to Compounds
15) Na or K Objectives
16) Mg or F Explain the difference between an element, a compound and a mixture

All images, unless otherwise stated, are created by the CK-12 Foundation and are under Substances and Mixtures
the Creative Commons license CC-BY-NC-SA. Matter can be classified into two broad categories: pure substances and mixtures. A
pure substance is a form of matter that has a constant composition (meaning it’s the same
everywhere) and properties that are constant throughout the sample (meaning there is only
one set of properties such as melting point, color, boiling point, etc throughout the matter).
Elements and compounds are both example of pure substances.

CC – Tracy Poulsen

Mixtures are physical combinations of two or more elements and/or compounds. The
term “physical combination” refers to mixing two different substances together where the
substances do not chemically react. The physical appearance of the substances may change
but the atoms and/or molecules in the substances do not change.
The chemical symbols are used not only to represent the elements; they are also used to
write chemical formulas for the millions of compounds formed when elements chemically
combine to form compounds. The law of constant composition states that the ratio by mass of
the elements in a chemical compound is always the same, regardless of the source of the
70 compound. The law of constant composition can be used to distinguish between compounds and
www.ck12.org mixtures. Compounds have a constant composition, and mixtures do not. For example,
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pure water is always 88.8% oxygen and 11.2% hydrogen by weight, regardless of the
source of the water. Because water is a compound, it will always have this exact
composition. Brass is an example of a mixture. Brass consists of two elements, copper and oxide requires two atoms of aluminum and three atoms of oxygen. Therefore, we write the
zinc, but it can contain as little 10% or as much as 45% zinc.
formula for aluminum oxide as Al2O3. The symbol Al tells us that the compound contains
Consider the following examples including elements, compounds, and mixtures.
aluminum, and the subscript 2 tells us that there are two atoms of aluminum in each
molecule. The O tells us that the compound contains oxygen, and the subscript 3 tells us that
there are three atoms of oxygen in each molecule. It was decided by chemists that when the
Matter with only one subscript for an element is 1, no subscript would be used at all. Thus the chemical formula
Pure substance
type of atom is called MgCl2 tells us that one molecule of this substance contains one atom of magnesium and two
(element) an element. atoms of chlorine. In formulas that contain parentheses, such as Ca(OH)2, the subscript 2
applies to everything inside the parentheses. Therefore, this formula (calcium hydroxide)
contains one atom of calcium and two atoms of oxygen and two atoms of hydrogen.
Although the chlorine
atoms are bonded in Lesson Summary
Pure substance
pairs, since there is
(element) only one type of atom, Matter can be classified into two broad categories: pure substances and mixtures.
this is an element. A pure substance is a form of matter that has a constant composition and properties
that are constant throughout the sample.
Mixtures are physical combinations of two or more elements and/or compounds.
When two or more Elements and compounds are both example of pure substances.
Pure substance elements are bonded
(compound) together, a compound
Compounds are substances that are made up of more than one type of atom.
is produced. Elements are the simplest substances made up of only one type of atom.
Vocabulary
Element: a substance that is made up of only one type of atom
When two or more Compound: a substance that is made up of more than one type of atom
pure substances (in
this case, two
bonded together
Mixture elements) are Mixture: a combination of two or more elements or compounds which have not
combined, but not reacted to bond together; each part in the mixture retains its own properties
bonded together, a
mixture is produced. Further Reading / Supplemental Links
You may listen to Tom Lehrer’s humorous song “The Elements” with animation at
When two or more The Element Song (http://www.privatehand.com/flash/elements.html)
pure substances are
Mixture combined, but not
The learner.org website allows users to view streaming videos of the
bonded together, a Annenberg series of chemistry videos. You are required to register before you
mixture is produced. can watch the videos but there is no charge. Video on Demand – The World of
Chemistry (http://www.learner.org/resources/series61.html?pop=yes&pid=793#)
CC – Tracy Poulsen
4.1: Review Questions
The last couple of chapters have focused on elements and their properties. This unit Classify each of the following as an element, compound, or mixture.
will focus on compounds, including what compounds form and how elements combine to 1) Salt, NaCl
make compounds. Later chapters will cover mixtures. 2) Oil
3) Gold
Compounds and Chemical Formulas 4) Sugar, C6H12O6
The formula for a compound uses the symbols to indicate the type of atoms involved and 5) Salad dressing
uses subscripts to indicate the number of each atom in the formula. For example, aluminum 6) Salt water
combines with oxygen to form the compound aluminum oxide. To form aluminum 7) Water
8) Copper
72 9) Air
10) Milk
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(pronounced CAT-ions). The ions that are formed when an atom gains electrons are
negatively charged because they have more electrons in the electron cloud than protons in
the nucleus. Negatively charged ions are called anions (pronounced AN-ions). When main group nonmetals gain electrons to form anions, their names are changed to end
Predicting Charges of Main Group Ions in “-ide”. For example, fluorine atoms gain electrons to become fluoride ions.
1
All the metals in family 1A have electron configurations ending with an s electron in Polyatomic Ions
the outer energy level. For that reason, all family 1A members will tend to lose exactly one
Thus far, we have been dealing with ions made from single atoms. Such ions are
electron when they are ionized, obtaining an electron configuration like the closest noble
+ + + called monatomic ions. There also exists a group of polyatomic ions, ions composed of a
gas. The entire family forms +1ions: Li , Na , K , etc. We need to note that while hydrogen group of atoms that are covalently bonded and behave as if they were a single ion. Almost
is in this same column, it is not considered to be an metal. There are times that hydrogen acts all the common polyatomic ions are negative ions.
as if it is a metal and forms +1ions; however, most of the time it bonds with other atoms as a
A table of many common polyatomic ions is given. The more familiar you become
nonmetal. In other words, hydrogen doesn’t easily fit into any chemical family. All members
of family 1A form ions with +1charge. with polyatomic ions, the better you will be able to write names and formulas of ionic
The metals in family 2A all have electron configurations ending with two electrons in compounds. It is also important to note that there are many polyatomic ions that are not on
2
an s position in the outermost energy level. To have an electron configuration like the this chart.
closest noble gas, each of the elements in this family will lost two valence electrons and
2+ 2+
form +2 ions; Be , Mg , etc. Other metal elements’ charges can be predicted using the Cations
same patterns. Members of family 3A form ions with 3+ charge. +1
Family 5A nonmetals at the top will gain electrons to form negative ions. By gaining Ammonium, NH4 + Anions
electrons, they are able to obtain the electron configuration of the noble gas closest to them
on the periodic table. Family 6A non-metals will gain two electrons to obtain a octet thus -1 -2 -3
- - - 2-
forming a -2 ion. Family 7A will form -1 ions: F , Cl , etc. Hypochlorite, ClO Sulfite, SO3 Phosphate, PO4 3-

Family 8A, of course, is the noble gases and has no tendency to either gain or lose Chlorite, ClO2 - Sulfate, SO4 2-

electrons so they do not form ions. Chlorate, ClO3 -

The charges ions form can be summarized as in the following table. Many of the Perchlorate, ClO4 -
transition elements have variable oxidation states so they can form ions with different Nitrite, NO2 - Carbonate, CO3 2-
charges, and, therefore, are left off of this chart. Nitrate, NO3 - -
Bicarbonate, HCO3
- 2-
Hydroxide, OH Peroxide, O2
Acetate, C2H3O2 - Oxalate, C2O42-
2-
Silicate, SiO3
2-
Thiosulfate, S2O3
Permanganate, MnO4 - Chromate, CrO4 2- 2-
-
Cyanide, CN Dichromate, Cr2O7
-
Thiocyanate, SCN

Naming Transition Metals

Some metals are capable of forming ions with various charges. These include most
2+
of the transition metals and many post transition metals. Iron, for example, may form Fe
3+
ions by losing 2 electrons or Fe ions by losing 3 electrons. The rule for naming these ions
is to insert the charge (oxidation number) of the ion with Roman numerals in parentheses
after the name. These two ions would be named iron (II) and iron (III). When you see that
CC – Tracy Poulsen the compound involves any of the variable oxidation number metals (iron, copper, tin, lead,
nickel, and gold), you must determine the charge (oxidation number) of the metal from the
formula and insert Roman numerals indicating that charge.
Consider FeO and Fe2O3. These are very different compounds with different properties.
When we name these compounds, it is absolutely vital that we clearly distinguish between them.
80 They are both iron oxides but in FeO iron is exhibiting a charge of 2+ and in
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Fe2O3, it is exhibiting a charge of 3+. The first, FeO, is named iron (II) oxide. The second,
Fe2O3, is named iron (III) oxide.

Lesson Summary octet rule: the tendency for atoms gain or lose the appropriate number of electrons so
that the resulting ion has either completely filled or completely empty outer energy
When an atom gains one or more extra electrons, it becomes a negative ion, an anion
levels, or 8 valance electrons.
When an atom loses one or more of its electrons, it becomes a positive ion, a cation.
Polyatomic ions are ions composed of a group of atoms that are covalently bonded Further Reading / Supplemental Links
and behave as if they were a single ion.
Some transition elements have fixed oxidation numbers and some have Website with lessons, worksheets, and quizzes on various high school chemistry
variable charges. When naming these charge variable ions, their charges are topics. Lesson 4-1 is on Electronegativity. Lesson 4-2 is on Types of Bonds.
included in Roman numerals. http://www.fordhamprep.org/gcurran/sho/sho/lessons/lesson31.htm

4.3: Review Questions

1) Define an ion.
2) Will an iron atom form a positive or negative ion? Why?
3) Will a bromine atom form a positive or negative ion? Why?

Predict the charge of each ion. Then give the name each ion would have.
4) Cl 5) Br 6) N

7) O 8) Ca 9) F

10) Mg 11) Li 12) I

13) Na 14) K 15) Al

16) How are transition metals that form ions named differently than other metals? Why is
this important? What does the Roman numeral tell you?

Name the following ions.

2+ 2+ 3+
17) Cu 18) Co 19) Co
+ 2+ 3+
20) Cu 21) Ni 22) Cr
2+ 3+
23) Fe 24) Fe

25) What are polyatomic ions?

CC Tracy Poulsen Name each of the following ions.
-
26) NO3 - 27) C2H3O2 -
28) OH
Vocabulary
29) PO4 30) SO3 31) CO32-
Ion: An atom or group of atoms with an excess positive or negative charge. 3- 2-

Cation: positive ion

Anion: negative ion
ionic bond: A bond between ions resulting from the transfer of electrons from one
of the bonding atoms to the other and the resulting electrostatic attraction between
the ions.
electrostatic attraction: The force of attraction between opposite electric charges.

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4.4: Writing Ionic Formulas
Objectives
Write the correct formula for an ionic compound Solution:
3+
Cation, aluminum: Al (you can find this charge using your periodic table)
Introduction 2-
Anion, sulfide: S (sulfide is sulfur as an ion, get its charge from your periodic table)
Ionic compounds do not exist as molecules. In the solid state, ionic compounds are in
crystal lattices containing many ions each of the cation and anion. An ionic formula, like To balance the charges you need 2·(+3) and 3·(-2). Giving:
NaCl, is an empirical formula. This formula merely indicates that sodium chloride is made of The final formula is: Al2S3
an equal number of sodium and chloride ions. Sodium sulfide, another ionic compound, has
Example: Write the formula for lead (IV) oxide.
the formula Na2S. This formula indicates that this compound is made up of twice as many
sodium ions as sulfide ions. This section will teach you how to find the correct ratio of ions, Solution:
4+
so that you can write a correct formula. Cation, lead (IV): Pb (the charge is given to you as Roman numerals, because this is a
metal with a variable charge)
2-
Ionic Formulas Anion, oxide: O (oxide is oxygen as an ion, get its charge from your periodic table)
When an ionic compound forms, the number of To balance the charges you need 1·(+4) and 2·(-2). Giving:
Carbon will share electrons with
electrons given off by the cations must be exactly the same hydrogen, or other atoms, to get an The final formula is: PbO2
as the number of electrons taken on by the anions. octet.
Therefore, if calcium, which gives off two electrons, is to combine with fluorine, which Example: Write the formula for calcium nitrate.
takes on one electron, then one calcium atom must combine with two fluorine atoms. The Solution:
2+
formula would be CaF2. Cation, calcium: Ca (you can find this charge using your periodic table)
-
To write the formula for an ionic compound: Anion, nitrate: (NO3) (this is a polyatomic ion)
To balance the charges you need 1·(+2) and 2·(-1). Giving:
1) Write the symbol and charge of the cation (first word) The final formula is: Ca(NO3)2
a) If the element is in group 1, 2, Al with a consistent charge, you can get the -
In this case you need to keep the parentheses. There are two of the group (NO3) . Without
charge using your periodic table. the parentheses, you are merely changing the number of oxygen atoms.
b) If the metal is a transition metal with a variable charge, the charge will be given
to you in Roman numerals. Example: Write the formula for magnesium sulfate.
2) Write the symbol and charge of the anion (second word). Solution:
a) Look at your polyatomic ion chart first. If your anion is a polyatomic ion, write the 2+
Cation, magnesium: Mg (you can find this charge using your periodic table)
ion in parentheses. 2-
Anion, sulfate: (SO4) (this is a polyatomic ion)
b) If the anion is not on the polyatomic chart, it is a nonmetal anion from your To balance the charges you need 1·(+2) and 1·(-2). Giving:
periodic table. You can get its charge using your table. The final formula is: MgSO4
3) Write the correct subscripts so that the total charge of the compound will be zero. In this case you do not need parentheses. They are only required if there is more than one of
4) Write the final formula. Leave out all charges and all subscripts that are 1. If there is the polyatomic ion.
only 1 of the polyatomic ion, leave off parentheses.
Example: Write the formula for copper (II) acetate.
Pay close attention to how these steps are followed in the given examples. Solution:
2+
Cation, copper (II): Cu (the charge is given to you in Roman numerals)
-
Anion, acetate: (C2H3O2) (this is a polyatomic ion)
Example: Write the formula for aluminum chloride.
To balance the charges you need 1·(+2) and 2·(-1). Giving:
Solution:
3+ The final formula is: Cu(C2H3O2)2
Cation, aluminum: Al (you can find this charge using your periodic table) -
-
Anion, chloride: Cl (chloride is chlorine as an ion, get its charge from your periodic table) In this case you need to keep the parentheses. There are two of the group (C2H3O2) .
Without the parentheses, you are merely changing the number of oxygen atoms.
To balance the charges you need 1·(+3) and 3·(-1). Giving:
The final formula is: AlCl3 Lesson Summary
Formulas for ionic compounds contain the lowest whole number ratio of subscripts
Example: Write the formula for aluminum sulfide.
such that the sum of the subscript of the more electropositive element times its
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 oxidation number plus the subscripts of the more electronegative element times its
oxidation number equals zero.
Introduction:
Vocabulary
We have already learned about naming individual ions, including main group ions,
Ionic Formula: includes the symbols and number of each atom present in a compound transition metal ions, and polyatomic ions. We have also learned how to put these into
in the lowest whole number ratio correct charge-balanced formulas. In this section, we will learn how to correctly naming a
compound, given its formula.
Further Reading / Supplemental Links
http://www.kanescience.com/_chemistry/5Ionic.htm Naming Ionic Compounds
http://visionlearning/library/module_viewer.php?mid=55 To name ionic compounds, we will need to follow these steps:
1. Split the formula into the cation and anion. The first metal listed will be the cation
4.4: Review Questions and the remaining element(s) will form the anion.
Copy and fill in the chart by writing formulas for the compounds that might form between 2. Name the cation. We learned two types of cations:
the ions in the columns and rows. Some of these compounds don’t exist but you can still a. Main group cations in which the name of the ion is the same as that of
write formulas for them. +
the element (for example, K is potassium).
+ 2+ 3+
Na Ca Fe b. Transition metals with variable charges with Roman numerals indicating the
NO3 - 1) 2) 3) When
charge of the ion (you will havean to
ionic
dobond forms,
a little bit electrons
of mathare lost from
to find thisacharge).
metal
and given to a nonmetal.
SO4 2- 4) 5) 6) 3. Name the anion. There are also two general types of anions:
Cl
- a. Main group anions in which the name of the anion ends in “-ide” (for
7) 8) 9) -
PO4 3- example, F is fluoride)
10) 11) 12)
- b. Polyatomic ions (as listed on the polyatomic ion chart)
OH 13) 14) 15)
CO3 2- 16) 17) 18) When writing the name of an ionic compound, it is important to note that the name
gives no information about the number of ions. The name only tells the types of ions present.
Write the formulas from the names of the following compounds. The formula uses subscripts to indicate how many of each ion there are.
19) Magnesium sulfide 20) Lead(II) Nitrate 21) Sodium Oxide
Example: What is the name of Na2O?
22) Calcium hydroxide 23) Potassium Carbonate 24) Aluminum Bromide Solution:
An ionic bond is similar to a tug-of-war in which the stronger
Split up the formula: Na2 | O (more electronegative) nonmetal atom pulls the electron away
25) Iron (III) nitrate 26) Iron(II) Chloride 27) Copper(II) Nitrate from the
Name the cation: Na is a group 1 metal with weaker (less electronegative)
a consistent metal not
charge. It does atom.need Roman
CC – Tracy Poulsen numerals. Its name is “sodium”
28) Magnesium oxide 29) Calcium Oxide 30) Copper(I) Bromide Name the anion: O is not polyatomic. When oxygen atoms get a -2 charge, the name
changes to end in –ide, so the anion is “oxide” Final answer: sodium oxide
31) Aluminum sulfide 32) Hydrogen Carbonate 33) Potassium
permanganate
34) Copper (I) dichromate 35) Iron(III) Chloride 36) Iron(II) Sulfate Example: What is the name of NaC2H3O2?
Solution:
Split up the formula: Na | C2H3O2
Name the cation: Na is a group 1 metal with a consistent charge. It does not need Roman
numerals. Its name is “sodium”
4.5: Naming Ionic Compounds Name the anion: C2H3O2 is polyatomic. Its name is “acetate”.
Objectives Final answer: sodium acetate
Correctly name binary ionic compounds, compounds containing metals with variable
oxidation numbers, and compounds containing polyatomic ions given the formulas. Example: Write the name of CuCl2?
Solution:
Split up the formula: Cu | Cl2
Name the cation: Cu is a transition metal with a variable charge. It needs Roman numerals.
86 To find the charge, consider the charge of the other ion and the number of both ions:
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 . The copper must have a charge of +2 to balance out the negatives: 1·(+2)
to cancel out 2·(-1). Its name is “copper (II)”
Name the anion: Cl is not polyatomic. When chlorine atoms get a -1 charge, the name Remember, when writing the name of an ionic compound, it is important to note
changes to end in –ide, so the anion is “chloride” Final answer: copper (II) chloride that the name gives no information about the number of ions. The name only tells the types
of ions present. The formula uses subscripts to indicate how many of each ion there are.
Example: Write the name of PbS2? Lesson Summary
Solution: Ionic bonds are formed by transferring electrons from metals to non-metals
Split up the formula: Pb | S2 after which the oppositely charged ions are attracted to each other.
Name the cation: Pb is a post-transition metal with a variable charge. It needs Roman Ionic compounds form crystal lattice structures rather than molecules.
numerals. To find the charge, consider the charge of the other ion and the number of both Binary ionic compounds are named by naming the metal first followed by the non-
ions: . The copper must have a charge of +4 to balance out the negatives: 1·(+4) metal with the ending of the non-metal changed to “ide.”
to cancel out 2·(-2). Its name is “lead (IV)” Compounds containing polyatomic ions are named with the name of the polyatomic
Name the anion: S is not polyatomic. When sulfur atoms get a -2 charge, the name changes ion in the place of the metal or non-metal or both with no changes in the name of the
to end in –ide, so the anion is “sulfide” polyatomic ion.
Final answer: Lead (IV) sulfide. Compounds containing variable oxidation number metals are named with Roman
numerals in parentheses following the name of the metal and indicating the
The most common error made by students in naming these compounds is to choose oxidation number of the metal.
the Roman numeral based on the number of atoms of the metal instead of the charge of the Vocabulary
metal. For example, in PbS2, the oxidation state of lead Pb is +4 so the Roman numeral Anion: An ion with a negative charge.
following the name lead is “IV.” Notice that there is no four in the formula. As in previous Cation: An ion with a positive charge.
examples, the formula is always the lowest whole number ratio of the ions involved. Think Chemical nomenclature: The system for naming chemical compounds.
carefully when you encounter variable charge metals. Make note that the Roman numeral Ionic bond: The electrostatic attraction between ions of opposite charge.
does not appear in the formula but does appear in the name. Polyatomic ion: A group of atoms bonded to each other covalently but possessing an
overall charge.
Example: Write the name of Mg3(PO4)2 ?
Solution: Further Reading / Supplemental Links
Split up the formula: Mg3 | (PO4)2 Matching Game: Naming Ionic Compounds with Polyatomic Ions:
Name the cation: Mg is a group 2 metal with a consistent charge. It does not need Roman http://www.quia.com/mc/65767.html
numerals. Its name is “magnesium”
Name the anion: PO4 is polyatomic. Its name is “phosphate”. 4.5: Review Questions
Final answer: sodium acetate Name the following compounds.
1) KCl 2) MgO 3) CuSO4
Example: Write the name of Cr(NO2)3 ?
Solution: 4) NaCl 5) CoBr2 6) MgF2
Split up the formula: Cr | (NO2)3
Name the cation: Cr is a transition metal with a variable charge. It needs Roman numerals. 7) Ni(OH)2 8) NaC2H3O2 9) CuO
To find the charge, consider the charge of the other ion and the number of both ions:
. The copper must have a charge of +3 to balance out the negatives: 1·(+3) 10) FeCl2 11) LiCl 12) MgBr2
to cancel out 3·(-1). Its name is “chromium (III)”
-
Name the anion: NO2 is polyatomic. Its name is “nitrite”. 13) NH4(OH) 14) Cu2O 15) CaF2
Final answer: chromium (III) nitrite
16) K2CO3 17) Na2O 18) PbO

19) Ca(NO3)2 20) Mg(OH)2 21) SnO2

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4.6: Covalent Compounds & Lewis Structures
Objectives
Explain what covalent bonds are.
Explain why covalent bonds are formed. Lewis Structures
Draw a Lewis structures for covalent compounds and polyatomic ions The Lewis structure of a molecule show how the valence electrons are arranged
among the atoms of the molecule. These representations are named after G. N. Lewis. The
Introduction
rules for writing Lewis structures are based on observations of thousands of molecules.
In ionic bonding, electrons leave metallic atoms and enter non-metallic atoms. This From experiment, chemists have learned that when a stable compound forms, the atoms
complete transfer of electrons changes both of the atoms into ions. Often, however, two usually have a noble gas electron configuration or eight valence electrons. Hydrogen forms
atoms combine in a way that no complete transfer of electrons occurs. Instead, electrons are stable molecules when it shares two electrons (sometimes called the duet rule). Other atoms
held in overlapping orbitals of the two atoms, so that the atoms are sharing the electrons. involved in covalent bonding typically obey the octet rule. (Note: Of course, there will be
The shared electrons occupy the valence orbitals of both atoms at the same time. The nuclei exceptions.)
of both atoms are attracted to this shared pair of electrons and the atoms are held together by
this attractive force. The attractive force produced by sharing electrons is called a covalent To draw a Lewis structure:
bond. 1. Determine the number of valence electrons that will be drawn in the Lewis structure.
a. Use your periodic table to determine the number of valence electrons in
Covalent Bond Formation
each atom. Add these to get the total electrons in the structure.
In covalent bonding, the atoms
b. If you are drawing the structure for a polyatomic ion, you must add or subtract
acquire a stable octet of electrons by
any electrons gained or lost. If an ion has a negative charge, electrons were
sharing electrons. The covalent
gained. If the ion has a positive charge, electrons were lost.
bonding process produces molecular
2. Draw a skeleton
substances as opposed to the lattice
a. Typically, the first element listed in the formula goes in the center, which
structures of ionic bonding. There are In a covalent bond, electrons are shared by atoms in
the remaining atoms surrounding.
far more covalently bonded overlapping orbitals.
b. Draw bonds to each of the surrounding atoms. Each bond is two valence
substances than ionic substances. electrons.
The diatomic hydrogen
3. Use the remaining electrons to give each atom an octet (except hydrogen which
molecule, H2, is one of the many
only gets a duet)
molecules that are covalently
a. Place electrons left over after forming the bonds in the skeleton in
bonded. Each hydrogen atom has a
unshared pairs around the atoms to give each an octet. *Remember, any
1s electron cloud containing one We can also show a covalent bond between atoms with an bonds they have formed already count as two valence electrons each.
electron. These 1s electron clouds electron dot formula where the shared pair of electrons are the
b. If you run out of electrons, and there are still atoms without an octet,
overlap and produce a common bonding electrons or with the bond represented by a dash.
move some of the electrons that are not being shared to form double,
volume which the two electrons
sometimes triple bonds.
occupy.
Example: Draw a Lewis structure for water, H2O.
Some Compounds Have Both Covalent and Ionic Bonds
If you recall the introduction of polyatomic ions, Solution:
you will remember that the bonds that hold the polyatomic 1) add up all available valence electrons: each H atom has 1, each oxygen atom has 6, so
ions together are covalent bonds. Once the polyatomic ion 2(1)+6=8
is constructed with covalent bonds, it reacts with other
2) Draw a skeleton. Although the first atom written typically goes in the middle, hydrogen
substances as an ion. The bond between a polyatomic ion
can’t, so O gets the middle spot. We need to draw bonds connecting atoms in the skeleton.
and another ion will be ionic. An example of this type of In this compound, the N and O
situation is in the compound sodium nitrate. Sodium nitrate atoms are covalently bonded,
We get:
is composed of a sodium ion and a nitrate ion. The nitrate sharing electrons. However, the
ion is held together by covalent bonds and the nitrate ion is NO3 anion is ionically bonded to 3) Use the remaining electrons to give each atom (except hydrogen) an octet. If we look at
attached to the sodium ion by an ionic bond.
the Na cation. our skeleton, we drew two bonds, which uses 4 of our 8 available electrons. We are left with
four more. Each H atom already has two valence electrons and O currently has 4 (each bond
counts as two for each atom that it connects). We will give the remaining four electrons to
90 O, in pairs. We get:
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 or
Check:
Is the total number of valence electrons correct? Yes. Our final picture has 8 valence e-.
16 electrons, we may get a picture such as:
Does each atom have the appropriate duet or octet of electrons? Yes

Example: Draw a Lewis structure for CO2

Solution:
1) add up all available valence But notice that the nitrogen atom still does not have an octet. We ran out of electrons so we
electrons: 1(4) + 2(6) = 16 must form a double bond. Use some of the electrons on an oxygen atom to share with the
2) Draw a skeleton. nitrogen. We get:
Carbon goes in the middle with the two oxygen atoms bonded to it:

3) Use the remaining electrons to give each atom (except hydrogen) an octet.
In this case, we have already used up four electrons to draw the two bonds in the skeleton, Check:
leaving 12 left. This is not enough to give everybody an octet. Our picture may look Is the total number of valence electrons correct? Yes. Our final picture has 24 valence e-.
something like this with 16 electrons: Does each atom have the appropriate duet or octet of electrons? Yes

Lesson Summary
Covalent bonds are formed by electrons being shared between two atoms. Half-
We have used up the 16 electrons, but neither O has an octet. The rules state that if you run filled orbitals of two atoms are overlapped and the valence electrons shared by the
out of electrons and still don’t have octets, then you must use some of the unshared pairs of atoms.
electrons as double or triple bonds instead. Move the electrons that are just on the carbon Covalent bonds are formed between atoms with relatively high electron affinity.
atom to share with the oxygen atom until everybody has an octet. We get:
Vocabulary
Covalent bond: A type of bond in which electrons are shared by atoms.
OR
Check: Further Reading / Supplemental Links
Is the total number of valence electrons correct? Yes. Our final picture has 16 valence e-. Tutorial on bonding:
Does each atom have the appropriate duet or octet of electrons? Yes http://visionlearning.org/library/module_viewer.php?mid=55&l=

Example: Draw a Lewis structure for nitric acid, HNO3. The skeleton is given below: 4.6: Review Questions
Which of the following compounds would you expect to be ionically bonded and
which covalently bonded?
1) CS2 4) PF3
Solution: 2) K2S 5) AlF3
1) add up all available valence 3) FeF3 6) BaS
electrons: 1(1) + 1(5) + 3(6)=24
2) Draw a skeleton. Draw a Lewis structure for each of the following compounds.
This was given to us, but we need to draw the bonds. 7) H2O 12) CO3 2-
8) CH4 13) CO2
9) CO 14) NH3
10) PCl3 15) CH2O
3) Use the remaining electrons to give each atom (except hydrogen) an octet. 11) C2H6 16) SO3
Each bond used up 2 electrons, so we have already used 8 electrons. If we use the reaining
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4.7: Molecular Geometry
Objectives
Predict the shape of simple molecules and their polarity from Lewis dot structures. We have a similar problem in the case of a molecule such as water, H2O. In water,
Explain the meaning of the acronym VSEPR and state the concept on which it is the oxygen atom in the middle is bonded to the two hydrogen atoms with two lone pairs.
based. Once again, we only consider the location of atoms when we discuss shape. When a
molecule has a central atom bonded to two other atoms with two lone pair of electrons,
Introduction
the overall shape is bent.
Although a convenient way for chemists to look at covalent compounds is to draw
Lewis structures, which shows the location of all of the valence electrons in a compound.
Although these are very useful for understanding how atoms are arranged and bonded, they
are limited in their ability to accurately represent what shape molecules are. Lewis structures
are drawn on flat paper as two dimensional drawings. However, molecules are really three
dimensional. In this section you will learn to predict the 3d shape of many molecules given
their Lewis structure.
Many accurate methods now exist for determining molecular structure, the three-
dimensional arrangement of the atoms in a molecule. These methods must be used if precise Various molecular geometries with four pairs of electrons around a central atom.
information about structure is needed. However, it is often useful to be able to predict the
approximate molecular structure of a molecule. A simple model that allows us to do this is As you can probably imagine, there are different combinations of bonds making
called the valence shell electron pair repulsion (VSEPR) theory. This model is useful in different shapes of molecules. Some of the possible shapes are listed in the table. However,
predicting the geometries of molecules formed in the covalent bonding of non-metals. The it is important to note that some molecules obtain geometries that are not included here.
main postulate of this theory is that in order to minimize electron-pair repulsion. In other Summary of Molecular Geometry
words, the electron pairs around the central atom in a molecule will get as far away from
# of atoms # of unshared
each other as possible.
bonded to pairs around Molecular Geometry
central atom central atom
2 0 Linear
3 0 Trigonal Planar
Predicting the Shape of Molecules
2 1 Bent
Consider, methane,
commonly known as natural gas. 3 1 *Trigonal pyramidal
In this molecule, carbon has four 2 2 *Bent
valence electrons and each 4 0 *Tetrahedral
hydrogen adds one more so the +
central atom in methane has four Example: Determine the shape of ammonium, NH4 , given by the following Lewis structure:
pairs of electrons in its valence
shell. The 3d shape of this
molecule is dictated by the A central atoms bonded to four other atoms with no lone pairs
repulsion of the electrons. Those has a tetrahedral shape.
four pairs of electrons get as far
away from each other as possible which forms a shape called tetrahedral. In the Solution: To answer this question, you need to count the number of atoms
tetrahedral shape, the bond angle between any two hydrogen atoms is 109.5°. around the central atom and the number of unshared pairs. In this example,
What if we look at ammonia instead, NH3? A molecule of ammonia has a nitrogen there are four atoms bonded to the N with zero unshared pairs of electrons.
atom in the middle with three bonds to the hydrogen atoms plus one lone pair of electrons. The shape must be tetrahedral.
That means there are four total pairs of electrons around the central atom, and the
electrons will still be close to 109.5° apart from each other. However, when discussing the
overall shape of the molecule, we only take into account the location of the atoms. When a Example: Determine the shape of carbon dioxide, CO2, given by the following
central atom is bonded to three atoms and has one lone pair of electrons, the overall shape Lewis structure:
is trigonal pyramidal.

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Solution: To answer this question, you need to count the number of atoms around the central
atom and the number of unshared pairs. In this example, there are two atoms bonded to the C
with zero unshared pairs of electrons. The shape must be linear, according to the table. 4) Water, H2O H H
.. | |
H—O—H H—N: H—C—H
`` | |
H H
Example: Determine the shape of carbon dioxide, SO2, given by the following
Lewis structure: 4.8: Polarity & Hydrogen Bonding
Objectives
Explain how polar compounds differ from nonpolar compounds
Determine if a molecule is polar or nonpolar
Identify whether or not a molecule can exhibit hydrogen bonding
Solution: To answer this question, you need to count the number of atoms List important phenomena which are a result of hydrogen bonding
around the central atom and the number of unshared pairs. In this example, Given a pair of compounds, predict which would have a higher melting or boiling
there are two atoms bonded to the S with one unshared pair of electrons. point
The shape must be bent, according to the table.
Introduction
The ability of an atom in a molecule to attract shared electrons is called
Vocabulary electronegativity. When two atoms combine, the difference between their
VSEPR model: A model whose main postulate is that the structure around a given electronegativities is an indication of the type of bond that will form. If the difference
atom in a molecule is determined by minimizing electron-pair repulsion. between the electronegativities of the two atoms is small, neither atom can take the shared
Molecular geometry: The specific three-dimensional arrangement of atoms in electrons completely away from the other atom and the bond will be covalent. If the
molecules. difference between the electronegativities is large, the more electronegative atom will take
the bonding electrons completely away from the other atom (electron transfer will occur)
Further Readings / Supplemental Links and the bond will be ionic. This is why metals (low electronegativities) bonded with
nonmetals (high electronegativities) typically produce ionic compounds.
http://www.up.ac.za/academic/chem/mol_geom/mol_geometry.htm
http://en.wikipedia.org/wiki/Molecular_geometry
Polar Covalent Bonds
An animation showing the molecular shapes that are generated by sharing various
numbers of electron pairs around the central atom (includes shapes when some pairs So far, we have discussed two extreme types of bonds. One case is when two
of electrons are non-shared pairs). The link must be copied and pasted into your identical atoms bond. They have exactly the same electronegativities, thus the two bonded
browser to go directly to the animation. atoms pull exactly equally on the shared electrons. The shared electrons will be shared
http://www.classzone.com/cz/books/woc_07/resources/htmls/ani_chem/chem_flash/p exactly equally by the two
opup.html?layer=act&src=qtiwf_act065.1.xml
atoms.
The other case is when
the bonded atoms have a very
4.7: Review Questions large difference in their
Predict the 3d shape each of the following molecules will have: electronegativities. In this case,
1) CH3Cl 2) Silicon tetrafluoride, SiF4 3) CHCl3 the more electronegative atom
H F Cl will take the electrons
| | | completely away from the other
Cl—C—H F—Si—F H—C—Cl atom and an ionic bond forms. A polar covalent bond is similar to a tug-of-war in which one
| | | What about the atom pulls more on the electrons and gains a partial negative
molecules whose charge. The weaker (less electronegative atom) has a partial
H F Cl
electronegativities are not the positive charge.
5) Ammonia, NH3 6) Methane, CH4 same but the difference is not CC – Tracy Poulsen
big enough to form an ionic
96 bond? For these molecules, the electrons remain shared by the two atoms but they are not
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shared equally. The shared electrons are pulled closer to the more electronegative atom. This
results in an uneven distribution of electrons over the molecule and causes slight charges on
opposite ends of the molecule. The negative electrons are around the more electronegative
atom more of the time creating a partial negative
side. The other side has a resulting partial
positive charge. These charges are not full +1 2) methanol, CH3OH:
and -1 charges, they are fractions of charges. For
small fractions of charges, we use the symbols 3) hydrogen cyanide, HCN:
δ+ and δ−. These molecules have slight opposite
charges on opposite ends of the molecule and 4) Oxygen, O2:
said to have a dipole or are called polar
molecules.
When atoms combine, there are three possible types of bonds that they can form. 5) Propane, C3H8:
In the figure, molecule A represents a covalent bond that would be formed between Solution:
identical atoms. The electrons would be evenly shared with no partial charges forming. 1) Water is polar. Any molecule with lone pairs of electrons around the central atom is polar.
This molecule is nonpolar. Molecule B is a polar covalent bond formed between atoms 2) Methanol is polar. This is not a symmetric molecule. The –OH side is different from the
whose electronegativities are not the same but whose electronegativity difference is less other 3 –H sides.
than 1.7, making this molecule polar. Molecule C is an ionic bond formed between atoms 3) Hydrogen cyanide is polar. The molecule is not symmetric. The nitrogen and hydrogen
whose electronegativity difference is greater than 1.7. have different electronegativities, creating an uneven pull on the electrons.
4) Oxygen is nonpolar. The molecule is symmetric. The two oxygen atoms pull on the
electrons by exactly the same amount.
5) Propane is nonpolar, because it is symmetric, with H atoms bonded to every side around
the central atoms and no unshared pairs of electrons.

While molecules can be described as "polar covalent", "non-polar covalent",

A) A nonpolar covalent bond in which two identical atoms are or "ionic", it must be noted that this is often a relative term, with one molecule simply
sharing electrons
being more polar or less polar than another. However, the following properties are typical of
B) a polar covalent bond in which the more electronegative atom pulls
the electrons more toward itself (forming partial negative and positive such molecules. Polar molecules tend to:
sides) have higher melting points than nonpolar molecules
C) an ionic bond in which an extremely electronegative atom is bonded have higher boiling points than nonpolar molecules
to a very weakly electronegative atom. be more soluble in water (dissolve better) than nonpolar
Polar molecules can be attracted to each other due attraction between opposite molecules have lower vapor pressures than nonpolar molecules
charges. Polarity underlies a number of physical properties including surface tension,
solubility, and melting- and boiling-points. The more attracted molecules are to other Hydrogen Bonding:
molecules, the higher the melting point, boiling point, and surface tension. We will discuss When a hydrogen atom is bonded to a very electronegative atom, including
in more detail later how polarity can affect how compounds dissolve and their solubility. fluorine, oxygen, or nitrogen, a very polar bond is formed. The electronegative atom
In order to determine if a molecule is polar or nonpolar, it is frequently useful to look obtains a negative partial charge and the hydrogen obtains a positive partial charge. These
a Lewis structures. Nonpolar compounds will be symmetric, meaning all of the sides around partial charges are similar to what happens in every polar molecule. However, because of
the central atom are identical – bonded to the same element with no unshared pairs of the big difference in electronegativities between these two atoms and the amount of
electrons. Polar molecules are assymetric, either containing lone pairs of electrons on a positive charge exposed by the hydrogen, the dipole is much more dramatic. These
central atom or having atoms with different electronegativities bonded. molecules will be attracted to other molecules which also have partial charges. This
Example: Label each of the following as polar or nonpolar. attraction for other molecules which also have a hydrogen bonded to a fluorine, nitrogen, or
oxygen atom is called a hydrogen bond.
1) Water, H2O:
Hydrogen bonds in water
The most important, most common, and perhaps simplest example of a hydrogen
bond is found between water molecules. This interaction between neighboring water
molecules is responsible for many of the important properties of water.
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 Hydrogen bonding strongly affects the crystal structure of ice, helping to create an
open hexagonal lattice. The density of ice is less than water at the same temperature; thus, the
solid phase of water floats on the liquid, unlike most other substances in which the solid form
would sink in the liquid form. have higher boiling points than polar molecules
Water also has a high boiling point (100°C) compared to the other compounds be more soluble in water (dissolve better) than polar molecules
of similar size without hydrogen bonds. Because of the difficulty of breaking these Example: Label each of the following as polar or nonpolar and indicate which have
bonds, water has a very high boiling point, melting point, and viscosity compared to hydrogen bonding.
otherwise similar liquids not conjoined by hydrogen bonds.
Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms,
meaning that the total number of bonds of a water molecule is up to four. For example, a) H2O,
hydrogen fluoride——which has three lone pairs on the F atom but only one H atom——
can form only two bonds; (ammonia has the opposite problem: three hydrogen atoms but
only one lone pair).
... ...

H-F H-F H-F.

Have you ever experienced a belly flop? This is also due to the hydrogen bonding b) Ammonia,
between water molecules, causing surface tension. On the surface of water, water molecules
are even more attracted to their neighbors than in the rest of the water. This attraction makes
it difficult to break through, causing belly flops. It also explains why water striders are able
to stay on top of water and why water droplets form on leaves or as they drip out of your
faucet.
c) CH4,

d) acetone, CH3COCH3
Solution:
A. Water beading on a leaf B. Water dripping from a tap C. Water striders stay atop the liquid due to surface tension a) This molecule is polar (the unshared pairs of electrons make a polar assymetric
shape), and hydrogen bonding (hydrogen is bonded to N, O, or F).
b) This molecule is polar (the unshared pairs of electrons make a polar assymetric
Hydrogen bonds in DNA and proteins Hydrogen shape), and hydrogen bonding (hydrogen is bonded to N, O, or F).
bonding also plays an important c) This molecule is nonpolar (the molecule is symmetric with H’s bonded to all four sides of
role in determining the three-dimensional structures the central atom), and does not have hydrogen bonding (hydrogen is not bonded to N, O,
adopted by proteins and nucleic bases, as found in or F).
your DNA. In these large molecules, bonding between d) This molecule is polar (the O is not the same as the CH3 bonded to the central atom)
parts of the same macromolecule cause it to fold into a and does not have hydrogen bonding (hydrogen is bonded DIRECTLY to N, O, or F).
specific shape, which helps determine the molecule's
physiological or biochemical role. The double helical Example: For each pair of molecules, indicate which you would expect to have a higher
structure of DNA, for example, is due largely to melting point. Explain why. Also, refer to the Lewis structures given to you in the previous
hydrogen bonding between the base pairs, which link example.
one complementary strand to the other and enable a) H2O vs. acetone
replication. It also plays an important b) CH4 vs. acetone
role in the structure of polymers, both synthetic and natural, such as nylon and many plastics. Solution:
As a result of the strong attraction between molecules that occurs in a hydrogen bond, a) H2O (polar, hydrogen bonding) vs. acetone (polar, no hydrogen bonding). H2O will have a
the following properties can be summarized. Molecules with hydrogen bonding tend to: higher melting point because compounds with hydrogen bonding tend to have higher
have higher melting points than polar molecules melting points than polar compounds.
b) CH4 (nonpolar, no hydrogen bonding) vs. acetone (polar, no hydrogen bonding). Acetone
100 will have a higher melting point because polar molecules tend to have higher melting
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Lesson Summary
Covalent bonds between atoms that are not identical will produce polar bonds. | | |
Molecules with polar bonds and non-symmetrical shapes will have a dipole. H F Cl
Hydrogen bonding is a special interaction felt between molecules, which is a stronger
interaction than polar-polar attraction.
Hydrogen bonding occurs between molecules in which a hydrogen atom is bonded 10) Water, H2O 11) Ammonia, NH3 12) Methane, CH4
to a very electronegative fluorine, oxygen, or nitrogen atom. .. H H
Compounds with hydrogen bonding tend to have higher melting points, higher boiling H—O—H | |
points, and greater surface tenstion. `` H—N: H—C—H
The unique properties of water are a result of hydrogen bonding | |
Hydrogen bonding plays roles in many compounds including DNA, proteins, and H H
polymers.
For each of the following, indicate which of the compounds in the pair has the
Vocabulary given property.
Electronegativity: The tendency of an atom in a molecule to attract shared electrons 13) higher melting point: ammonia or methane
to itself. 14) higher boiling point: water or CH3Cl
15) more soluble in water: ammonia or CHCl3
Polar covalent bond: A covalent bond in which the electrons are not shared equally 16) higher melting point: SiF4 or ammonia
because one atom attracts them more strongly that the other.

Further Reading / Supplemental Links All images, unless otherwise stated, are created by the CK-12 Foundation and are under
http://learner.org/resources/series61.html; The learner.org website allows users to the Creative Commons license CC-BY-NC-SA.
view streaming videos of the Annenberg series of chemistry videos. You are
required to register before you can watch the videos but there is no charge. The
website has one video that relates to this lesson called Molecular Architecture.
Vision Learning: Water Properties & Behaviors
http://visionlearning.com/library/module_viewer.php?mid=57&l=&c3=

4.8: Review Questions

1) Explain the differences among a nonpolar covalent bond, a polar covalent bond, and an
ionic bond.
2) Predict which of the following bonds will be more polar and explain why; P-Cl or S-Cl.
3) What does it mean for a molecule to be “polar”?
4) Which three elements, when bonded with hydrogen, are capable of forming hydrogen
bonds?
5) Molecules that are polar exhibit dipole-dipole interaction. What’s the difference between
dipole-dipole interactions and hydrogen bonding? Which interaction is stronger?
6) Define hydrogen bonding. Sketch a picture of several water molecules and how they
interact.
Given each of the following Lewis structures, indicate whether each is polar or nonpolar.
Then indicate whether or not that compound exhibits hydrogen bonding.
7) CH3Cl 8) Silicon tetrafluoride, SiF4 9) CHCl3
H F Cl
| | |
Cl—C—H F—Si—F H—C—Cl

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 Chapter 5: Problem Solving & the Mole
5.1: Measurement Systems centi- is placed in front of gram, as in centigram, the measure is now of a gram. When
Objectives
State the different measurement systems used in chemistry. milli- is placed in front of meter, as in millimeter, the measure is now of a meter.
Describe how prefixes are used in the metric system and identify how the Common prefixes are shown in the table.
prefixes milli-, centi-, and kilo- compare to the base unit Explain the difference Common Prefixes in the International System
between mass and weight. Prefix Meaning Symbol
Identify SI units of mass, distance (length), volume, temperature, and time. -6
micro- 10 μ
Describe absolute zero. milli- m
10-3
-2
Introduction centi- 10 c
3
Even in ancient times, humans needed measurement systems for commerce. Land kilo- 10 k
ownership required measurements of length and the sale of food and other commodities
required measurements of mass. Mankind’s first elementary efforts in measurement required These prefixes are used for all metric units of measurement, including units for
convenient objects to be used as standards and the human body was certainly convenient. volume, time, distance, etc. Common metric units, symbols, and relationships to a base unit
The names of several measurement units reflect these early efforts. Inch and foot are are shown below.
-6
examples of measurement units that are based on parts of the human body. The inch is based 1 micrometer = 1 μm = 1x10 m
-6
on the width of a man’s thumb, and the foot speaks for itself. 1 microliter = 1 μL = 1x10 L
3
It should be apparent that measurements of a foot by two people could differ by a few 1 kilometer = 1 km = 1x10 m
3
inches. To achieve more consistency, everyone could use the king’s foot as the standard. The 1 kilogram = 1 kg = 1x10 g
length of the king’s foot could be marked on pieces of wood and everyone who needed to
measure length could have a copy. Of course, this standard would change when a new king SI Units
was crowned. What was needed were objects that could be safely stored so they didn’t The International System of Units, abbreviated SI from the French Le Système
change over time. Copies could be made of these objects and distributed so that everyone International d’ Unites, is the main system of measurement units used in science. Since the
was using exactly the same units of measure. The requirements of science in the 1600s, 1960s, the International System of Units has been internationally agreed upon as the standard
1700s, and 1800s necessitated even more accurate, reproducible measurements. metric system. The SI base units are based on physical standards. The definitions of the SI
base units have been and continue to be modified and new base units added as advancements
The Metric System in science are made. Each SI base unit except the kilogram is described by stable properties
The metric system is an international decimal-based system of measurement. of the universe.
Because the metric system is a decimal system, making conversions between different units
of the metric system are always done with factors of ten. Let’s consider the English system – Mass
that is, the one that is in everyday use in the US– to explain why the metric system is so Mass and weight are not the same thing.
much easier to manipulate. For instance, if you need to know how many inches are in a foot, Although we often use the terms mass and weight
you only need to remember what you at one time memorized: 12 inches = 1 foot. But now interchangeably, each one has a specific definition
you need to know how many feet are in a mile. What happens if you never memorized this and usage. The mass of an object is a measure of
fact? Of course you can look it up online or elsewhere, but the point is that this fact must be the amount of matter in it. The mass (amount of
given to you, as there is no way for you to derive it out yourself. This is true about all parts of matter) of an object remains the same regardless
the English system: you have to memorize all the facts that are needed for different of where the object is placed. For example,
measurements. moving a brick to the moon does not cause any
matter in it to disappear or be removed.
Metric Prefixes and Equivalents The weight of an object is the force of
The metric system uses a number of prefixes along with the base units. A base unit attraction between the object and the earth (or
is one that cannot be expressed in terms of other units. The base unit of mass is the gram (g), whatever large body it is resting on). We call this
that of length is the meter (m), and that of volume is the liter (L). Each base unit can be force of attraction the force of gravity. Since the force of gravity is not the same at every
combined with different prefixes to define smaller and larger quantities. When the prefix point on the earth’s surface, the weight of an object is not constant. The gravitational pull on
the object varies depending on where the object is with respect to the Earth or other gravity-
104 producing object. For example, a man who weighs 180 pounds on Earth would weigh only
www.ck12.org 45 pounds if he were in a stationary position, 4,000 miles above the Earth's surface. This
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same man would weigh only 30 pounds on the moon
because the moon's gravity is only one-sixth that of Earth.
The mass of this man, however, would be the same in each
The cubic meter is the SI unit of volume. The cubic meter is a very large unit and is not very
situation because the amount of matter in him is constant.
convenient for most measurements in chemistry. A more common unit is the liter (L) which
Consistency requires that scientists use mass and not
weight in its measurements of the amount of matter. of a cubic meter. Another commonly used volume measurement is the milliliter, which
The basic unit of mass in the International System is equal to of a liter.
of Units is the kilogram. A kilogram is equal to 1000
grams. A gram is a relatively small amount of mass and so Temperature
larger masses are often expressed in kilograms. When very
When used in a scientific context, the words heat and temperature do NOT mean the
tiny amounts of matter are measured, we often use A polar molecule has partially positive and
partially negative charges on opposite sides of
milligrams which are equal to 0.001 gram. There the molecule. same thing. Temperature represents the average kinetic energy of the particles that make up
are numerous larger, smaller, and intermediate mass a material. Increasing the temperature of a material increases its thermal energy. Thermal
units that may also be appropriate. energy is the sum of the kinetic and potential energy in the particles that make up a material.
th Objects do not “contain” heat; rather they contain thermal energy. Heat is the movement of
At the end of the 18 century, a kilogram was thermal energy from a warmer object to a cooler object. When thermal energy moves from
the mass of a liter of water. In 1889, a new international
one object to another, the temperature of both objects change.
prototype of the kilogram was made of a platinum-iridium alloy. The kilogram is equal to the
A thermometer is a device that measures temperature. The name is made up of
mass of this international prototype, which is held in Paris, France.
“thermo” which means heat and “meter” which means to measure. The temperature of a
Length substance is directly proportional to the average kinetic energy it contains. In order for the
average kinetic energy and temperature of a substance to be directly proportional, it is
Length is the measurement of anything from end to end. In science, length usually
necessary that when the temperature is zero, the average kinetic energy must also be zero.
refers to how long an object is. There are many units and sets of standards used in the world for
This is not true with either the Fahrenheit or Celsius temperature scales. Most of are
measuring length. The ones familiar to you is probably inches, feet, yards, and miles.
familiar with temperatures that are below the freezing point of water. It should be apparent
Most of the world, however, measures
distances in meters and kilometers for that even though the air temperature may be -5°C, the molecules of air are still moving.
Substances like oxygen gas and nitrogen gas have already melted and boiled to vapor at
longer distances, and centimeters and
temperatures below -150°C.
millimeters for shorter distances. For
consistency and ease of communication, It was necessary for use in calculations in science for a third temperature scale in
scientists around the world have agreed to which zero degrees corresponds with zero kinetic energy, that is, the point where molecules
use the SI system of standards regardless of cease to move. This temperature scale was designed by Lord Kelvin. Lord Kelvin stated that
the length standards used by the general there is no upper limit of how hot things can get, but there is a limit as to how cold things can
get. Kelvin developed the idea of Absolute Zero, which is the temperature at which
The standard meter used in France in the 18th century. public.
molecules stop moving and therefore, have zero kinetic energy. The Kelvin temperature
The SI unit of length is the meter.
scale has its zero at absolute zero (determined to be -273.15°C), and uses the same size
In 1889, the definition of the meter was a bar of platinum-iridium alloy stored under
degree as the Celsius scale. Therefore, the mathematical relationship between the Celsius
conditions specified by the International Bureau of Standards. In 1960, this definition of the
scale and the Kelvin scale is K=°C + 273. In the case of the Kelvin scale, the degree sign is
standard meter was replaced by a definition based on a wavelength of krypton-86 radiation.
not used. Temperatures are expressed, for example, simply as 450 K.
In 1983, that definition was replaced by the following: the meter is the length of the path
traveled by light in a vacuum during a time interval of of a second. Time
The SI unit for time is the second. The second was originally defined as a tiny
Volume fraction of the time required for the Earth to orbit the Sun. It has since been redefined
The volume of an object is the amount of several times. The definition of a second (established in 1967 and reaffirmed in 1997) is: the
space it takes up. In the SI system, volume is a duration of 9,192,631,770 periods of the radiation corresponding to the transition between
derived unit, that is, it is based on another SI unit. In the two hyperfine levels of the ground state of the cesium-133 atom.
the case of volume, a cube is created with each side
of the cube measuring 1.00 meter. The volume of this Lesson Summary
3 The metric system is an international decimal-based system of measurement.
cube is 1.0m·1.0m·1.0m=1.0 m or one cubic meter.
The metric system uses a number of prefixes along with the base units.
106 The prefixes in the metric system are multiples of 10.
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 The International System of Units, abbreviated SI from the French Le Système
International d’ Unites is, since the 1960s, internationally agreed upon as the standard
metric system.
10) What is the basic unit of measurement in the metric system for mass?
The mass of an object remains the same regardless of where the object is placed.
11) What unit is used in the metric system to measure volume? How is this unit related to
The basic unit of mass in the International System of Units is the kilogram.
the measurement of length?
The SI unit of length is the meter.
12) Would it be comfortable to swim in a swimming pool whose water temperature was 275
Temperature represents the average kinetic energy of the particles that make up a
K? Why or why not?
material.
Absolute Zero is the temperature at which molecules stop moving and therefore,
have zero kinetic energy.
5.2: Scientific Notation
The Kelvin temperature scale has its zero at absolute zero (determined to be - Objectives
273.15°C), and uses the same size degree as the Celsius scale. Identify when scientific notation is useful to record
The mathematical relationship between the Celsius scale and the Kelvin scale is measurements Convert measurements to scientific notation.
K=°C + 273. Convert quantities from scientific notation to their standard numerical form.
The SI unit for time is the second.
Introduction
Vocabulary Work in science frequently involves very large and very small numbers. The speed
Metric system: an international decimal-based system of measurement. of light, for example, is 300,000,000 meters/second; the mass of the earth is
International System of Units (Le Système International d’ Unites): 6,000,000,000,000,000,000,000,000 kg; and the mass of an electron is
the internationally agreed upon standard metric system 0.0000000000000000000000000000009 kg. It is very inconvenient to write such numbers
and even more inconvenient to attempt to carry out mathematical operations with them.
Mass: a measure of the amount of matter in an object
Weight: the force of attraction between the object and the earth (or whatever What is Scientific Notation?
large body it is resting on)
Scientists and mathematicians have designed an easier method for dealing with such
Temperature: the average kinetic energy of the particles that make up a material numbers. This more convenient system is called exponential notation by mathematicians and
Absolute Zero: the temperature at which molecules stop moving and therefore, have scientific notation by scientists.
zero kinetic energy In scientific notation, very large and very small numbers are expressed as the
product of a number between 1 and 10 multiplied by some power of 10. The number
Further Reading / Supplemental Links 9,000,000 for example, can be written as the product of 9 times 1,000,000 and 1,000,000 can
http://en.wikipedia.org/wiki/Si_units 6 6
be written as 10 . Therefore, 9,000,000 can be written as 9x10 . In a similar manner,
-8
5.1: Review Questions 0.00000004 can be written as or 4x10 .
1) List three advantages to using the metric system over the English system or other
measurement systems.

Identify which is bigger in each set of measurements:

2) 1 kg or 1 g Examples of Scientific Notation
3) 10 mg or 10 g Decimal Notation Scientific Notation
4
4) 100 cg or 100 mg 95,672 9.5672x10
3
8,340 8.34x10
Fill in the missing information in the following equivalencies: 100 1x10
2
5) ? g = 1 kg Hydrogen bonding between guanine (G) 0
and cytosine(C), one of two types 7.21 7.21x10
6) 100?=1L -2
of base pairs in DNA. The hydrogen 0.014 1.4x10
7) 1 m = ? cm bonds are shown by dotted lines. -9
0.000000008 8.0x10
-12
8) Why is it important for scientists to use the same system to make measurements? 0.00000000000975 9.75x10
9) What is the basic unit of measurement in the metric system for length? As you can see from the examples above, to convert a number from decimal form to
scientific notation, you count the spaces that you need to move the decimal and that number
108 becomes the exponent of 10. If you are moving the decimal to the left, the exponent is positive
www.ck12.org and if you are moving the decimal to the right, the exponent is negative.
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Calculators and Scientific Notation

 Because you will be performing calculations using scientific notation, it is important
that you understand how your calculator uses scientific notation. For this course, you will
need a scientific or graphing calculator. These calculators have an “exponential” button. It
is typically labeled “EXP” or “EE” and can be read “times ten to the…” Consider the
23
following number: 6.02x10 . If you were to read this out loud, you would say “6.02 times 5.3: Math in Chemistry
rd Objectives
ten to the 23 ”. To type this in your calculator, you would put “6.02EE23” or “6.02EXP23”.
Most calculators would print 6.02E23. By properly using the exponential button, you will Use the factor-label method to solve problems.
avoid common mistakes made by students when multiplying or dividing these numbers. Perform metric conversions using the factor label method for conversions.

Example: Perform the following calculation correctly using the exponential button. Introduction
24 23 During your studies of chemistry (and physics also), you will note that
1.20x10 / 6.02x10
Solution: To type this in my calculator, I would type: mathematical equations are used in a number of different applications. Many of these
OR equations have a number of different variables with which you will need to work. You
The answer is 2.0. *Be careful to look on the far right of your calculator screen. If you see should also note that these equations will often require you to use measurements with their
E47 or a 47 typed offset, you typed it in wrong. Try again. units. Algebra skills become very important here!

Lesson Summary Conversion Factors

Very large and very small numbers in science are expressed in scientific notation. A conversion factor is a factor used to convert one unit of measurement into another.
A simple conversion factor can be used to convert meters into centimeters, or a more
Vocabulary complex one can be used to convert miles per hour into meters per second. Since most
calculations require measurements to be in certain units, you will find many uses for
Scientific notation: a shorthand method of writing very large and very small numbers conversion factors. What always must be remembered is that a conversion factor has to
in terms of a decimal number between 1 and 10 multiplied by 10 to a power. represent a fact; this fact can either be simple or much more complex. For instance, you
already know that 12 eggs equal 1 dozen. A more complex fact is that the speed of light is
5.2: Review Questions 5
1.86 x 10 miles/sec. Either one of these can be used as a conversion factor depending on
1) When is it useful to use scientific notation? what type of calculation you might be working with.
Write the following numbers in scientific notation. Factor-Label Method of Problem Solving
2) 0.0000479 3) 4260
Frequently, it is necessary to convert units measuring the same quantity from one
4) 251,000,000 5) 0.00206 form to another. For example, it may be necessary to convert a length measurement in
Write each of the following numbers in standard notation. meters to millimeters. This process is quite simple if you follow a standard procedure called
6) 2.3x10
4
7) 9.156x10
-4 unit analysis or dimensional analysis. The Factor-Label Method is a technique that
-3 6 involves the study of the units of physical quantities. It affords a convenient means of
8) 7.2x10 9) 8.255x10 checking mathematical equations. This method involves considering both the units you
presently have (given measurement), the units you wish to end up with, and designing
Write the buttons you would push on your calculator to type in the following numbers
conversion factors than will cancel units you don’t want and produce units you do want. The
using the exponential button. conversion factors are created from the equivalency relationships between the units or ratios
14
10) 7.3x10 of how units are related to each other.
-6
11) 6.01x10 In terms of making unit conversions, suppose you want to convert 0.0856 meters into
5
12) 7.98x10 millimeters. In this case, you need only one conversion factor and that conversion factor must
cancel the meters unit and create the millimeters unit. The conversion factor will be created
Using the exponential button, perform each calculation on your calculator.
3 4 from the relationship1000 millimeters (mm) = 1 meter (m).
13) 2.0x10 · 3.0x10
3 4
14) 2.0x10 / 3.0x10
-4 -2
15) 4.2x10 / 3.0x10 Remember that when you multiply fractions and you have the same number on top of one
-7 -3
16) 7.3x10 · 8.0x10 fraction and the bottom of another fraction, the numbers will cancel out leaving one. The
same is true for units. When the above expression is multiplied as indicated, the meters units
110 will cancel and only millimeters will remain. The unit analysis process involves creating
www.ck12.org conversion factors from equivalencies between various units.

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The given table contains many useful conversion factors.

 English Units Metric Units
1 ounces (oz) (weight) 28.35 grams (g)
1 fluid ounce (oz) (volume) 29.6 mL Sometimes, it is necessary to insert a series of conversion factors. Suppose we need
2.205 pounds (lb) 1 kilograms (kg) to convert miles to kilometers and the only equivalencies we know are 1 mile = 5280 feet, 12
1 inch (in) 2.54 centimeters (cm) inches = 1 foot, 2.54 cm = 1 inch, 100 cm = 1m, and 1000 m = 1 km. We will set up a series
.6214miles (mi) 1 kilometer (km) of conversion factors so that each conversion factor produces the next unit in the sequence.
1 quart (qt) 0.95 liters (L)
Example: Convert 12 miles to kilometers.
Metric Prefix Base unit equivalency Solution: Although we have a ratio for miles to kilometers given in the table, we will solve
1000 milli (base unit) 1 base unit this problem using other units to see what a longer process looks like. The answer would be
100 centi (base unit) 1 base unit the same.
1 kilo (base unit) 1000 base units
Of course, there are other ratios which are not listed in this table. They may include:
Ratios embedded in the text of the problem (using words such as per or in each, or In each step, the previous unit is cancelled and the next unit in the sequence is produced.,
using symbols such as / or %) each successive unit cancelling out until only the unit needed in the answer is left.
Conversions in the metric system, as covered earlier in this chapter.
Common knowledge ratios (such as 60 seconds = 1 minute) Conversion factors for area and volume can also be produced by this method.

The general steps you must take in order to solve these problems include: 2 2
Example: Convert 1500 cm to m .
1. Identify the “given” information in the problem. Look for a number with units to Solution:
start this problem with.
2. What is the problem asking you to “find”? In other words, what unit will your
answer have? OR
3. Use ratios and conversion factors to cancel out the units that aren’t part of your
answer, and leave you with units that are part of your answer.
4. When your units cancel out correctly, you are ready to do the math. You are
multiplying fractions, so you multiply the top numbers and divide by the bottom
numbers in the fractions. Lesson Summary
Look for each of these steps in the following examples. Conversion factors are used to convert one unit of measurement into another.
Example: Convert 1.53 grams to centigrams. The factor-label method involves considering both the units you presently have, the
Solution: The equivalency relationship is 100cg=1g (given in the second table), so the units you wish to end up with and designing conversion factors than will cancel units
conversion factor is constructed from this equivalency to cancel grams and produce you don’t want and produce units you do want.
centigrams.
Vocabulary
Conversion factor: a ratio used to convert one unit of measurement into another.

Further Reading / Supplemental Links

Example: Convert 1000 inches to feet.
Tutorial: Vision Learning: Unit Conversion & Dimensional Analysis
Solution: The equivalency between inches and feet is 12 inches = 1 ft. The conversion
factor is designed to cancel inches and produce feet. The balance on the left measures mass, a
property that does not change based on
location. The scale on the right measures
weight, which differs depending on where in
Each conversion factor is designed specifically for the problem. In this case, we know we the universe an object is.
need to cancel inches so we know we need the inches component in the conversion factor to http://visionlearning.com/library/module_viewer.php?mid=144&l=&c3=
be in the denominator.
5.3: Review Questions
112 For each of the following, first A) identify the given, find, and ratios within the
www.ck12.org problem. Then, B) solve the problem using the factor-label method. Show all work and
unit cancellations.
1) What is the diameter of a 9” cake pan in centimeters?
2) It is approximately 52 miles from Spanish Fork to Salt Lake. If I drive 65 miles/hr, how
many minutes will it take to drive there?
 www.ck12.org
113

3) If there are 35 g of sugar in 8 oz of soda, what mass (in grams) of sugar is in an entire atom to one sulfur atom was 12 amu to 32 amu. They realized that if they massed out 12
2 liter bottle? grams of carbon and 32 grams of sulfur, they would have the same number of atoms of each
4) My car gets about 35 miles per gallon. Right now, gas costs $3.69 per gallon. How element. They didn’t know how many atoms were in each pile but they knew the number in
much does it cost me to drive to Salt Lake (52 miles away)? each pile had to be the same. This is the same logic as knowing that if a basketball has twice
5) What is your mass in grams? (Start with your weight in pounds) the mass of a soccer ball and you massed out 100 lbs of basketballs and 50 lbs of soccer
6) Nervous Ned paced for 3 hours while his wife was in the delivery room. If he paces at 5 balls, you would have the same number of each ball. Many years later, when it became
possible to count particles using electrochemical reactions, the number of atoms turned out to
paces every 3 seconds, how far did he go, in miles? (In this case, one pace is 2.2 feet.) 23
(There are 5280 feet per mile.) be 6.02x10 particles. Eventually chemists decided to call that number of particles a mole.
23
7) If I drive 75 miles/hr, how long in minutes will it take me to drive 500 km? The number 6.02x10 is called Avogadro’s number. Avogadro, of course, had no hand in
8) A male elephant seal weighs about 4 tons. What is the mass of the seal in grams? (There determining this number, rather it was named in honor of Avogadro.
are 2000 lbs in one ton)
9) My car gets about 37 miles per gallon. How many km/liter Converting Between Molecules to Moles
The standard kilogram copy stored
is this? (There are 4 quarts in a gallon) and used in Denmark. We can use Avogadro’s number as a conversion factor, or ratio, in dimensional
10) In a nuclear chemistry experiment, an alpha particle is analysis problems. If we are given a number of molecules of a substance, we can convert it
found to have a velocity of 14,285 m/s. Convert this measurement into miles/hour. into moles by dividing by Avogadro’s number and vice versa.
9
5.4: The Mole Example: How many moles are present in 1 billion (1x10 ) molecules of water?
Objectives
Use Avogadro's number to convert to moles and vice versa given the number
of particles of a substance. You should note that this amount of water is too small for even our most delicate balances to
Use the molar mass to convert to grams and vice versa given the number of moles of
determine the mass. A very large number of molecules must be present before the mass is
a substance.
large enough to detect with our balances.
Introduction
Example: How many molecules are present in 0.00100 mol?
When objects are very small, it is often inconvenient or inefficient, or even impossible
to deal with the objects one at a time. For these reasons, we often deal with very small objects Solution:
in groups, and have even invented names for various numbers of objects. The most common of
these is “dozen” which refers to 12 objects. We frequently buy objects in groups of 12, like Converting Grams to Moles and Vice Versa
doughnuts or pencils. Even smaller objects such as straight pins or staples are usually sold in 23
1.00 mol of carbon-12 atoms has a mass of 12.0 g and contains 6.02x10 atoms.
boxes of 144, or a dozen dozen. A group of 144 is called a “gross.”
This problem of dealing with things that are too small to operate with as single items 1.00 mole of any element or compound has a mass equal to its molecular mass in grams and
also occurs in chemistry. Atoms and molecules are too small to see, let alone to count or 23
contains 6.02x10 particles. The mass, in grams, of 1 mole of particles of a substance is
measure. Chemists needed to select a group of atoms or molecules that would be convenient now called the molar mass (mass of 1.00 mole).
to operate with. To quickly find the molar mass of a substance, you need to look up the masses on the
periodic table and add them together. For example, water has the formula H2O. Hydrogen
Avogadro's Number has a mass of 1.0084 g/mol (see periodic table) and oxygen has a mass of 15.9994 g/mol.
In chemistry, it is impossible to deal with a single atom or molecule because we The molar mass of H2O=2(1.0084g/mol) + 15.9994g/mol = 18.0162g/mol. This means that 1
can’t see them or count them or weigh them. Chemists have selected a number of particles mole of water has a mass of 18.0162 grams.
with which to work that is convenient. Since molecules are extremely small, you may
suspect that this number is going to be very large and you are right. The number of particles We can also convert back and forth between grams of substance and moles. The
23 conversion factor for this is the molar mass of the substance. The molar mass is the ratio
in this group is 6.02x10 particles and the name of this group is the mole (the abbreviation
23 giving the number of grams for each one mole of a substance. This ratio is easily found by
for mole is mol). One mole of any object is 6.02x10 of those objects. There is a very adding up the atomic masses of the elements within a compound using the periodic table.
particular reason that this number was chosen and we hope to make that reason clear to you. This ratio has units of grams per mole or g/mol.
When chemists are carrying out chemical reactions, it is important that the
relationship between the numbers of particles of each reactant is known. Chemists looked at
the atomic masses on the periodic table and understood that the mass ratio of one carbon

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Example: Find the molar mass of each of c) F2
the following: d) H2SO4

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a) S e) Al2(SO4)3
b) H2O
Solution: You will need a periodic table to solve these problems. Look for each element’s
mass.
a) Look for sulfur on the periodic table. Its molar mass is 32.065 g/mol. That means that
one mole of sulfur has a mass of 32.065 grams.
b) This compound contains two hydrogen atoms and one oxygen atoms. To find the molar
mass of H2O, we need to add the mass of two hydrogen atoms plus the mass of one
oxygen atom. We get: 2(1.008) + 16.00 = 18.016 g/mol. That means that one mole of
water has a mass of just over 18 grams.
c) This compound contains two fluorine atoms. To find the molar mass of F2, we need to add
the mass of two fluorine atoms. We get: 2(19.00) = 38.00 g/mol
d) This compound contains two hydrogen atoms, one sulfur atom, and four oxygen atoms.
To find the molar mass of H2SO4, we need to add the mass of two hydrogen atoms
plus the mass of one sulfur atom plus the mass of four oxygen atoms. We get: 2(1.008)
+ 32.065 + 4(16.00) = 100.097 g/mol
e) This compound contains two aluminum atoms, three sulfur atoms, and twelve oxygen
atoms. To find the molar mass of Al2(SO4)3, we need to add the mass of all of these
atoms. We get: 2(26.98) + 3(32.065) + 12(16.00) = 342.155 g/mol

To convert the grams of a substance into moles, we use the ratio molar mass. We
divide by the molar mass and to convert the moles of a substance into grams, we multiply
by the molar mass.

Example: How many moles are present in 108 grams of water?

Solution:
To get the ratio 1 mol H2O=18.02 g, we added up the molar mass of H2O using the masses
on a periodic table.

Example: What is the mass of 7.50 mol of CaO?

Solution:
To get the ratio 1 mol CaO=56.0 g, we added up the molar mass of CaO using the masses
on a periodic table.

We will be using these ratios again to solve more complex problems in the next
chapters. Being able to use these ratios is a very important skill for later math problems.

Lesson Summary
23
There are 6.02x10 particles in 1.00 mole. This number is called
Avogadro’s number.
The molar mass of a substance can be found by adding up the masses on a
periodic table.
Using the factor-label method, it is possible to convert between grams, moles, and the
number of atoms or molecules.
 What is the molar mass of each of the following substances? Include units with your answer.
5) H2O 6) NaOH 7) NH4Cl
Vocabulary 8) H2SO4 9) Al2(CO3)3 10) PbO2
23
Avogadro's number: The number of objects in a mole; equal to 6.02x10 . Convert the following to moles.
Mole: An Avogadro’s number of objects. 11) 60.0 g NaOH 12)5.70 g H2SO4
Molar Mass: The mass, in grams, of 1 mole of a substance. This can be found by adding up the masses on the periodic
table. 13) 2.73 g NH4Cl 14) 10.0g PbO2
Convert the following to grams.
Further Reading / Supplemental Links 15) 0.100 mol CO2 16)0.500 mol (NH4)2CO3
http://learner.org/resources/series61.html The learner.org website allows users to view streaming videos of the Annenberg
series of chemistry videos. You are required to register before you can watch the videos but there is no charge. The website 17) 0.437 mol NaOH 18) 3.00 mol H2O
has one video that relates to this lesson called The Mole. How many molecules are present in the following masses?
Using Avogadro's law, the mass of a substance can be related to the number of particles contained in that mass. The
Mole: (http://www.learner.org/vod/vod_window.html?pid=803) 19) 1.00 g Na2CO3 20) 1000. g H2O
Vision Learning tutorial: The Mole Convert the following to grams.
http://visionlearning.com/library/module_viewer.php?mid=53&l=&c3= 23 24
21) 2.0x10 molecules H2 22)1.75x10 molecules NaCl
22
5.4: Review Questions 23) 8.6x10 molecules NaOH
How many molecules are present in the following quantities? All images, unless otherwise stated, are created by the CK-12 Foundation and are under
1) 0.250 mol H2O 2) 0.0045 mol Al2(CO3)3 the Creative Commons license CC-BY-NC-SA.
How many moles are present in the following quantities?
20
3) 1.0x10 molecules H2O 4) 5 billion atoms of carbon
117
www.ck12.org In this image, the dark particles are the
solvent particles as there are fewer of
them. The solute particles are the
lighter colored particles.
Chapter 6: Mixtures & Their Properties and visible places in which there are more zinc atoms). CC – Tracy Poulsen
Thus the combination of zinc filings and copper pieces in a pile does not represent a
6.1: Solutions, Colloids, and Suspensions homogeneous mixture, but is, instead a heterogeneous mixture. In a solution, the particles
are so small that they cannot be distinguished by the naked eye. In a solution, the mixture
Objectives would have the same appearance and properties in all places throughout the mixture.
Define a solution.
The point should be made that because solutions have the same composition
Identify the solute and solvent in a solution.
throughout does not mean you cannot vary the composition. If you were to take one cup of
Explain the differences among solutions and heterogeneous mixtures, such as colloids
water and dissolve ¼ teaspoon of table salt in it, a solution would form. The solution would
and suspensions.
have the same properties throughout, the particles of salt would be so small that they would
Introduction not be seen and the composition of every milliliter of the solution would be the same. But
you can vary the composition of this solution to a point. If you were to add another ½
In chapter 4, we distinguished between pure substances and mixtures. Remember, a
teaspoon of salt to the cup of water, you would make
mixture contains two or more pure substances that are not bonded together. These
another solution, but this time there would be a
substances remain unbounded to each other, but are mixed within the same container. They
different composition than the last. You still have a
also retain their own properties, such as color, boiling point, etc. There are two types of
mixtures: mixtures in which the substances are evenly mixed together (called a solution) and solution where the salt particles are so small that they
a mixture in which the substances are not evenly mixed (called a heterogeneous mixture). would not be seen and the solution has the same
In this chapter we begin our study of solution chemistry. We all probably think we properties throughout, thus it is homogeneous.
The solvent and solute are the two basic parts
know what a solution is. We might be holding a can of soda or a cup of tea while reading
of a solution. The solvent is the substance present in
this book and think … hey this is a solution. Well, you are right. But you might not realize the greatest amount, whichever substance there is
that alloys, such as brass, are also classified as solutions, or that air is a solution. Why are more of in the mixture. The solvent is frequently, but
these classified as solutions? Why wouldn’t milk be classified as a solution? To answer these not always, water. The solute, then, is the substance
questions, we have to learn some specific properties of solutions. Let’s begin with the present in the least amount. Let’s think for a minute
definition of a solution and view some of the different types of mixtures. that you are making a cup of hot chocolate. You take a
teaspoon of cocoa powder and dissolve it in one cup of
Homogeneous Mixtures
 A solution is an even (or homogeneous) hot water. Since the cocoa powder is in the lesser
mixture of substances. When you consider that the amount it is said to be the solute; and the water is the
prefix “homo” means “same”, this definition makes solvent since it is in the greater amount.
perfectly good sense. Solutions carry the same
properties throughout. Take, for example, vinegar that Example: Name the solute and solvent in each of the following solutions.
is used in cooking is approximately 5% acetic acid in
(a) salt water
water. This means that every teaspoon of vinegar that is Solutions must have particles that are (b) air
removed from the container contains 5% acetic acid uniformly spread throughout. In this Solution:
and 95% water. This ratio of mixing is carried out image, A and B are not solutions as the
(a) solute = salt; solvent = water
throughout the entire container of vinegar. red particles are not evenly spread.
A point should be made here that when a Mixture C, however, is a solution. (b) solute = oxygen; solvent = nitrogen
CC – Tracy Poulsen
solution is said to have uniform properties throughout, the definition is referring to properties Colloids and Suspensions
at the particle level. Well, what does this mean? Let's consider brass as an example. The Two other types of mixtures that we will compare to solutions include colloids
brass is an alloy made from copper and zinc. To the naked eye a brass coin seems like it is and suspensions. These mixtures are frequently confused with solutions, but these are
just one substance but at a particle level two substances are present (copper and zinc) and the heterogeneous, not homogeneous, mixtures.
copper and zinc atoms are evenly mixed at the atomic level. So the brass represents a Recall that a solution is a mixture of substances in such a way that the final product
homogeneous mixture. Now, consider a handful of zinc filings and copper pieces. Is this has the same composition throughout. Remember the example of vinegar that is 5%, by
now a homogeneous solution? The properties of any scoop of the “mixture” you are holding mass, acetic acid in water. This clear liquid is a solution since light easily passes through it
would not be consistent with any other scoop you removed from the mixture. The ratio of and it never separates. All liquid solutions have this shared property, in which the particles
copper and zinc may be different. Additionally, you would see differences in the color at are so small that light goes straight through. In other words, the mixture is clear or see-
different places in the mixture (there are visible places in which there are more copper atoms through. It is important to note, however, that clear does not necessarily mean colorless.

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Neither colloids nor suspensions are
classified as solutions, but are special types of
heterogeneous mixtures instead. In order to be as
On the other hand, colloids are mixtures in which the size of the particles is between
3 6 -9 -6 solution, you must have very small particles
1x10 pm and 1x10 pm. In meters, these sizes translate to 1x10 m to 1x10 m – between evenly distributed, so that the mixture has the
10 and 1000 times smaller than a small grain. These particles, although sounding small, are same properties throughout. Colloids and
still much bigger than the particles in a solution. suspensions have particles that are too big to be
A common example of a colloid is milk. One way to tell that milk is a colloid is by considered a solution.
the Tyndall effect. The Tyndall effect is the scattering of light by particles. This involves
shining a light through the mixture: when the light is shined through a colloid, the light is Example: Label each of the following mixtures as a solution, colloid, or suspension.
does not go straight through, but is instead scattered. Note that milk is not see-through, but
a) Italian salad dressing
has a cloudy appearance. Because light not allowed to pass through the mixture, the mixture
b) Mustard
is considered a colloid. When light is passed through a solution, the particles are so small
c) Apple juice
that they do not obstruct the light. However, when light is passed through a colloid, since the
Solution:
particles are larger, they will act as an obstruction to the light and the light is scattered. The
a) suspension – when left to sit, it separates into layers
particles in a colloid, while able to scatter light, are still small enough so that they do not
b) colloid – although it does not separate into
settle out of solution.
layers like suspensions do, mustard does not let
It is amazing just how common colloids are to us in our everyday lives. Some
light go through
common colloids you may have seen include milk of magnesia, mayonnaise, jell-o,
c) solution – apple juice doesn’t separate into
and marshmallows.
layers like suspensions do, but apple juice will let
Suspensions are mixtures which contain even bigger particles than solutions or
light through so it is a solution and not a colloid.
colloids do. In suspensions, particles settle into layers within a container if they are left
standing. This means that the particles in a suspension are large enough so that gravity pulls
Lesson Summary
them out of solution. With suspensions, filtration can usually be used to separate the excess
particles from the solution. If a suspension is passed through a piece of filter paper (or a Generally speaking, in a solution, a solute
coffee filter) some particles will go through and others will be stopped in the filter paper. A is present in the least amount (less than
common example of a suspension is muddy water. If you had a beaker of water and added a 50% of the solution) whereas the solvent
handful of fine dirt, even if you stirred it, when you let it stand, dirt would settle to the is present in the greater amount (more
than 50% of the solution).
bottom.
 120
www.ck12.org A solution is a mixture that has the same properties throughout.
Common examples of colloids include milk, butter, Jell-O, and clouds.
Suspensions are mixtures in which the particles are large enough so that they settle to
the bottom of the container and can be filtered using filter paper.

Vocabulary
Solution: a homogeneous mixture of substances
Solvent: the substance in a solution present in the greatest amount
Solute: the substance in a solution present in the least amount
3
Colloid: type of mixture in which the size of the particles is between 1x10 pm and
8
1x10 pm
Suspension: type of mixture in which the particles settle to the bottom of the
container and can be separated by filtration

6.1: Review Questions

1) Distinguish between a solution, a colloid, and a suspension.
2) What is one way to tell you have a colloid and not a solution?

Multiple choice
3) The biggest difference between a colloid and a suspension is that:
a) In colloids, the solute is permanently dissolved in the solvent.
b) In colloids the particles eventually settle to the bottom.
c) In suspensions the particles eventually settle to the bottom.
d) None of these are correct
4) Karen was working in the lab with an unknown solution. She noticed that there was no
precipitate in the bottom of the beaker even after it had been on the lab bench for
several days. She tested it with a light and saw that light scattered as it passed through
the solution. Karen concluded that the liquid was what type of a mixture?
a) colloid
b) suspension
c) solution

6.2: Solution Formation

Objectives
Explain why solutions form.
Explain the significance of the statement “like dissolves
like.” Discuss the idea of water as the “universal solvent”.
Explain how water molecules attract ionic solids when they dissolve in water.

Introduction
We have learned that solutions can be formed in a variety of combinations using
solids, liquids, and gases. We also know that solutions have constant composition and we can
also vary this composition up to a point to maintain the homogeneous nature of the solution.
But how exactly do solutions form? Why is it that oil and water will not form a solution and
yet vinegar and water will? Why could we dissolve table salt in water but not in vegetable

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oil? The reasons why solutions form will be explored in this section, along with a discussion
of why water is used most frequently to dissolve substances of various types.
to be attracted to water as well. Because of these partial charges, polar molecules are able to
Ionic Compounds in Solution
dissolve in other polar compounds.
Recall that metals form positive ions by losing electrons and nonmetals form
negative ions by gaining electrons. In ionic compounds, the ions in the solid are held If you mix a nonpolar compound with a polar compound, they will not form an even
together by the attraction of these oppositely charges particles. Since ionic compounds can mixture. The polar compound is more attracted to the other molecules of the same compound
dissolve in polar solutions, specifically water, we can extend this concept to say that ions than they are attracted to the nonpolar compound. If you have tried to mix oil and water
themselves are attracted to the water molecules because the ions of the ionic solid are together you may have witnessed this. Water is much more polar than oil, so the oil does not
attracted to the polar water molecule. When you dissolve table salt in a cup of water, the dissolve in the water. Instead, you will see two different layers form.
table salt dissociates into sodium ions and chloride ions: However, when a nonpolar compound is mixed with another nonpolar compound,
+ - neither of them have partial charges to be attracted to. They are instead attracted by London
NaCl(s) → Na (aq) + Cl (aq) dispersion forces and are able to dissolve together, forming a solution. The similarity in type
How does salt dissolve, though? Dissolving is based on electrostatic attraction, that and strength of intermolecular forces allows two nonpolar compounds such as CO2 and
is, the attraction between positive and negative charges. The sodium ions get attracted to the
partially negative ends of the water molecule and the chloride ions get attracted to the benzene, C6H6.
partially positive end of the water molecule. When we studied how ionic solids dissolve, we said that as they dissolve in solution,
these solids separate into ions. More specifically, ionic solids separate into their positive ions
and negative ions in solution. This is not true for molecular compounds. Molecular
compounds are held together with covalent bonds meaning they share electrons. When they
share electrons, their bonds do not easily break apart, thus the molecules stay together even
in solution. For example, when you dissolve a spoonful of sugar into a glass of water, the
intermolecular forces are broken but not the bonds. You can write the following equation for
the dissolution of sugar in water.
C12H22O11(s) → C12H22O11(aq)
Notice how the molecules of sugar are now separated by water molecules (aq). In other
words, sugar molecules are separated from neighboring sugar molecules due to attraction for
Most ionic compounds dissolve in water as the positive ions are attracted to the negative side of
the water, but the molecules themselves have not. The bonds within the molecules have not
water molecules and the negative ions are attracted to the positive sides of water molecules.
broken.
Example: Which compounds will dissolve in solution to separate into ions?
(a) LiF
To understand why salt will dissolve in water, we first must (b) P2F5
remember what it means for water to be polar. The more (c) C2H5OH
electronegative oxygen atom pulls the shared electrons away from Solution:
the hydrogen atoms in a water molecule causing an unequal LiF will separate into ions when dissolved in solution, because it is an ionic compound.
distribution of electrons. The hydrogen end of the water molecule P2F5 and C2H5OH are both covalent and will stay as molecules in a solution.
will be slightly positive and the oxygen end of the water molecule
will be slightly negative. These partial charges allow water to be A simple way to predict which compounds will dissolve in other compounds is the
attracted to the various ions in salt, which pulls the salt crystal apart. Water is a polar phrase “like dissolves like”. What this means is that polar compounds dissolve polar
The same is process true for any ionic compound dissolving molecule, meaning it compounds, nonpolar compounds dissolve nonpolar compounds, but polar and nonpolar do
in water. The ionic compound will separate into the positive and has a partial positive not dissolve in each other.
negative ions and the positive ion will be attracted to the partially and negative side.
These partial charges
Even some nonpolar substances dissolve is water but only to a limited degree.
negative end of the water molecules (oxygen) while the negative ion allow water to Have you ever wondered why fish are able to breathe? Oxygen gas, a nonpolar molecule,
will be attracted to the partially positive end of the water molecules dissolve other does dissolve in water and it is this oxygen that the fish take in through their gills. Or, one
(hydrogen). substances with more example of a nonpolar compound that dissolves in water is the reason we can enjoy
charges. carbonated sodas. Pepsi-cola and all the other sodas have carbon dioxide gas, CO2, a
Covalent Compounds in Solution nonpolar compound, dissolved in a sugar-water solution. In this case, to keep as much gas
Some other covalent compounds, aside from water, are also polar. Having these partial in solution as possible, the sodas are kept under pressure.
positive and negative charges within the molecule gives polar compounds the ability
This general trend of “like dissolves like” is summarized in the following table:

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 Combination Solution Formed? Define concentration, and list the common units used to express the concentration
Polar substance in a polar substance. Yes of solutions.
Non-polar substance in a non-polar substance. Yes Calculate concentration in units of molarity or molality.
Polar substance in a non-polar substance. No Calculate the amount of solute needed to make a given amount of solution with
Non-polar substance in a polar substance. No a given concentration.
Ionic substance in a polar substance. Yes
Introduction
Ionic substance in a non-polar substance. No
Concentratio
Note that every time charged particles (ionic compounds or polar substances) are n is the measure of
how much of a
mixed, a solution is formed. When particles with no charges (nonpolar compounds) are
given substance is
mixed, they will form a solution. However, if substances with charges are mixed with other
mixed with another
substances without charges a solution does not form. substance. Solutions
can be said to be
Lesson Summary
dilute or
Whether or not solutions are formed depends on the similarity of polarity or the concentrated. When
“like dissolves like” rule. we say that vinegar
Polar molecules dissolve in polar solvents, non-polar molecules dissolve in non- The solution on the left is more concentrated than the solution on the right
is 5% acetic acid in because there is a greater ratio of solute (red) to solvent (blue) particles. The
polar solvents. water, we are giving solute particles are closer together. The solution on the right is more dilute (less
Ionic compounds dissolve in polar solvents, especially water. This occurs when the the concentration. If concentrated)
positive cation from the ionic solid is attracted to the negative end of the water we said the mixture
CC – Tracy Poulsen

molecule (oxygen) and the negative anion of the ionic solid is attracted to the positive was 10% acetic acid, this would be more concentrated than the vinegar solution.
end of the water molecule (hydrogen). A concentrated solution is one in which there is a large amount of solute in a given
Water is considered as the universal solvent since it can dissolve both ionic and amount of solvent. A dilute solution is one in which there is a small amount of solute in a
polar solutes, as well as some non-polar solutes (in very limited amounts). given amount of solvent. A dilute solution is a concentrated solution that has been, in
essence, watered down. Think of the frozen juice containers you buy in the grocery store.
Vocabulary
What you have to do is take the frozen juice from inside these containers and usually empty
Miscible: liquids that have the ability to dissolve in each other 3 or 4 times the container size full of water to mix with the juice concentrate and make your
Immiscible: liquids that do not have the ability to dissolve in each other container of juice. Therefore, you are diluting the concentrated juice. When we talk about
electrostatic attraction: the attraction of oppositely charged particles solute and solvent, the concentrated solution has a lot of solute verses the dilute solution that
would have a smaller amount of solute.
6.2: Review Questions The terms “concentrated” and “dilute” provide qualitative methods of describing
1) What does the phrase “like dissolves like” mean? Give an example. concentration. Although qualitative observations are necessary and have their place in every
2) Why will LiCl not dissolve in CCl4? part of science, including chemistry, we have seen throughout our study of science that
3) In which compound will you expect benzene, C6H6, to dissolve? there is a definite need for quantitative measurements in science. This is particularly true in
a) Carbon tetrachloride, CCl4 solution chemistry. In this section, we will explore some quantitative methods of expressing
b) water solution concentration.
c) none of the above
4) Thomas is making a salad dressing for supper using balsamic vinegar and oil. He Molarity
shakes and shakes the mixture but cannot seem to get the two to dissolve. Explain to Of all the quantitative measures of concentration, molarity is the one used most
Thomas why they will not dissolve. frequently by chemists. Molarity is defined as the number of moles of solute per liter of
solution. The symbol given for molarity is M or moles/liter. Chemists also used square
brackets to indicate a reference to the molarity of a substance. For example, the expression
6.3: Concentration +
[Ag ] refers to the molarity of the silver ion. Solution concentrations expressed in
Objectives molarity are the easiest to calculate with but the most difficult to make in the lab.
Define the terms "concentrated" and "dilute".

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 To solve these problems, we will set them using the factor-label method. To review
these steps: Example: Calculate the molality of a solution of hydrochloric acid where 12.5 g of
1. Identify the “given” information in the problem. Look for a number with units to hydrochloric acid, HCl, has been dissolved in 115 g of water.
start this problem with. Solution:
2. What is the problem asking you to “find”? In other words, what unit will your Given: 12.5 g HCl, 115 g H2O
answer have? Find:
3. Use ratios and conversion factors to cancel out the units that aren’t part of
your answer, and leave you with units that are part of your answer.
4. When your units cancel out correctly, you are ready to do the math. You are
multiplying fractions, so you multiply the top numbers and divide by the bottom
numbers in the fractions.
Although these units of concentration are those which chemists most frequently use,
Example: What is the concentration, in mol/L, where 137 g of NaCl has been dissolved in they are not the ones you are most familiar with. Most commercial items you buy at the
enough water to make 500. mL of solution? grocery store have concentrations reported as percentages. For example, hydrogen peroxide
Solution: you buy is approximately 3% hydrogen peroxide in water; a fruit drink may be 5% real fruit
Given: 137 g NaCl, 500. mL solution juice. This unit is convenient for these purposes, but not very useful for many chemistry
problems. Molarity and molality are preferred because these units involve moles, or how
Find: many solute particles there are in a given amount of solution. This comes in handy when
performing calculations involving reactions between solutions.
Another common unit of concentration is parts per million (ppm) or parts per billion
(ppb). If you have ever looked at the annual water quality report for your area, contaminants
in water are typically reported in these units. These units are very useful for concentrations
Example: What mass of potassium sulfate is in 250. mL of 2.50 M potassium that are really low. A concentration of 1 ppm says that there is 1 gram of the solute for every
sulfate, K2SO4, solution? million grams of the mixture. Because we will not deal with concentrations this low
Solution: throughout most of this course, we will not use this unit in our calculations. However, you
Given: 250 mL solution should be aware of it and understand it when you see it.
Find: g K2SO4
Ratios: 2.50 M or 2.50 mol K2SO4/1 L solution Lesson Summary
Concentration is the measure of how much of a given substance is mixed with another
substance.
Molarity is the number of moles of solute per liter of solution.
Molality is calculated by dividing the number of moles of solute by the kilograms
Molality of solvent. It is less common than molarity but more accurate because of its lack of
Molality is another way to measure concentration of a solution. It is calculated by dependence on temperature.
dividing the number of moles of solute by the number of kilograms of solvent. Molality has
Vocabulary
the symbol, m.
Concentration: the measure of how much of a given substance is mixed with another
substance
Molarity, if you recall, is the number of moles of solute per volume of solution. Volume Concentrated: a solution in which there is a large amount of solute in a given amount
is temperature dependent. As the temperature rises, the molarity of the solution will of solvent
actually decrease slightly because the volume will increase slightly. Molality does not Dilute: a solution in which there is a small amount of solute in a given amount of
involve volume, and mass is not temperature dependent. Thus, there is a slight advantage solvent
to using molality over molarity when temperatures move away from standard conditions. Molarity: the number of moles of solute per liter of solution
Molality: the number of moles of solute per kilograms of solution
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6.3: Review Questions
1) Most times when news reports indicate the amount of lead or mercury found in foods,
they use the concentration measures of ppb (parts per billion) or ppm (parts per
million). Why use these over the others we have learned? Boiling Point Elevation
2) What is the molarity of a solution prepared by dissolving 2.5g of LiNO3 in Water boils at 100°C at 1 atm of pressure but a solution of salt water does not. When
sufficient water to make 60.0 mL of solution?
table salt is added to water the resulting solution has a higher boiling point than the water did
3) Calculate the molality of a solution of copper(II) sulfate, CuSO4, where 11.25g of
by itself. The ions form an attraction with the solvent particles that then prevent the water
the crystals has been dissolved in 325 g of water.
4) What is the molarity of a solution made by mixing 3.50g of potassium chromate, molecules from going into the gas phase. Therefore, the salt-water solution will not boil at
K2CrO4, in enough water to make 100. mL of solution? 100°C. In order to cause the salt-water solution to boil, the temperature must be raised above
5) What is the molarity of a solution made by mixing 50.0g of magnesium 100°C in order to allow the solution to boil. This is true for any solute added to a solvent; the
nitrate, Mg(NO3)2, in enough water to make 250. mL of solution? boiling point of the solution will be higher than the boiling point of the pure solvent (without
the solute). In other words, when anything is dissolved in water the solution will boil at a
6) Find the mass of aluminum nitrate, Al(NO3)3, required to mix with 750g of water to higher temperature than pure water would.
make a 1.5m solution.
The boiling point elevation due to the presence of a solute is also a colligative
7) The Dead Sea contains approximately 332 g of salt per kilogram of seawater. Assume
property. That is, the amount of change in the boiling point is related to number of
this salt is all NaCl. What is the molality of the solution?
particles of solute in a solution and is not related to chemical composition of the solute. A
8) What is the molarity of a solution prepared by mixing 12.5 grams FeCl3 in enough
water to make 300 mL of solution? 0.20 m solution of table salt and a 0.20 m solution of hydrochloric acid would have the
9) If 5 grams of NaCl are mixed in enough water to make .5L of solution. What is the same effect on the boiling point.
molarity of the solution?
Freezing Point Depression
10) What is the molality of a solution made by mixing 15 grams of Ba(OH)2 in 250 grams
of water? The effect of adding a solute to a solvent has the opposite effect on the freezing point
of a solution as it does on the boiling point. A solution will have a lower freezing point than a
11) A solution is made by mixing 10.2 grams of CaCl2 in 250 grams of water. What is the
molality of the solution? pure solvent. The freezing point is the temperature at which the liquid changes to a solid. At
a given temperature, if a substance is added to a solvent (such as water), the solute-solvent
6.4: Colligative Properties interactions prevent the solvent from going into the solid phase. The solute-solvent
interactions require the temperature to decrease further in order to solidify the solution. A
Objectives common example is found when salt is used on icy roadways. Here the salt is put on the
Explain what the term "colligative" means, and list the colligative properties.
roads so that the water on the roads will not freeze at the normal 0°C but at a lower
Indicate what happens to the boiling point and the freezing point of a solvent when
temperature, as low as -9°C. The de-icing of planes is another common example of freezing
a solute is added to it.
point depression in action. A number of solutions are used but commonly a solution such as
Calculate boiling point elevations and freezing point depressions for a solution
ethylene glycol, or a less toxic monopropylene glycol, is used to de-ice an aircraft. The
Introduction aircrafts are sprayed with the solution when the temperature is predicted to drop below the
freezing point. The freezing point
People who live in colder climates have seen the trucks put salt on the roads when
depression is the difference in the
snow or ice is forecast. Why do they do that? As a result of the information you will explore
freezing points of the solution
in this section you will understand why these events occur. You will also learn to calculate
from the pure solvent. This is true
exactly how much of an effect a specific solute can have on the boiling point or freezing
for any solute added to a solvent;
point of a solution.
the freezing point of the solution
Colligative Properties will be lower than the freezing
point of the pure solvent (without
The example given in the introduction is an example of a colligative property.
the solute). In other words, when
Colligative properties are properties that differ based on the concentration of solute in a
anything is dissolved in water the
solvent, but not on the type of solute. What this means for the example above is that people
solution will freeze at a lower
in colder climate don’t necessary need salt to get the same effect on the roads – any solute
temperature than pure water
will work. However, the higher the concentration of solute, the more these properties will
would.
change.
The freezing point
depression elevation due to the
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colligative property. That is, the amount of change in the freezing point is related to number
of particles of solute in a solution and is not related to chemical composition of the solute. A
0.20 m solution of table salt and a 0.20 m solution of hydrochloric acid would have the same
0.1m CaI2 will have the lowest freezing point, followed by 0.1m NaCl, and the highest of the
effect on the freezing point.
three solutions 0.1m C6H12O6, but all of them will have a lower freezing point than pure
Comparing the Freezing and Boiling Point of Solutions water.
Recall that covalent and ionic compounds do not dissolve in the same way. Ionic
compounds break up into cations and anions when they dissolve. Covalent compounds do The Mathematics of Boiling Point and Freezing Point Changes
not break up. For example a sugar/water solution stays as sugar + water with the sugar The boiling point of a solution is higher than the boiling point of a pure solvent and
molecules staying as molecules. Remember that colligative properties are due to the number the freezing point of a solution is lower than the freezing point of a pure solvent. However,
of solute particles in the solution. Adding 10 molecules of sugar to a solvent will produce 10 the amount to which the boiling point increases or the freezing point decreases depends on
solute particles in the solution. When the solute is ionic, such as NaCl however, adding 10 the amount solute that is added to the solvent. A mathematical equation is used to
formulas of solute to the solution will produce 20 ions (solute particles) in the solution. calculate the boiling point elevation or the freezing point depression.
Therefore, adding enough NaCl solute to a solvent to produce a 0.20 m solution will have The boiling point elevation is the amount the boiling temperature increases compared
twice the effect of adding enough sugar to a solvent to produce a 0.20 m solution. to the original solvent. For example, the boiling point of pure water at 1.0 atm is 100°C
Colligative properties depend on the number of solute particles in the solution. while the boiling point of a 2% salt-water solution is about 102°C. Therefore, the boiling
“i” is the number of particles that the solute will dissociate into upon mixing with the point elevation would be 2°C. The freezing point depression is amount the freezing
solvent. For example, sodium chloride, NaCl, will dissociate into two ions so for NaCl i = 2, temperature decreases.
for lithium nitrate, LiNO3, i = 2, and for calcium chloride, CaCl2, i = 3. For covalent Both the boiling point elevation and the freezing point depression are related to the
compounds, i is always equal to 1. molality of the solutions. Looking at the formulas for the boiling point elevation and freezing
By knowing the molality of a solution and the number of particles a compound point depression, we can see similarities between the two. The equation used to calculate the
will dissolve to form, it is possible to predict which solution in a group will have the increase in the boiling point is:
lowest freezing point. ΔTb=kb·m·i
To compare the boiling or freezing points of solutions, follow these general steps: Where:
1. Label each solute as ionic or covalent. ΔTb = the amount the boiling temperature increased
2. If the solute is ionic, determine the number of ions in the formula. Be careful to look kb = the boiling point elevation constant which depends on the solvent (for water,
for polyatomic ions. this number is 0.515ºC/m)
3. Multiply the original molality (m) of the solution by the number of particles formed m = the molality of the solution
when the solution dissolves. This will give you the total concentration of particles i = the number of particles formed when that compound dissolves. (for covalent
dissolved. compounds, this number is always 1)
4. Compare these values. The higher total concentration will result in a higher boiling
point and a lower freezing point. The following equation is used to calculate the decrease in the freezing point:
ΔTf=kf·m·i
Example: Rank the following solutions in water in order of increasing (lowest to highest) Where:
freezing point: ΔTf = the amount the freezing temperature decreased
0.1 m NaCl 0.1 m C6H12O6 0.1 m CaI2 kf = the freezing point depression constant which depends on the solvent (for
Solution: water, this number is 1.86ºC/m)
To compare freezing points, we need to know the total concentration of all particles when the m = the molality of the solution
solute has been dissolved. i = the number of particles formed when that compound dissolves. (for covalent
0.1m NaCl: this compound is ionic (metal with nonmetal), and will dissolve into 2 compounds, this number is always 1)
parts. The total final concentration is: (0.1m)(2) = 0.2m
0.1m C6H12O6: this compound is covalent (nonmetal with nonmetal), and will stay as Example: Antifreeze is used in automobile radiators to keep the coolant from freezing. In
1 part. The total final concentration is: (0.1m)(1) = 0.1m geographical areas where winter temperatures go below the freezing point of water, using
0.1m CaI2: this compound is ionic (metal with nonmetal), and will dissolve into 3 pure water as the coolant could allow the water to freeze. Since water expands when it
freezes, freezing coolant could crack engine blocks, radiators, and coolant lines. The main
parts. The total final concentration is: (0.1m)(3) = 0.3m
Remember, the greater the concentration of particles, the lower the freezing point will be. component in antifreeze is ethylene glycol, C2H4(OH)2. What is the concentration of
ethylene glycol in a solution of water, in molality, if the freezing point dropped by 2.64°C?
130 The freezing point constant, kf for water is -1.86°C/m.
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 ΔTf=kf·m·i
Substituting in the appropriate values we get:
2.64°C=(1.86°C/m)(m)(1) d) the freezing point decreases.
Solve for m by dividing both sides by 1.86°C/m. 2) Why do we put salt on ice on the roads in the winter? What effect does it have on the
m=1.42 ice? (Do NOT say that it melts the ice. What does it REALLY do?)
3) Besides adding flavor, what effect does adding salt to water that you cook spaghetti in?
Example: A solution of 10.0g of sodium chloride is added to 100.0g of water in an attempt 4) How do covalent and ionic compounds differ in how they dissolve? How does this
to elevate the boiling point. What is the boiling point of the solution? kb for water is change the molality of the particles in the solution?
0.52°C/m.
Solution: Label each of the following compounds as ionic or covalent. Then indicate the number of
ΔTb=kb·m·i particles formed when dissolved (i) for each compound.
W need to be able to substitute each variable into this equation. 5) Salt, NaCl
kb=0.52°C/m. 6) Acetone, C2H2O
m: We must solve for this using stoichiometry. Given: 10.0 g NaCl and 100. g H2O. 7) Benzene, C6H6
Find: mol NaCl/kg H2O. Ratios: molar mass of NaCl, 1000 g = 1 kg 8) Copper(II) nitrate, Cu(NO3)2
9) AlCl3
10) Potassium hydroxide, KOH
For NaCl, i = 2
Substitute these values into the equation ΔTb=kb·m·i. We get: For each pair of solutions, indicate which would have a lower freezing point:
11) 0.2 m KI or 0.2 m CaCl2
12) 0.1 m KI or 0.1 m C6H12O6
Water normally boils at 100°C, but our calculation shows that the boiling point increased 13) 0.2 m NaCl or 0.3 m C6H12O6
by 1.78°C. Our new boiling point is 101.78°C.
14) If 25.0g of sucrose, C12H22O11, is added to 500.g of water, the boiling point is
Lesson Summary increased by what amount? (Kb for water is 0.52°C/m)
Colligative properties are properties that are due only to the number of particles in
15) For a sample of seawater (an aqueous solution of NaCl), the concentration of salt is
solution and not related to the chemical properties of the solute.
Boiling points of solutions are higher that the boiling points of the pure solvents. approximately 0.50m. Calculate the freezing point of seawater. (K f for water is 1.86°C/m)
Freezing points of solutions are lower than the freezing points of the pure solvents. 16) Calcium chloride is known to melt ice faster than sodium chloride but is not used on
Ionic compounds split into ions when they dissolve, forming more particles.
roads, because the salt itself attracts water. If 15g of CaCl2 was added to 250g of
Covalent compounds stay as complete molecules when they dissolve.
water, what would be the new freezing point of the solution? Kf for water is 1.86°C/m
Vocabulary
Colligative property: a property that is due only to the number of particles in solution
and not the type of the solute
Boiling point elevation: the amount the boiling point of a solution increases from the
boiling point of a pure solvent
Freezing point depression: the amount the freezing point of a solution decreases from
the boiling point of a pure solvent

6.4: Review Questions

1) Which of the following statements are true when a solute is added to a solvent: (you
may choose more than 1)
a) the boiling point increases.
b) the boiling point decreases.
c) the freezing point increases.
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 Chapter 7: Describing Chemical Reactions
7.1: Chemical & Physical Change For the most part, physical changes tend to be reversible – in other words, they can
Objectives: occur in both directions. You can turn liquid water into solid water through cooling; you can
also turn solid water into liquid water through heating. However, as we will later learn, some
chemical changes can also be reversed.
Chemical changes occur when
Label a change as chemical or physical
bonds are broken and/or formed between
List evidence that can indicate a chemical change occurred
molecules or atoms. This means that one
Introduction substance with a certain set of properties
(such as melting point, color, taste, etc) is
Change is happening all around us all of the time. Just as chemists have classified
turned into a different substance with
elements and compounds, they have also classified types of changes. Changes are either
difference properties. Chemical changes
classified as physical or chemical changes.
are frequently harder to reverse than
Physical & Chemical Change physical changes.
One good example of a chemical Water forming from the elements hydrogen and oxygen
Chemists learn a lot about the nature of matter by studying the changes that matter
can undergo. Chemists make a distinction between two different types of changes that they change is burning paper. In contrast to the is a chemical change, because different substances are
act of ripping paper, the act of burning present at the end than the beginning.
study – physical changes and chemical changes. Physical changes are changes in which no
bonds are broken or formed. This means that the same types of compounds or elements that paper actually results in the formation of
were there at the beginning of the new chemicals (carbon dioxide and water, to be exact). Another example of chemical change
change are there at the end of the occurs when water is formed. Each molecule contains two atoms of hydrogen and one atom
of oxygen chemically bonded.
change. Because the ending materials
are the same as the beginning materials,
the properties (such as color, boiling
point, etc) will also be the same.
Physical changes involve moving
molecules around, but not changing
them. Some types of physical changes
include: Water changing from a liquid to gas (boiling) is a physical
Changes of state (changes from change, because H2O molecules are present at the
a solid to a liquid or a gas and beginning and end of the change. Natural gas burning is an example of a chemical change, because Firework displays are an example of a
vice versa) bonds are broken and formed to make different molecules. chemical change.
Separation of a mixture Another example of a chemical change is what occurs when natural gas is burned in
Physical deformation (cutting, denting, stretching) your furnace. This time, on the left we have a molecule of methane, CH4, and two molecules
Making solutions (special kinds of mixtures) of oxygen, O2, while on the right we have two molecules of water, H2O, and one molecule
When we heat the liquid water, it changes to water vapor. of carbon dioxide, CO2. In this case, not only has the appearance changed, but the structure
But even though the physical properties have changed, the of the molecules has also changed. The new substances do not have the same chemical
molecules are exactly the same as before. We still have each water properties as the original ones. Therefore, this is a chemical change.
molecule containing two hydrogen atoms and one oxygen atom
covalently bonded. When you have a jar containing a mixture of Evidence of Chemical Change
pennies and nickels and you sort the mixture so that you have one We can’t actually see molecules breaking and forming bonds, although that’s what
pile of pennies and another pile of nickels, you have not altered the defines chemical changes. We have to make other observations to indicate that a chemical
identity of either the pennies or the nickels – you’ve merely change has happened. Some of the evidence for chemical change will involve the energy
separated them into two groups. This would be an example of a changes that occur in chemical changes, but some evidence involves the fact that new
physical change. Similarly, if you have a piece of paper, you don’t substances with different properties are formed in a chemical change.
Melting snow is an change it into something other than a piece of paper by ripping it
example of a physical up. What was paper before you starting tearing is still paper when Observations that help to indicate chemical change include:
change. you’re done. Again, this is an example of a physical change. Temperature changes (either the temperature increases or decreases)

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 Light is given off
Unexpected color changes (a substance with a different color is made, rather than
just mixing the original colors together)
Bubbles are formed (but the substance is not boiling – you made a substance that is a 7.1: Review Questions
gas at the temperature as the beginning materials, instead of a liquid) Label each of the following as a physical or chemical change.
Different smell or taste (do not taste your chemistry experiments, though!) 1) Water boils at 100°C.
A solid forms if two clear liquids are mixed (look for floaties – technically called
a precipitate) 2) Water is separated by electrolysis
(running electricity through it) into
Example: Label each of the following changes as a physical or chemical change. Given hydrogen gas and oxygen gas. 8)
evidence to support your answer.
a) boiling water 3) Sugar dissolves in water.
b) a nail rusts
c) a green solution and colorless solution are mixed. The resulting mixture is a solution with 4) Vinegar and baking soda react to
a pale green color. produce a gas.
9)
d) two colorless solutions are mixed. The resulting mixture has a yellow precipitate
Solution: 5) Yeast acts on sugar to form
a) physical: boiling and melting are physical changes. When water boils no bonds are broken carbon dioxide and ethanol.
or formed. The change could be written: H2O(l) → H2O(g)
b) chemical: because the dark gray metal nail changes color to form an orange flakey 6) Wood burns producing several new 10)
substance (the rust) this must be a chemical change. Color changes indicate color change. substances.
The following reaction occurs: Fe + O2 → Fe2O3
c) physical: because none of the properties changed, this is a physical change. The green 7) A cake is baked.
mixture is still green and the colorless solution is still colorless. They have just been spread
together. No color change occurred or other evidence of chemical change. 7.2: Reaction Rate
d) chemical: the formation of a precipitate and the color change from colorless to yellow Objectives
indicates a chemical change. Describe the conditions for successful collisions that cause reactions
Describe the rate in terms of the conditions of successful collisions.
Lesson Summary Describe how changing the temperature, concentration of a reactant, or surface
Chemists make a distinction between two different types of changes that they study – area of a reaction affects the rate of a reaction
physical changes and chemical changes. Define a catalyst and how a catalyst affects the rate of a reaction
Physical changes are changes that do not alter the identity of a substance
Chemical changes are changes that occur when one substance is turned into another Introduction
-
substance. We know that a chemical system can be made up of atoms (H 2, N2, K, etc), ions (NO3 ,
- +
Chemical changes are frequently harder to reverse than physical changes. Cl , Na , etc), or molecules (H2O, C12H22O11, etc). We also know that in a chemical system,
Observations that indicate a chemical change occurred include color change, these particles are moving around in a random motion. The collision theory explains why
temperature change, light given off, formation of bubbles, formation of a precipitate, etc reactions occur at this particle level between these atoms, ions, and/or molecules. It also explains
how it is possible to speed up or slow down reactions that are occurring.

Vocabulary Collision Theory

The collision theory provides us with the ability to predict what conditions are necessary
Physical changes: changes that do not alter the identity of a substance, the same types for a successful reaction to take place. These conditions include:
of molecules are present at the beginning and end of the change. 1. The particles must collide with each other.
Chemical changes: changes that occur when one substance is turned into another 2. The particles must collide with sufficient energy to break the old bonds.
substance; different types of molecules are present at the beginning and end of 3. The particles must have proper orientation.
the change.

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 A chemical reaction involves breaking bonds in the reactants, re-arranging the
atoms into new groupings (the products), and the formation of new bonds in the products.

Effect of Temperature on Rate of Reaction

The rate of reaction was discussed in terms of three factors: collision frequency, the
collision energy, and the geometric orientation. Remember that the collision frequency is
the number of collisions per second. The collision frequency is dependent, among other
factors, on the temperature of the reaction.
When the temperature is increased, the average velocity of the particles is increased.
The average kinetic energy of these particles is also increased. The result is that the particles
This collision is successful and results in reaction. (Source: Richard Parsons. CC-BY-SA) will collide more frequently, because the particles move around faster and will encounter
Therefore, not only must a collision occur between reactant particles, but the collision has to more reactant particles. However, this is only a minor part of the reason why the rate is
have sufficient energy to break all the reactant bonds that need to be broken in order to form increased. Just because the particles are colliding more frequently does not mean that the
the products. Some reactions need less collision energy than others. The amount of energy reaction will definitely occur.
the reactant particles must have in order to break the old bonds for a reaction to occur is The major effect of increasing the temperature is that more of the particles that collide
called the activation energy, abbreviated Ea. Another way to think of this is to look at an will have the amount of energy needed to have an effective collision. In other words, more
energy diagram, as particles will have the necessary activation energy.
shown in the figure. At room temperature, the hydrogen and oxygen in the atmosphere do not
Particles must be have sufficient energy to attain the activation energy needed to produce water:
able to get over the O2(g) + H2(g) No reaction
“bump”, the At any one moment in the atmosphere, there are many collisions occurring between these
activation energy, if two reactants. But what we find is that water is not formed from the oxygen and hydrogen
they are going to molecules colliding in the atmosphere because the activation energy barrier is just too high
react. If the reactant and all the collisions are resulting in rebound. When we increase the temperature of the
particles collide with reactants or give them energy in some other way, the molecules have the necessary activation
less than the Each reaction has its own activation energy, Ea. The smaller the “bump”, energy and are able to react to produce water:
activation energy, the the less energy particles must have to react. O2(g) + H2(g) H2O(l)
particles will rebound There are times when the rate of a reaction needs to be slowed down. Lowering the
(bounce off each other), and no reaction will occur. temperature could also be used to decrease the number of collisions that would occur and
lowering the temperature would also reduce the kinetic energy available for activation
Reaction Rate energy. If the particles have insufficient activation energy, the collisions will result in
Chemists use reactions to generate a product for which they have use. For the most rebound rather than reaction. Using this idea, when the rate of a reaction needs to be
part, the reactions that produce some desired compound are only useful if the reaction occurs lower, keeping the particles from having sufficient activation energy will definitely keep
at a reasonable rate. For example, using a reaction to produce brake fluid would not be useful the reaction at a lower rate.
if the reaction required 8,000 years complete the product. Such a reaction would also not be Society uses the effect of temperature on rate every day. Food storage is a prime
useful if the reaction was so fast that it was explosive. For these reasons, chemists wish to be example of how the temperature effect on reaction rate is utilized by society. Consumers
able to control reaction rates. In some cases, chemists wish to speed up reactions that are too store food in freezers and refrigerators to slow down the processes that cause it to spoil. The
slow and slow down reactions that are too fast. In order to gain any control over reaction decrease in temperature decreases the rate at which the food will break down or be broken
th
rates, we must know the factors that affect reaction rates. Chemists have identified many down by bacteria. In the early years of the 20 century, explorers were fascinated with
factors that affect the rate of a reaction. trying to be the first one to reach the South Pole. In order to attempt such a difficult task at a
The rate, or speed, at which a reaction occurs depends on the frequency of time without most of the technology we take for granted today, they devised a variety of
successful collisions. Remember, a successful collision occurs when two reactants collide ways of surviving. One method was to store their food in the snow to be used later during
with enough energy and with the right orientation. That means if we can do things that will their advances to the pole. On some explorations, they buried so much food, that they didn’t
increase the number of collisions, increase the number of particles that have enough energy need to use all of it and it was left. Many years later, when this food was located and
to react and/or increase the number of particles with the correct orientation we will increase thawed, it was found to still be edible.
the rate of a reaction. When milk, for instance, is stored in the refrigerator, the molecules in the milk have
less energy. This means that while molecules will still collide with other molecules, few of
138 them will react (which means in this case “spoil”) because the molecules do not have
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sufficient energy to overcome the activation energy barrier. The molecules do have energy
and are colliding, however, and so, over time, even in the refrigerator, the milk will spoil.
Eventually the higher energy molecules will gain the energy needed to react and when
will continue to glow in air for a period of time. If we insert that glowing splint into any gas
enough of these reactions occur, the milk becomes “soured”.
that does not contain oxygen, the splint will immediately cease to glow - that is the reaction
However, if that same carton of milk was at room temperature, the milk would react stops. Oxygen is the only gas that will support combustion. Air is approximately, 20%
(in other words “spoil’) much more quickly. Now most of the molecules will have sufficient oxygen gas. If we take that glowing splint and insert it into pure oxygen gas, the reaction
energy to overcome the energy barrier and at room temperature many more collisions will be will increase its rate by a factor of five - since pure oxygen has 5 times the concentration of
occurring. This allows for the milk to spoil in a fairly short amount of time. This is also the
oxygen that is in air. When the reaction occurring on the glowing splint increases its rate by
reason why most fruits and vegetables ripen in the summer when the temperature is much
a factor of five, the glowing splint will suddenly burst back into full flame. This test, of
warmer. You may have experienced this first hand if you have ever bitten into an unripe
thrusting a glowing splint into a gas, is used to identify the gas as oxygen. Only a greater
banana – it was probably sour tasting and might even have felt like biting into a piece of
concentration of oxygen than that found in air will cause the glowing splint to burst into
wood! When a banana ripens, numerous reactions occur that produce all the compounds that
flame.
we expect to taste in a banana. But this can only happen if the temperature is high enough to
allow these reactions to make those products. Effect of Surface Area on Rate of Reaction
Effect of Concentration on Rate of Reaction The very first requirement for a reaction to occur between reactant particles is that
the particles must collide with each other. The previous section pointed out how increasing
If you had an enclosed space, like a classroom, and there was one red ball and one
the concentration of the reactants increases reaction rate because it increased the frequency
green ball flying around the room with random motion and undergoing perfectly elastic
of collisions between reactant particles. It can be shown that the number of collisions that
collisions with the walls and with each other, in a given amount of time, the balls would
occur between reactant particles is also dependent on the surface area of solid reactants.
collide with each other a certain number of times determined by probability. If you now put
Consider a reaction between reactant RED and reactant BLUE in which reactant blue is in
two red balls and one green ball in the room under the same conditions, the probability of a
the form of a single lump. Then compare this to the same reaction where reactant blue has
collision between a red ball and the green ball would exactly double. The green ball would
been broken up into many smaller pieces.
have twice the chance of encountering a red ball in the same amount of time.
In the diagram, only the blue
particles on the outside surface of
the lump are available for collision
In terms of chemical reactions, with reactant red. The blue particles
a similar situation exists. Particles of on the interior of the lump are
two gaseous reactants or two reactants protected by the blue particles on the
in solution have a certain probability of surface. In Figure A, if you count the In these figures, only the particles on the outside of the
undergoing collisions with each other number of blue particles available solid blue reactant have a chance to collide with the red
in a reaction vessel. If you double the for collision, you will find that only reactant. In figure B, the same amount of solid reactant as
concentration of either reactant, the 20 blue particles could be struck by a used in A was crushed into smaller particles. This means
probability of a collision doubles. The that more particles on the outside of the reactant have an
particle of reactant red. In Figure A,
opportunity to collide with the red reactant and speeds up
there are a number of blue particles
the reaction.
rate of reaction is proportional to the The reaction mixture on the left is less concentrated, so on the interior of the lump that cannot be struck. In Figure B, however, the lump has been
particles will not collide as often. The reaction will be
number of collisions per unit time. If broken up into smaller pieces and all the interior blue particles are now on a surface and
slower.
one concentration of one of the Because the reacting particles on the right image have a available for collision. In diagram B, more collisions between blue and red will occur, and
reactants is doubled, the number of greater change of colliding, the reaction will go faster. therefore, the reaction in Figure B will occur at faster rate than the same reaction in Figure A.
collisions will also double. Assuming CC – Tracy Poulsen Increasing the surface area of a reactant increases the frequency of collisions and increases
that the percent of collision that are successful does not change, then having twice as many the reaction rate.
collisions will result in twice as many successful collisions. The rate of reaction is Several smaller particles have more surface area than one large particle. The more
proportional to the number of collisions over time and increasing the concentration of either surface area that is available for particles to collide, the faster the reaction will occur. You can
reactant increases the number of collisions and therefore, increases the number of see an example of this in everyday life if you have ever tried to start a fire in the fireplace. If you
successful collisions and the reaction rate. hold a match up against a large log in an attempt to start the log burning, you will find it to be an
For example, the chemical test used to identify a gas as oxygen or not relies on the fact unsuccessful effort. Holding a match against a large log will not cause enough reactions to occur
that increasing the concentration of a reactant increases reaction rate. The reaction we call in order to keep the fire going by providing sufficient activation energy for further reactions. In
combustion refers to a reaction in which a flammable substance reacts with oxygen. If we light a order to start a wood fire, it is common to break a log up into many small, thin sticks called
wooden splint (a thin splinter of wood) on fire and then blow the fire out, the splint kindling. These thinner sticks of wood provide many times the

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surface area of a single log. The match will successfully cause enough reactions in the kindling
so that sufficient heat is given off to provide activation energy for further reactions.
There have been, unfortunately, cases where serious accidents were caused by the In the figure on the
failure to understand the relationship between surface area and reaction rate. One such
left, the endothermic
example occurred in flour mills. A grain of wheat is not very flammable. It takes a
reaction shows the catalyst
significant effort to get a grain of wheat to burn. If the grain of wheat, however, is pulverized
reaction in red with the
and scattered through the air, only a spark is necessary to cause an explosion. When the
lower activation energy,
wheat is ground to make flour, it is pulverized into a fine powder and some of the powder
designated E’a. The new
gets scattered around in the air. A small spark then, is sufficient to start a very rapid reaction
reaction pathway has lower
which can destroy the entire flour mill. In a 10-year period from 1988 to 1998, there were
activation energy but has no
129 grain dust explosions in mills in the United States. Efforts are now made in flour mills to
effect on the energy of the
have huge fans circulate the air in the mill through filters to remove the majority of the flour A catalyst speeds up a reaction by lowering the activation energy,
reactants, the products, or
Ea,. Because less energy is required to react, more particles have the
the value of ΔH. The same
necessary energy.
is true for the exothermic
dust particles. reaction in Right Figure. The activation energy of the catalyzed reaction is lower than that
Another example is in the operation of coal mines. Coal, of course, will burn but it of the uncatalyzed reaction. The new reaction pathway provided by the catalyst affects the
takes an effort to get the coal started and once it is burning, it burns slowly because only the energy required for reactant bonds to break and product bonds to form.
surface particles are available to collide with oxygen particles. The interior particles of coal While many reactions in the laboratory can be increased by increasing the
have to wait until the outer surface of the coal lump burns off before they can collide with temperature, that is not possible for all the reactions that occur in our bodies throughout our
oxygen. In coal mines, huge blocks of coal must be broken up before the coal can be entire lives. In fact, the body needs to be maintained at a very specific temperature: 98.6°F or
brought out of the mine. In the process of breaking up the huge blocks of coal, drills are 37°C. Of course there are times, for instance, when the body is fighting an infection, when
used to drill into the walls of coal. This drilling produces fine coal dust that mixes into the the body temperature may be increased. But generally, in a healthy person, the temperature
air and then a spark from a tool can cause a massive explosion in the mine. There are is quite consistent. However, many of the reactions that a healthy body depends on could
explosions in coal mines for other reasons but coal dust explosions contributed to the death never occur at body temperature. The answer to this dilemma is catalysts or what are also
of many miners. In modern coal mines, lawn sprinklers are used to spray water through the referred to as enzymes. Many of these enzymes are made in your cells since your DNA
air in the mine and this reduces the coal dust in the air and eliminates coal dust explosions. carries the directions to make them. However, there are some enzymes that your body must
have but are not made in your cells. These catalysts must be supplied to your body in the
Effect of a Catalyst on Rate of Reaction food you eat and are called vitamins.
The final factor that affects the rate of the reaction is the effect of the catalyst. A
catalyst is a substance that speeds up the rate of the reaction without itself being Lesson Summary
consumed by the reaction. The collision theory explains why reactions occur between atoms, ions, and/or
In the reaction of potassium chlorate breaking down to potassium chloride and molecules
oxygen, a catalyst is available to make this reaction occur much faster than it would occur by In order for a reaction to be effective, particles must collide with enough energy
itself under room conditions. The reaction is: and having the correct orientation.
With an increase in temperature, there is an increase in energy that can be converted
The catalyst is manganese dioxide and its presence causes the reaction shown above to run into activation energy in a collision and therefore there will be an increase in the
many times faster than it occurs without the catalyst. When the reaction has reached reaction rate. A decrease in temperature would have the opposite effect.
completion, the MnO2 can be removed from the reaction vessel and its condition is exactly With an increase in temperature there is an increase in the number of collisions.
the same as it was before the reaction. This is part of the definition of a catalyst . . . that it Increasing the concentration of a reactant increases the frequency of collisions
is not consumed by the reaction. You should note that the catalyst is not written into the between reactants and will, therefore, increase the reaction rate.
equation as a reactant or product but is noted above the yields arrow. This is standard Increasing the surface area of a reactant (by breaking a solid reactant into smaller
notation for the use of a catalyst. particles) increases the number of particles available for collision and will
Some reactions occur very slowly without the presence of a catalyst. In other words increase the number of collisions between reactants per unit time.
the activation energy for these reactions is very high. When the catalyst is added, the The catalyst is a substance that speeds up the rate of the reaction without itself being
activation energy is lowered because the catalyst provides a new reaction pathway with lower consumed by the reaction. When the catalyst is added, the activation energy is
activation energy. lowered because the catalyst provides a new reaction pathway with lower activation
energy.

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Vocabulary
Catalyst: A substance that speeds up the rate of the reaction without itself being
consumed by the reaction a) The kinetic energy increases.
Surface area to volume ratio: The comparison of the volume inside a solid to the area b) The activation energy increases.
exposed on the surface. c) The number of successful collisions increases.
d) All of the above.
Further Readings / Supplemental Links
Activation Energy: Choose the substance with the greatest surface in the following groupings:
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/activa2.swf 6) a block of ice or crushed ice
http://learner.org/resources/series61.html The learner.org website allows users to 7) sugar crystals or sugar cubes
view streaming videos of the Annenberg series of chemistry videos. You are 8) a piece of wood or wood shavings
required to register before you can watch the videos but there is no charge. The
website has one video that relates to this lesson called Molecules in Action. 9) Why, using the collision theory, do reactions with higher surface area have faster reaction
http://www.vitamins-guide.net rates?
http://en.wikipedia.org/wiki
Observing molecules during chemical reactions helps explain the role of catalysts. Limestone (calcium carbonate) reacts with hydrochloric acid in an irreversible reaction,
Dynamic equilibrium is also demonstrated. Molecules in Action to form carbon dioxide and water as described by the following equation:
(http://www.learner.org/vod/vod_window.html?pid=806) CaCO3(s) + 2 HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)
Surface science examines how surfaces react with each other at the molecular level. What is the effect on the rate if:
On the Surface (http://www.learner.org/vod/vod_window.html?pid=812) 10) The temperature is lowered?
http://en.wikipedia.org/wiki 11) Limestone chips are used instead of a block of limestone?
12) A more dilute solution of HCl is used?
7.2: Review Questions
Multiple Choice 7.3: Chemical Reactions and Equations
1) According to the collision theory, what must happen in order for a reaction to Objectives
be successful? (Select all that apply.) Identify the reactants and products in any chemical reaction.
a) particles must collide Convert word equations into chemical equations.
b) particles must have proper geometric orientation Use the common symbols, (s), (l), (g), (aq), and , appropriately when writing a
c) particles must have collisions with enough energy chemical reaction
2) What would happen in a collision between two particles if particles did not have enough Explain the roles of subscripts and coefficients in chemical equations.
energy or had the incorrect orientation? Balance a chemical equation when given the unbalanced equation.
a) the particles would rebound and there would be no reaction Explain the role of the Law of Conservation of Mass in a chemical reaction.
b) the particles would keep bouncing off each other until they eventually react,
therefore the rate would be slow Introduction
c) the particles would still collide but the byproducts would form In a chemical change, new substances are formed. In order for this to occur, the
d) the temperature of the reaction vessel would increase chemical bonds of the substances break, and the atoms that compose them separate and re-
3) Why does higher temperature increase the reaction rate? arrange themselves into new substances with new chemical bonds. When this process occurs,
a) more of the reacting molecules will have higher kinetic energy we call it a chemical reaction. A chemical reaction is the process in which one or more
b) increasing the temperature causes the reactant molecules to heat up substances are changed into one or more new substances.
c) the activation energy will decrease
4) When the temperature is increased, what does not change? Reactants and Products
a) number of collisions In order to describe a chemical reaction, we need to indicate what substances are
b) activation energy present at the beginning and what substances are present at the end. The substances that are
c) number of successful collisions present at the beginning are called reactants and the substances present at the end are
d) all of the above change called products.
5) Why is the increase in concentration directly proportional to the rate of the reaction? Sometimes when reactants are put into a reaction vessel, a reaction will take place to
produce products. Reactants are the starting materials, that is, whatever we have as our initial
144 ingredients. The products are just that, what is produced or the result of what happens to the
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reactants when we put them together in the reaction vessel. If we think about baking
chocolate chip cookies, our reactants would be flour, butter, sugar, vanilla, some baking
soda, salt, egg, and chocolate chips. What would be the products? Cookies! The reaction
3. They could write the equation in chemical shorthand.
vessel would be our mixing bowl.
Flour + Butter + Sugar + Vanilla + Baking Soda + Eggs + Chocolate Chips Cookies 2 H2(g) + O2(g) → 2 H2O(g)
In the symbolic equation, chemical formulas are used instead of chemical names for
Writing Chemical Equations reactants and products and symbols are used to indicate the phase of each substance. It
When sulfur dioxide is added to oxygen, sulfur trioxide is produced. Sulfur dioxide should be apparent that the chemical shorthand method is the quickest and clearest method
and oxygen, SO2 + O2, are reactants and sulfur trioxide, SO3, is the product. for writing chemical equations.
We could write that an aqueous solution of calcium nitrate is added to an aqueous
solution of sodium hydroxide to produce solid calcium hydroxide and an aqueous solution
of sodium nitrate. Or in shorthand we could write:
Ca(NO3)2(aq) + 2 NaOH(aq) → Ca(OH)2(s) + 2 NaNO3(aq)
Reactants → Products How much easier is that to read? Let's try it in reverse? Look at the following reaction
In chemical reactions, the reactants are found before the symbol “→” and the products and in shorthand notation and write the word equation for the reaction.
found after the symbol “→”. The general equation for a reaction is: Cu(s) + AgNO3(aq) → Cu(NO3)2(aq) + Ag(s)
Reactants → Products The word equation for this reaction might read something like "solid copper reacts with an
There are a few special symbols that we need to know in order to “talk” in chemical aqueous solution of silver nitrate to produce a solution of copper (II) nitrate with solid
shorthand. In the table below is the summary of the major symbols used in chemical equations. silver".
You will find there are others but these are the main ones that we need to know. In order to turn word equations into symbolic equations, we need to follow the given
Common Symbols in Chemical Reactions steps:
Symbol Meaning Example 1. Identify the reactants and products. This will help you know what symbols go on
each side of the arrow and where the + signs go.
Used to separate reactants from products;
2. Write the correct formulas for all compounds. You
can be read as "to produce" or "yields. 2H2+O2→2H2O will need to use the rules you learned in chapter 4 Diatomic Elements
Used to separate reactants from each (including making all ionic compounds charge Element name Formula
+ other or products from each other; can be AgNO3 + NaCl → AgCl + NaNO3 balanced). Hydrogen H2
read as "is added to" or “also forms”. 3. Write the correct formulas for all elements. Usually, Nitrogen N2
(s) in the solid state sodium in the solid state = Na(s) this is given straight off of the periodic table. Oxygen O2
(l) in the liquid state water in the liquid state = H2O(l) However, there are seven elements that are considered Fluorine F2
carbon dioxide in the gaseous state diatomic, meaning they are always found in pairs in Chlorine Cl2
(g) in the gaseous state nature. They include those elements listed in the table. Bromine Br
= CO2(g) 2

sodium chloride solution = Iodine I2

(aq) in the aqueous state, dissolved in water NaCl(aq)
Example: Transfer the following symbolic equations into word equations or word
Chemists have a choice of methods for describing a chemical reaction. equations into symbolic equations.
1. They could draw a picture of the chemical reaction.
(a) HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
(b) Gaseous propane, C3H8, burns in oxygen gas to produce gaseous carbon dioxide
and liquid water.
(c) Hydrogen fluoride gas reacts with an aqueous solution of potassium carbonate to produce an
aqueous solution of potassium fluoride, liquid water, and gaseous carbon dioxide.
Solution:
(a) An aqueous solution of hydrochloric acid reacts with an aqueous solution of
sodium hydroxide to produce an aqueous solution of sodium chloride and liquid water.
2. They could write a word equation for the chemical reaction: (b) Reactants: propane (C3H8) and oxygen (O2)
“Two molecules of hydrogen gas react with one molecule of oxygen gas to produce Products: carbon dioxide (CO2) and water (H2O)
two molecules of water vapor. “ C3H8(g) + O2(g) → CO2(g) + H2O(l)
(c) Reactants: hydrogen fluoride and potassium carbonate
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Products: potassium fluoride, water, and carbon dioxide
HF(g) + K2CO3(aq) → KF(aq) + H2O(l) + CO2(g)
Once you have written a symbolic equation from words, it is important to balance the
7.3: Review Questions
equation. It is very important to note that these steps must be carried out in the correct order.
Convert the following equations from word equations into symbolic equations. Be sure You must have the correct formulas for your reactants and products before you can use
to look up charges of ionic compounds to write the correct formula for the compound. coefficients to balance the equation.
1) Solid calcium metal is placed in liquid water to produce aqueous calcium hydroxide
When you learned how to write formulas, it was made clear that when only one atom
and hydrogen gas.
of an element is present, the subscript of "1" is not written - so that when no subscript
2) Gaseous sodium hydroxide is mixed with gaseous chlorine to produce aqueous solutions
appears for an atom in a formula, you read that as one atom. The same is true in writing
of sodium chloride and sodium hypochlorite plus liquid water.
balanced chemical equations. If only one atom or molecule is present, the coefficient of "1" is
3) Iron reacts with sulfur when heated to form iron(II) sulfide.
omitted. Coefficients are inserted into the chemical equation in order to balance it; that is, to
4) When aluminum is added to sulfuric acid (H2SO4), the solution reacts to form make the total number of each atom on the two sides of the equation equal. Equation
hydrogen gas and aluminum sulfate.
balancing is accomplished by changing coefficients, never by changing subscripts.
5) When aluminum is mixed with iron(III) oxide, they react to produce aluminum oxide
and iron.
Example: Balance the following skeletal equation. (The term "skeletal equation" refers to an
6) Fluorine is mixed with sodium hydroxide to form sodium fluoride, oxygen, and water.
equation that has the correct formulas but has not yet had the proper coefficients added.)
7) A solid chunk of iron is dropped into a solution of copper(I) nitrate forming iron(II)
nitrate and solid copper. Fe(NO3)3 + NaOH → Fe(OH)3 + NaNO3
Solution: We can balance the hydroxide ion by inserting a coefficient of 3 in front of
the NaOH on the reactant side.
Fe(NO3)3 + 3 NaOH → Fe(OH)3 + NaNO3
7.4: Balancing Chemical Equations Then we can balance the nitrate ions by inserting a coefficient of 3 in front of the
Even though chemical compounds are broken up and new compounds are formed sodium nitrate on the product side.
during a chemical reaction, atoms in the reactants do not disappear nor do new atoms appear Fe(NO3)3 + 3 NaOH → Fe(OH)3 + 3 NaNO3
to form the products. In chemical reactions, atoms are never created or destroyed. The same Counting the number of each type of atom on the two sides of the equation will now show
atoms that were present in the reactants are present in the products – they are merely re- that this equation is balanced.
organized into different arrangements. In a complete chemical equation, the two sides of the
equation must be balanced. That is, in a complete chemical equation, the same number of Example: Write a balanced equation for the reaction that occurs between chlorine gas
each atom must be present on the reactants and the products sides of the equation. and aqueous sodium bromide to produce liquid bromine and aqueous sodium chloride.
There are two types of numbers that appear in chemical equations. There are
Solution:
subscripts, which are part of the chemical formulas of the reactants and products and
Step 1: Write the word equation (keeping in mind that chlorine and bromine refer to the
there are coefficients that are placed in front of the formulas to indicate how many
diatomic molecules).
molecules of that substance is used or produced.
Chlorine + sodium bromide → bromine + sodium chloride
Step 2: Substitute the correct formulas into the equation.
Cl2 + NaBl → Br2 + NaCl
Step 3: Insert coefficients where necessary to balance the equation.
By placing a coefficient of 2 in front of the NaBr, we can balance the bromine atoms and
by placing a coefficient of 2 in front of the NaCl, we can balance the chlorine atoms.
Cl2 + 2 NaBl → Br2 + 2 NaCl
The subscripts are part of the formulas and once the formulas for the reactants and products A final check (always do this) shows that we have the same number of each atom on the two
are determined, the subscripts may not be changed. The coefficients indicate the number of
sides of the equation and we do not have a multiple set of coefficients so this equation is
each substance involved in the reaction and may be changed in order to balance the
properly balanced.
equation. The equation above indicates that one mole of solid copper is reacting with two
moles of aqueous silver nitrate to produce one mole of aqueous copper (II) nitrate and two
Example: Write a balanced equation for the reaction between aluminum sulfate and
atoms of solid silver.
calcium bromide to produce aluminum bromide and calcium sulfate. (You may need to refer
to a chart of polyatomic ions.)
148 Solution:
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 Aluminum sulfate + calcium bromide → aluminum bromide + calcium sulfate
Step 2: Replace the names of the substances in the word equation with formulas.
Al2(SO4)3 + CaBr2 → AlBr3 + CaSO4 Conservation of Mass in Chemical Reactions
Step 3: Insert coefficients to balance the equation.
In chapter 2 we discussed the development of the atomic theory, or the idea that
In order to balance the aluminum atoms, we must insert a coefficient of 2 in front of
everything is made of atoms. A strong piece of evidence for this theory was experimentally
the aluminum compound in the products.
determined by Antoine Lavoisier, a French chemist. The Law of Conservation of Mass, as
Al2(SO4)3 + CaBr2 → 2 AlBr3 + CaSO4
he states it, says that mass is conserved in chemical reactions. In other words, the mass of
In order to balance the sulfate ions, we must insert a coefficient of 3 in front of the CaSO4 in
the starting materials (reactants) is always equal to the
the products. mass of the ending materials (products).
Al2(SO4)3 + CaBr2 → 2 AlBr3 + 3 CaSO4 But what does this really mean? Dalton used this
In order to balance the bromine atoms, we must insert a coefficient of 3 in front of the
finding to support the idea of atoms. If the mass isn’t
CaBr2 in the reactants.
changing, then the particles that carry the mass (atoms)
Al2(SO4)3 + 3 CaBr2 → 2 AlBr3 + 3 CaSO4 aren’t created or destroyed, but are only rearranged in a
The insertion of the 3 in front of the CaBr2 in the reactants also balances the calcium atoms in chemical reaction. Both the numbers of each type of atom
the CaSO4 in the products. A final check shows 2 aluminum atoms on each side, 3 sulfur and the mass are conserved during chemical reactions. An
atoms on each side, 12 oxygen atoms on each side, 3 calcium atoms on each side, and examination of a properly balanced equation will
6 bromine atoms on each side. This equation is balanced.
demonstrate that mass is conserved. Consider the
following reaction.
Chemical equations should be balanced with the simplest whole number coefficients
Fe(NO3)3 + 3 NaOH → Fe(OH)3 + 3 NaNO3
that balance the equation. Here is the properly balanced equation from the previous section. You should check that this equation is balanced by
Al2(SO4)3 + 3 CaBr2 → 2 AlBr3 + 3 CaSO4 counting the number of each type of atom on each side of
Note that the equation in the previous section would have the same number of atoms of each the equation.
type on each side of the equation with the following set of coefficients. We can also demonstrate that mass is conserved in
2 Al2(SO4)3 + 6 CaBr2 → 4 AlBr3 + 6 CaSO4 this reaction by determining the total mass on the two sides Obtained from:
Count the number of each type of atom on each side of the equation to confirm that this http://upload.wikimedia.org/wikiped
of the equation. We will use the molar masses to add up ia/commons/7/78/Antoine_laurent_l
equation is "balanced". While this set of coefficients does "balance" the equation, they are the masses of the atoms on the reactant side and compare avoisier.jpg
not the lowest set of coefficients possible that balance the equation. We could divide each of this to the mass of the atoms on the product side of the
the coefficients in this equation by 2 and get another set of coefficients that are whole reaction:
numbers and also balance the equation. Since it is required that an equation be balanced with
the lowest whole number coefficients, the last equation is NOT properly balanced. When Reactant Side Mass
you have finished balancing an equation, you should not only check to make sure it is 1 moles of Fe(NO3)3 x molar mass = (1mol)(241.9 g/mol) = 241.9 g
3 moles of NaOH x molar mass = (3mol)(40.0 g/mol) = 120. g
balanced, you should also check to make sure that it is balanced with the simplest set of
whole number coefficients possible. Total mass for reactants = 241.9 g + 120. g = 361.9 g
Product Side Mass
Example: Balance each of the following reactions. 1 moles of Fe(OH)3 x molar mass= (1mol)(106.9g/mol) = 106.9g
3 moles of NaNO3 x molar mass = (3mol)(85.0 g/mol) = 255 g
(a) CaCO3(s) → CaO(s) + CO2(g)
Total mass for products = 106.9 g + 255 g = 361.9 g
(b) H2SO4(aq) + Al(OH)3(aq) → Al2(SO4)3(aq) + H2O(l)
(c) Ba(NO3)2(aq) + Na2CO3(aq) → BaCO3(aq) + NaNO3(aq) As you can see, both number of atom types and mass are conserved during chemical
(d) C2H4(g) + O2 → CO2(g) + H2O(l) reactions. A group of 20 objects stacked in different ways will still have the same total
Solutions
mass no matter how you stack them.
(a) CaCO3(s) → CaO(s) + CO2(g) (In this case, the equation balances with all
coefficients being 1) Lesson Summary
(b) 3 H2SO4(aq) + 2 Al(OH)3(aq) → Al2(SO4)3(aq) + 6 H2O(l) A chemical reaction is the process in which one or more substances are changed
(c) Ba(NO3)2(aq) + Na2CO3(aq) → BaCO3(aq) + 2 NaNO3(aq) into one or more new substances.
(d) C2H4(g) + 3 O2 → 2 CO2(g) + 2 H2O(l) Chemical reactions are represented by chemical equations.
Chemical equations have reactants on the left, an arrow that is read as "yields,"
and the products on the right.
150 To be useful, chemical equations must always be balanced.
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 Balanced chemical equations have the same number and type of each atom on both
sides of the equation.
The coefficients in a balanced equation must be the simplest whole number ratio.
13) H3PO4 + NH4OH → HOH + (NH4)3PO4
Mass is always conserved in chemical reactions.
14) C3H8 + O2 → CO2 + H2O
Vocabulary 15) Al + O2 → Al2O3
16) CH4 + O2 → CO2 + H2O
Chemical reaction: the process in which one or more substances are changed into one
or more new substances 17) When the following equation is balanced, what is the coefficient found in front of the
Reactants: the starting materials in a reaction O2? P4 + O2 + H2O → H3PO4
Products: materials present at the end of a reaction 18) When properly balanced, what is the sum of all the coefficients in the following chemical
Balanced chemical equation: a chemical equation in which the number of each equation? SF4 + H2O → H2SO3 + HF
type of atom is equal on the two sides of the equation 19) Explain in your own words why it is essential that subscripts remain constant but
Subscripts: part of the chemical formulas of the reactants and products that indicate coefficients can change when balancing a reaction.
the number of atoms of the preceding element
Coefficient: a small whole number that appears in front of a formula in a balanced 7.5: Types of Reactions
chemical equation Objectives
Classify a chemical reaction as a synthesis, decomposition, single replacement,
Further Reading / Supplemental Links double replacement, or a combustion reaction. Predict the products of simple
For a Bill Nye video on reactions, go to http://www.uen.org/dms/. Go to the k-12 reactions.
library. Search for “Bill Nye reactions”. (you can get the username and password
from your teacher) Introduction
Chemical reactions are classified into types to help us analyze them and also to
For videos and clips on reactions, go to http://www.uen.org/dms/. Go to the k-12
help us predict what the products of the reaction will be. The five major types of chemical
library. Search for “reactions” or “chemical equations”. (you can get the username reactions are synthesis, decomposition, single replacement, double replacement, and
and password from your teacher) combustion.

Vision Learning: Chemical Equations Synthesis Reactions

http://visionlearning.com/library/module_viewer.php?mid=56&l=&c3= A synthesis reaction is one in which two or
Balancing Equations Tutorial: more reactants combine to make one type of product.
http://www.mpcfaculty.net/mark_bishop/balancing_equations_tutorial.htm General equation: A + B → AB
Balancing Equations Tutorial: http://www.wfu.edu/~ylwong/balanceeq/balanceq.html Synthesis reactions occur as a result of two or more
simpler elements or molecules combining to form a
Law of Conservation of Mass (YouTube): http://www.youtube.com/watch%3Fv
more complex molecule. Look at the example below.
%3DdExpJAECSL8
Here two elements (hydrogen and oxygen) are
combining to form one product (water). A synthesis reaction is similar to
7.4 Review Questions forming a couple, which behaves
Example: 2 H2(g) + O2(g) → 2 H2O(l)
Balance the following equations. We can always identify a synthesis reaction because
and acts differently than the two
1) Cu + O2 → CuO single individuals.
2) H2O →H2 + O2 there is only one product of the reaction. CC – Tracy Poulsen
3) Fe + H2O → H2 + Fe2O3 You should be able to write the chemical
4) NaCl → Na + Cl2 equation for a synthesis reaction if you are given a product by picking out its elements and
5) AsCl3 + H2S → As2S3 + HCl writing the equation. Also, if you are given elemental reactants and told that the reaction is a
6) CaCO3 → CaO + CO2 synthesis reaction, you should be able to predict the products.
7) H2S + KOH → HOH + K2S Example:
8) XeF6 + H2O → XeO3 + HF (a) Write the chemical equation for the synthesis reaction of silver bromide, AgBr.
9) Cu + AgNO3 → Ag + Cu(NO3)2 (b) Predict the products for the following reaction: CO2(g) + H2O(l)
10) Fe + O2 → Fe2O3 Solution:
11) Al(OH)3 + Mg3(PO4)2 → AlPO4 + Mg(OH)2 (a) 2 Ag + Br2 → 2 AgBr
12) Al + H2SO4 → H2 + Al2(SO4)3 (b) CO2(g) + H2O(l) → H2CO3

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Decomposition Reactions
Notice that the metal element, Zn, replaced the metal in the compound Cu(NO3)2. A metal
When one type of reactant breaks down to form two or more products, we have a
element will always replace a metal in an ionic compound. Also, note that the charges of the
decomposition reaction. The best way to remember a decomposition reaction is that for all
ionic compounds must equal zero. To correctly predict the formula of the ionic product, you
reactions of this type, there is only one reactant.
must know the charges of the ions you are combining.
General Equation: AB → A + B
Zn(s) + 2 HBr(aq) → ZnCl2(aq) + H2(g)
Look at the example below for the decomposition of
When a metal element is mixed with acid, the metal will replace the hydrogen in the acid and
ammonium nitrate to dinitrogen oxide and water.
release hydrogen gas a product. Once again, note that the charges of the ionic compounds
Example: NH4NO3 → N2O + 2 H2O must equal zero. To correctly predict the formula of the ionic product, you must know the
Notice the one reactant, NH4NO3, is on the left of the 2+ -
arrow and there is more than one on the right side of charges of the ions you are combining, in this case Zn and Cl .
Cl2(g) + 2 KI(aq) → 2 KCl(aq) + I2(s)
the arrow. This is the exact opposite of the synthesis
When a nonmetal element is added to an ionic compound, the element will replace the
reaction type. A decomposition reaction is similar to
nonmetal in the compound. Also, to correctly write the formulas of the products, you must
When studying decomposition reactions, we breaking up a couple, in which the
can predict reactants in a similar manner as we did for individuals have different properties
first identify the charges of the ions that will be in the ionic compound.
synthesis reactions. Look at the formula for from the couple they started out in.
Example: What would be the products of the reaction between solid aluminum and iron(III)
magnesium nitride, Mg3N2. What elements do you see CC – Tracy Poulsen oxide? The reactants are: Al + Fe2O3 →
in this formula? You see magnesium and nitrogen. Now we can write a Solution: In order to predict the products we need to know that aluminum will replace iron
decomposition reaction for magnesium nitride. Notice there is only one reactant. and form aluminum oxide (the metal will replace the metal ion in the compound). Aluminum
Mg3N2 → 3 Mg + N2 has a charge of +3 and oxygen has a charge of -2. The compound formed between aluminum
and oxygen, therefore, will be Al2O3. Since iron is replaced in the compound by aluminum,
Example: Write the chemical equation for the decomposition of:
the iron will now be the single element in the products. The unbalanced equation will be:
(a) Al2O3
Al + Fe2O3 → Al2O3 + Fe
(b) Ag2S and the balanced equation will be:
(c) MgO
2 Al + Fe2O3 → Al2O3 + 2 Fe
Solution:
(a) 2 Al2O3 → 4 Al + 3 O2 Example:
(b) Ag2S → 2 Ag + S
(a) Write the chemical equation for the single replacement reaction between zinc solid and
(c) 2 MgO → 2 Mg + O2
lead(II) nitrate solution to produce zinc nitrate solution and solid lead. (*Note: zinc forms
ions with a +2 charge)
(b) Predict the products for the following reaction: Fe + CuSO4 (in this reaction, assume iron
Single Replacement Reactions forms ions with a +2 charge)
A third type of reaction is the single (c) Predict the products for the following reaction: Al + CuCl2
replacement reaction. In single replacement (d) Complete the following reaction. Then balance the equation: Al + HNO3 →
reactions one element reacts with one Solution:
compound to form products. The single element (a) Zn + Pb(NO3)2 → Pb + Zn(NO3)2
is said to replace an element in the compound (b) Fe + CuSO4 → Cu + FeSO4
when products form, hence the name single (c) 2 Al + 3 CuCl2 → 3 Cu + 2 AlCl3
replacement. Metal elements will always In a single replacement reaction, a single element
(d) 2 Al + 6 HNO3 → 2 Al(NO3)3 + 3 H2
replace other metals in ionic compounds or (individual) takes the place of an element within a
hydrogen in an acid. Nonmetal elements will compound (couple), leaving a different element
(individual) separate. Double Replacement
always replace another nonmetal in an ionic CC – Tracy Poulsen For double replacement reactions two ionic compound reactants will react by having
compound. the cations exchange places, forming two new ionic compounds. The key to this type of
General equation: A + BC → B + AC reaction, as far as identifying it over the other types, is that it has two compounds as
Consider the following examples. reactants. This type of reaction is more common than any of the others and there are many
Zn(s) + Cu(NO3)2(aq) → Zn(NO3)2(aq) + Cu(s) different types of double replacement reactions. Precipitation and neutralization reactions are
two of the most common double replacement reactions. Precipitation reactions are ones

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where two aqueous compound reactants combine to form products where one of the products
is an insoluble solid. A neutralization reaction is one where the two reactant compounds are
an acid and a base and the two
products are a salt and water (i.e.
acid + base salt + water).
precipitate of calcium hydroxide.
General equation: AB + CD → AD
(b) Predict the products for the following reaction: AgNO 3(aq) + NaCl(aq) →
(c) Predict the products for the following reaction: FeCl3(aq) + KOH(aq) →
Solution:
(a) CaCl2(aq) + 2 KOH(aq) → Ca(OH)2(s) + 2 KCl(aq)
In a double replacement reaction, the cations (girls) in two (b) AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
compounds (couples) trade places to form two new ionic (c) FeCl3(aq) + KOH(aq) → Fe(OH)3(s) + KCl(aq)
compounds (couples)
CC – Tracy Poulsen
Combustion
In a combustion reaction
+ CB oxygen reacts with another
For example, when solutions of silver nitrate and sodium chloride are mixed, the following substance to produce carbon
reaction occurs: dioxide and water. This is what
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) happens when fuel burns. In a In a combustion reaction, a fuel (CxHy compound) reacts
This is an example of a precipitate reaction. Notice that two aqueous reactants form one particular branch of chemistry, with oxygen (O2) to form CO2 and H2O. CC – Tracy Poulsen
solid, the precipitate, and another aqueous product. known as organic chemistry, we
An example of a neutralization reaction occurs when sodium hydroxide, a base, is study compounds known as
mixed with sulfuric acid: hydrocarbons. A hydrocarbon is compound consisting of only hydrogen and carbon.
2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l) Hydrocarbons represent the major components of all organic material including fuels.
Combustion reactions usually have the same products, CO2 and H2O, and one of its reactants
In order to write the products for a double displacement reaction, you must be able to
is always oxygen. In other words, the only part that changes from one combustion reaction to
determine the correct formulas for the new compounds. Remember, the total charge of all
the next is the actual hydrocarbon that burns. The general equation is given below. Notice the
ionic compounds is zero. To correctly write the formulas of the products, you must know the oxygen, carbon dioxide, and water parts of the reaction are listed for you to show you how
charges of the ions in the reactants. Let’s practice with an example or two. these reactants and products remain the same from combustion reaction to combustion
reaction.
Example: A common laboratory experiment involves the reaction between lead(II) nitrate General equation: CxHy (hydrocarbon) + O2 → CO2 + H2O
and sodium iodide, both colorless solutions. The reactants are given below. Predict the Look at the reaction for the combustion of octane, C8H18, below. Octane has 8 carbon
products. atoms hence the prefix “oct”.
Pb(NO3)2(aq) + NaI(aq) → Example: 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O
Solution: This reaction is referred to as complete combustion. Complete combustion reactions occur
We know that the cations exchange anions. We now have to look at the charges of each of the when there is enough oxygen to burn the entire hydrocarbon. This is why there are only
-
cations and anions to see what the products will be. In Pb(NO 3)2, the nitrate, NO3 has a carbon dioxide and water as products.
2+
charge of -1. This means the lead must be +2, Pb . In the sodium iodide, we are combining
+ -
Na and I . Example: Write the balanced reaction for the complete combustion of propane, C3H8.
Solution: The reactants of all combustion reactions include the fuel (a compound with
Now we switch ions and write the correct subscripts so the total charge of each compound is carbon and hydrogen) reacting with oxygen. The products are always carbon dioxide and
2+ - + -
zero. The Pb will combine with the I to form PbI2. The Na will combine with the NO3 to water.
form NaNO3. C3H8+5O2→3CO2+4H2O
Only after we have the correct formulas can we worry about balancing the two sides of Lesson Summary
the reaction. The final balanced reaction will be: The Five Types of Chemical Reactions
Pb(NO3)2(aq) + 2 NaI(aq) → PbI2(s) + 2 NaNO3(aq)

Example:
(a) Write a chemical equation for the double replacement reaction between calcium chloride
solution and potassium hydroxide solution to produce potassium chloride solution and a

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 Reaction Name Reaction Description
Synthesis: two or more reactants form one product.
Decomposition: one type of reactant forms two or more products.
Single replacement: one element reacts with one compound to form products.

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17) BaCl2 + Na2SO4 →
18) Ca + HCl →
19) FeS + HCl →
Double replacement: two compounds act as reactants. 20) NaI + Br2 →
Combustion: a fuel reactant reacts with oxygen gas.
158
Vocabulary www.ck12.org
Synthesis reaction: a reaction in which two or more reactants combine to make one
product
Decomposition reaction: a reaction in which one reactant breaks down to form two
or more products
Single replacement reaction: a reaction in which an element reacts with a compound
to form products
Double replacement reaction: a reaction in which two reactants form products by
having the cations exchange places with the anions
Combustion reaction: a reaction in which oxygen reacts with another substance to
produce carbon dioxide and water
Hydrocarbon: an organic substance consisting of only hydrogen and carbon

7.4: Review Questions

Classify each type of reaction as synthesis, decomposition, single replacement,
double replacement or combustion.
1) Cu + O2 → CuO
2) H2O →H2 + O2
3) Fe + H2O → H2 + Fe2O3
4) AsCl3 + H2S → As2S3 + HCl
5) Fe2O3 + H2 → Fe + H2O
6) CaCO3 → CaO + CO2
7) H2S + KOH → HOH + K2S
8) NaCl → Na + Cl2
9) Al + H2SO4 → H2 + Al2(SO4)3
10) CH4 + O2 → CO2 + H2O

11) Distinguish between synthesis and decomposition reactions.

12) When dodecane, C10H22, burns in excess oxygen, what will be the products?

13) When iron rods are placed in liquid water, a reaction occurs. Hydrogen gas evolves from
the container and iron(III) oxide forms onto the iron rod. Classify the type of reaction
and write a balanced chemical equation for the reaction.

Classify each of the following reactions and predict products for each reaction.
14) H3PO4 + NH4OH →
15) C3H8 + O2 →
16) Al + O2 →
 2 CuSO4 + 4 KI → 2 CuI + 2 K2SO4 + I2
2 formula units 4 formula 2 formula 2 formula units
CuSO4 + units KI → units CuI + K2SO4 + 1 molecule I2
7.6: Stoichiometry 2 moles CuSO4 + 4 moles KI → 2 moles CuI + 2 moles K2SO4 + 1 moles I2
Objectives The coefficients used, as we have learned, tell us the relative amounts of each substance in
Explain the meaning of the term “stoichiometry”. the equation. So for every 2 units of copper (II) sulfate (CuSO4) we have, we need to have
Determine mole ratios in chemical equations. 4 units of potassium iodide (KI). For every two dozen copper(II) sulfates, we need 4 dozen
Calculate the number of moles of any reactant or product from a balanced equation given the number of moles of one potassium iodides. Because the unit “mole” is also a counting unit, we can interpret this
reactant or product. equation in terms of moles, as well: For every two moles of copper(II) sulfate, we need 4
Calculate the mass of any reactant or product from a balanced equation given the mass of one reactant or product. moles potassium iodide.
Look at the chemical equation below. This reaction can be interpreted many ways.
Introduction N2O3 + H2O 2 HNO2
You have learned that chemical equations provide us with information about the types of particles that react to form
One molecule of dinitrogen trioxide plus one molecule of water yields two molecules
products. Chemical equations also provide us with the relative number of particles and moles that react to form products. In this
of hydrogen nitrite.
chapter you will explore the quantitative relationships that exist between the quantities of reactants and products in a balanced
equation. This is known as stoichiometry. One mole of dinitrogen trioxide plus one mole of water yields two moles of hydrogen
Stoichiometry, by definition, is the calculation of the quantities of reactants or products in a chemical reaction using nitrite.
the relationships found in the balanced chemical equation. The word stoichiometry is actually Greek coming from two Example: For each of the following equations, indicate the number of formula units or
words stoikheion, which means element and metron, which means measure. molecules, and the number of moles present in the balanced chemical equation.
(a) 2C2H6+7O2 4CO2+6H2O
Interpreting Chemical Equations (b) KBrO3 + 6 KI + 5 HBr 7 KBr + 3 I2 + 3 H2O
The mole, as you remember, is a quantitative measure that is equivalent to Avogadro’s number of particles. So how 159
does this relate to the chemical equation? Look at the chemical equation below.
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2 H2(g) + O2(g) 2 H2O(l)
2 moles of H2 react with 1 mole of O2 to produce 2
Solution: moles of H2O. Or, an alternate method to represent
(a) Two molecules of C2H6 plus seven molecules of O2 yields four molecules of CO2 plus this information is with mole ratios. The following
mole ratios can be obtained from this reaction:
six molecules of H2O.
Two moles of C2H6 plus seven moles of O2 yields four moles of CO2 plus six moles of H2O.
(b) Two formula units of KBrO3 plus six formula units of KI plus six formula units of HBr
yields seven formula units of KBr plus three molecules of I2 and three molecules of H2O.
Two moles of KBrO3 plus six moles of KI plus six moles of HBr yields seven moles of
Using the coefficients of a balanced
KBr plus three moles of I2 and three moles of H2O.
reaction, you can compare any two substances in
Stoichiometry the reaction you are interested in, whether they are
reactants or products. The correct mole ratios of
In chemistry, we “talk” to each other using chemical equations, the same way
the reactants and products in a chemical equation
mathematicians talk to each other using mathematical equations. In chemistry, we also want
to talk about quantities. Using stoichiometry, you can predict the quantities of reactants as are determined by the balanced equation.
products that can be used and produced in a chemical reaction. This requires working with Therefore, the chemical equation MUST always be
balanced chemical equations. balanced before the mole ratios are used for
In the previous section we explored mole relationships in balanced chemical calculations.
equations. In this section, we will use the mole as a conversion factor to calculate moles of
product from a given number of moles of reactant or moles of reactant from a given number Mole-Mole Calculations
of moles of product. This is called a “mole-mole” calculation. We will also perform “mass- We have already learned the process
mass” calculations, which allow you to determine the mass of reactant you require to through which chemists solve many math
produce a given amount of product or to calculate the mass of product you can obtain from a problems, the factor-label method. The mole-
given mass of reactant. mole ratio we obtain from a balanced reaction can
be used as a ratio in part of that process.
Mole Ratios
A mole ratio is the relationship of the number of moles of the substances in a
reaction. For instance, in the following reaction we read the coefficients as molecules (or 160
formula units) and moles: www.ck12.org
 For example, 15.0 g of chlorine gas is bubbled over liquid sulfur to produce disulfur
dichloride. How much sulfur, in grams, is needed according to the balanced equation:
Cl2(g) + 2 S(l) S2Cl2(l)
Example: If only 0.050 mol of magnesium hydroxide, Mg(OH)2, is present, how many
1. Identify the given: 15.0 g Cl2
moles of phosphoric acid, H3PO4, would be required for the reaction? 2. Identify the find: g S
2 H3PO4 + 3 Mg(OH)2 Mg3(PO4)2 + 6 H2O 3. Next, use the correct ratios that allow you to cancel the units you don’t want and get
Solution: We need to set up this problem using the same steps of dimensional analysis. to the unit you are calculating for.
Given: 0.050 mol Mg(OH)2
Find: mol H3PO4
The ratio we need is one that compares mol Mg(OH)2 to mol H3PO4. This is the ratio
obtained in the balanced reaction. Note that there are other reactants and products in this reaction, but we don’t need to use them
to solve this problem. If we combine the mole-mole ratio with ratios we learned previously, when we
first learned about the mole, we have several ratios we can use to solve a wide variety of
problems. The mole map is a tool we can use to help us to know which ratios to use when
solving problems.
Notice if the equation was not balanced, the amount of H3PO4 would have been different. The reaction MUST be balanced to use the You use this map much like you would use a road map. You must first find out
reaction in any calculations. As you can see, the mole ratios are useful for converting between the number of moles of one substance where you are on the map (your given units) and where you would like to go (your “find”
and another. units). The map will then let you know which roads (ratios) to take to get there. Let’s see
how this works with a couple of example problems.
Calculations Using a Mole Map
Being able to perform mass-mass calculations allows you to determine the mass of reactant (how many grams) you require to
produce a given amount of product; or to calculate the mass of product you can obtain from a given mass of reactant or the mass of reactant 161
needed to react with a specific amount of another reactant. Just as when working with mole ratios, it is important to make sure you have a www.ck12.org
balanced chemical equation before you begin.
These types of problems can be done using dimensional analysis, also called the factor-label method. This is simply a method that
uses conversion factors to convert from one unit to another. In this method, we can follow the cancellation of units to the correct answer.
Fe2O3(s) + 2 Al(s) → 2 Fe(l) + Al2O3(s)
If 5.00 g of iron is produced, how much iron(III) oxide was placed
in the original container?
The Mole Map Solution:
1) Identify the “given”: 5.00 g iron.
(Even though this is a product, it is
still the measurement given to us in
the problem.)
2) Identify the units of the “find”: g Fe2O3
(remember, mass is measured in grams)
3) Ratio
s: This
is
where
the map
comes
in
handy.
To start
with,
we are
at 5.00
g Fe.
CC Tracy Poulsen For this
proble
m, then,
Example: The thermite reaction is a very exothermic reaction which produces liquid iron,
given by the following balanced equation:
“A” on the map stands for Fe. We start at grams
A.
We want to know g Fe2O3. For this problem, “B”
Example: Ibuprofen is a common painkiller used by many people around the globe. It has the
stands for Fe2O3. We are heading to grams B.
formula C13H18O2. If 200.g of Ibuprofen is combusted how much carbon dioxide is
Our map tells us this problem will take 3 ratios (3 roads from g A to g B): molar mass of A, produced? The balanced reaction is:
mol:mol ratio from a balanced reaction, and molar mass of B. To solve our problem, the 2 C13H18O2 + 33 O2 → 26 CO2 + 18 H2O
work will look like: Solution:
Given: 200. g C13H18O2 (g A on the map)
Find: g CO2 (g B on the map)

Ratios: The map says we need to use the molar mass

of C13H18O2, then the coefficients of the balanced
162 reaction, then the molar mass of CO2.
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Example: If sulfuric acid is mixed with sodium cyanide, the deadly gas hydrogen cyanide
is produced. How many moles of sulfuric acid would have been placed in the container to
produce 12.5 g of hydrogen cyanide? The balanced reaction is:
2 NaCN + H2SO4 → Na2SO4 + 2 HCN
Solution:
Given: 12.5 g HCN (g A on map)
Find: mol H2SO4 (mol A on map)

Ratios: The mole map says we need the molar mass

of HCN and the coefficients of the balanced reaction.

Example: How many atoms of carbon would be released from the complete dehydration of
18.0 g of sugar (C6H12O6) with sulfuric acid? The balanced reaction is:
C6H12O6 + H2SO4→ 6 C + 7 H2O + SO3
Solution:
Given: 18 g C6H12O6
Find: atoms C

Ratios: the mole map says we need the molar mass

of the sugar, the balanced reaction, and finally
Avogadro’s number.

Lesson Summary

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 Stoichiometry is the calculation of the quantities of reactants or products in a
chemical reaction using the relationships found in the balanced chemical
equation. 10) Determine the mass of lithium hydroxide produced when 0.38 grams of lithium nitride
The coefficients in a balanced chemical equation represent the reacting ratios of reacts with water according to the following equation: Li3N + 3H2O → NH3 + 3LiOH
the substances in the reaction.
The coefficients of the balanced equation can be used to determine the ratio of moles 23
11) If 3.01x10 formulas of cesium hydroxide are produced according to this
of all the substances in a reaction. reaction: 2Cs + 2H2O → 2CsOH + H2, how many grams of cesium reacted?
Vocabulary 12) How many liters of oxygen are necessary for the combustion of 425 g of sulfur, assuming
Stoichiometry: the calculation of quantitative relationships of the reactants and that the reaction occurs at STP? The balanced reaction is:
products in a balanced chemical equation S + O2 → SO2 (hint: one mole of oxygen is 22.4 Liters at STP)
formula unit: the empirical formula of an ionic compound
Mole ratio: the ratio of the moles of one reactant or product to the moles of another 13) If I have 2.0 grams of carbon monoxide, how many molecules of carbon monoxide are
reactant or product according to the coefficients in the balanced chemical equation there?

Further Reading / Supplemental Links 14) What mass of oxygen is needed to burn 3.5 g of propane (C3H8) is burned according
Stoichiometry: http://www.lsua.us/chem1001/stoichiometry/stoichiometry.html to the following equation: C3H8 + 5O2 → 4H2O + 3CO2

7.6: Review Questions 15) How many grams of water are produced if 5 moles of oxygen react according to the
1) Given the reaction between ammonia and oxygen to produce nitrogen monoxide, following reaction? 2H2 + O2 → 2H2O
how many moles of water vapor can be produced from 2 mol of ammonia? The
balanced reaction is: 4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
7.7: Reversible reaction & Equilibrium
2) When properly balanced, how many moles of bismuth(III) oxide can be produced
Objectives
from 0.625 mol of bismuth? The unbalanced reaction is: Bi(s) + O2(g) → Bi2O3(s) Describe the three possibilities that exist when reactants come together.
3) Solid lithium reacts with an aqueous solution of aluminum chloride to produce aqueous Describe what is occurring in a system at equilibrium.
lithium chloride and solid aluminum. The reaction is: 3 Li + AlCl3 → 3 LiCl + Al. How Introduction
many moles of lithium chloride are formed if 5.0 mol aluminum were produced?
Think for a minute about sitting down to a table to eat dinner. There are three
possibilities that could happen when you eat dinner. You could (1) finish your entire dinner,
For the given balanced reaction: Ca3(PO4)2 + 3 SiO2 + 5 C → 3 CaSiO3 + 5 CO + 2 P
(2) you could not want any of it and leave it all on your plate, or (3) you could eat some of it
4) How many moles of silicon dioxide are required to react with 0.35 mol of carbon?
and leave some of it. Reactions have the same possibilities. Reactions also do not always
5) How many moles of calcium phosphate are required to produce 0.45 mol of calcium
proceed all the way from start to finish. You may have reactions that (1) go to completion so
silicate?
that at the end the reaction vessel contains all products and only products. Some reactions
(2) may not start at all so at the end the reaction vessel contains all reactants and only
For the given balanced reaction, 4 FeS + 7 O2 → 2 Fe2O3 + 4 SO2
reactants. And some reactions (3) may start but not go to completion, that is, the reaction
6) How many moles of iron(III) oxide are produced from 1.27 mol of oxygen?
might start but not go completely to products. In this last case, at the end, the reaction vessel
7) How many moles of iron(II) sulfide are required to produce 3.28 mol of sulfur dioxide?
would contain come reactants and some products. In this chapter, we are going to take a
8) Given the reaction between copper (II) sulfide and nitric acid, how many grams of closer look at the third type of reaction.
nitric acid will react with 2.00 g of copper(II) sulfide? Reversible Reactions and Equilibrium
3 CuS(s) + 8 HNO3(aq) → 3 Cu(NO3)2(aq) + 2 NO(g) + 4 H2O(l) + 3 S(s)
Consider the hypothetical reaction: A + B → C + D. If we looked at this reaction
9) When properly balanced, what mass of iodine was needed to produce 2.5 g of sodium using what we have learned, this reaction will keep going, forming C and D until A and
B run out. This is what we call an “irreversible reaction” or a “reaction that goes to
iodide in the equation below? I2(aq) + Na2S2O3(aq) → Na2S4O6(aq) + NaI(aq)
completion”.
Some reactions, however, are reversible, meaning the reaction can go backwards in
which products react to form reactants, so that: A + B C + D. The direction of the arrow
shows that C and D are reacting to form A and B. What if the two reactions, the forward
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reaction and the reverse reaction, were occurring at the same time? What would this look like? If
you could peer into the reaction, you would be able to find A, B, C, and D particles. A and B
would react to form C and D at the same time that C and D are reacting to form A and B. If the
Irreversible reactions (those that only go in one direction from reactants to products
forward and reverse reactions are happening at the same rate, the reaction is said to be at
and cannot reach a state of equilibrium) is more like a game of sharks and minnows. In
equilibrium or dynamic equilibrium. At this point, the concentrations of A, B, C, and D are
sharks and minnows almost everybody starts out as a minnow. Once tagged, they become a
not changing (or are constant) and we would see no difference in our reaction container, but
shark. However, the difference here is that once you are a shark you are always a shark;
reactions are still occurring in both directions. It is important to point out that having constant
there is no way to go back to becoming a minnow. The game continues until everybody has
amounts of reactants and products does NOT mean that the concentration of the reactants is
been tagged and becomes a shark. This is similar to irreversible reactions in that the
equal to the concentration of the products. It means they are not changing. These reactions
reactants turn into products, but can’t change back. Furthermore, the reaction will proceed
appear to have stopped before one of the reactants has run out.
until the reactants have been used up and there isn’t any more left. We could write the
reaction as:
Minnow → Shark

Lesson Summary
There are a few possible ways a reaction can go: It can go to completion (reactants →
products); it can occur but not go to completion. Instead it would reach chemical
equilibrium (reactants ֖ products).
Chemical equilibrium occurs when the number of particles becoming products is
equal to the number of particles becoming reactants.
A dynamic equilibrium is a state where the rate of the forward reaction is equal to the
In reversible reactions, not all of the reactants are used to make products. Instead, both reactants rate of the reverse reaction.
and products are left over at the end.
Further Reading / Supplemental Links
Chemists use a double-headed arrow, ֖, to show that a reaction is at equilibrium.
We would write the example reaction as: A + B ֖ C + D. The arrow indicates that both http://en.wikipedia.org/wiki/Chemical_equilibrium
directions of the reaction are happening.
Another way to think about reversible and irreversible reactions is to compare them Vocabulary
to two types of games of tag. Reversible reactions are in many ways like a traditional game Equilibrium: A state that occurs when the rate of forward reaction is equal to the rate
of tag: The “it” person can become “not it” and somebody who is “not it” is tagged and of the reverse reaction.
becomes “it”. In this way it is a reversible change. It is also like a reaction at equilibrium,
7.7: Review Questions
1) For the reaction PCl5(g) ֖ PCl3(g) + Cl2(g), describe what is happening to make this an
equilibrium reaction. #2
because overall no change is
Time (min)[HC2H3O2] mol/L
occurring. There is always a 2) If the following table of concentration vs. time was 0 0.100
constant number of “it” people and provided to you for the ionization of acetic acid. When does 0.5 0.099
“not it” people in the game. Also, the reaction reach equilibrium? How do you know? 1.0 0.098
having constant numbers of “it” and 1.5 0.097
“not it” people in our game does not 3) The word “equilibrium” comes from the word “equal”. What 2.0 0.096
mean that the number of “it” people does the term equal mean in this definition? 2.5 0.095
(reactants) is equal to the number of 3.0 0.095
“not it” people. Furthermore, this is Indicate whether each of the following statements is true or 3.5 0.095
similar to equilibrium in that this false for a system in equilibrium. 4.0 0.095
game never 4) The amount of products is equal to the amount of reactants.
truly ends (unless everybody gets Pretend a bridge connects two cities separated by a river. 5) The amount of product is not changing.
tired of playing). The game could This situation models equilibrium if the rate that the cars
6) The amount of reactant is not changing.
go on forever. We could write this move between the cities is the same. This does not mean that
the same number of cars are in City A as are in City B. 7) Particles (atoms/molecules) are not reacting.
as the following reversible reaction: CC Tracy Poulsen 8) The rate of the forward reaction is equal to the rate of the reverse reaction.
“It” ֕ “Not it”

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7.8: Equilibrium Constant
Objectives Example: Write the equilibrium expression for:
Write equilibrium constant expressions. P4(s) + 6 Cl2(g) ֖ 4 PCl3(s)
Use equilibrium constant expressions to solve for unknown concentrations.
Solution:
Use known concentrations to solve for the equilibrium constants.
Explain what the value of K means in terms of relative concentrations of
reactants and products.
*Note that the only product is a solid, which is left out. That leaves just 1 on top in the
Introduction numerator
In the previous section, you learned about reactions that can reach a state of
equilibrium, in which the concentration of reactants and products aren’t changing. If these Example: Write the equilibrium expression for:
+ -
amounts are changing, we should be able to make a relationship between the amount of H2O(l) ֖ H (aq) + OH (aq)
product and reactant when a reaction reaches equilibrium. Solution:

The Equilibrium Constant

Equilibrium reactions are those that do not go to completion but are in a state where Mathematics with Equilibrium Expressions
the reactants are reacting to yield products and the products are reacting to produce reactants. The equilibrium constant value is the ratio of the concentrations of the products
In a reaction at equilibrium, the equilibrium concentrations of all reactants and products can over the reactants. This means we can use the value of K to predict whether there are more
be measured. The equilibrium constant (K) is a mathematical relationship that shows how products or reactants at equilibrium for a given reaction.
the concentrations of the products vary with the concentration of the reactants. Sometimes, If the equilibrium constant is "1" or nearly "1", it indicates that the molarities of the
subscripts are added to the equilibrium constant symbol K, such as Keq, Kc, Kp, Ka, Kb, and reactants and products are about the same. If the equilibrium constant value was a large
15
Ksp. These are all equilibrium constants and are subscripted to indicate special types of number, like 100, or a very large number, like 1x10 , it indicates that the products
equilibrium reactions. (numerator) is a great deal larger than the reactants. That means that at equilibrium, the
There are some rules about writing equilibrium constant expressions that you must great majority of the material is in the form of products and we say the "products are
-12
learn: strongly favored". If the equilibrium constant is small, like 0.10, or very small, like 1x10 ,
1. Concentrations of products are multiplied on the top of the expression. it indicates that the reactants are much larger than the products and the reactants are strongly
Concentrations of reactants are multiplied together on the bottom. favored. With large K values, most of the material at equilibrium is in the form of products
2. Coefficients in the equation become exponents in the equilibrium expression. and with small K values, most of the material at equilibrium is in the form of the reactants.
3. Leave out solids and liquids, as their concentrations do not change in a reaction The equilibrium expression is an equation that we can use to solve for K or for
the concentration of a reactant or product.
Example: Write the equilibrium expression for:
CO(g) + 3 H2(g) ֖ CH4(g) + H2O(g) Example : For the reaction, SO2(g) + NO2(g) ֖ SO3(g) + NO(g)
Solution: determine the value of K when the equilibrium concentrations are: [SO2]=1.20 M,
[NO2]=0.60 M, [NO]=1.6 M, and [SO3]=2.2 M.
Solution:
*Note that the coefficients become exponents. Also, note that the concentrations of products Step 1: Write the equilibrium constant expression:
in the numerator are multiplied. The same is true of the reactants in the denominator

Example: Write the equilibrium expression for: Step 2: Substitute in given values and solve:
2 TiCl3(s) + 2 HCl(g) ֖ 2 TiCl4(s) + H2(g)
Solution:

Example: Consider the following reaction: CO(g) + H2O(g) ֖ H2(g) + CO2(g); K=1.34
*Note that the solids are left out of the expression completely If the [H2O]=0.100 M, [H2]=0.100 M, and [CO2]=0.100 M at equilibrium, what is the
equilibrium concentration of CO?
Solution:
Step 1: Write the equilibrium constant expression:
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Step 2: Substitute in given values and solve:

15) For the reaction: MgCl2(s) + ½ O2(g) ֖ 2 MgO(s) + Cl2(g). The equilibrium constant was
found to be 3.86 at a certain temperature. If [O2]=0.560 M at equilibrium, what is the
Solving for [CO], we get: [CO]=0.0746 M concentration of Cl2(g)?

Lesson Summary 16) Consider the equilibrium: CO(g) + H2O(g) ֖ H2(g) + CO2(g).
The equilibrium expression is a mathematical relationship that shows how the a) Write an equilibrium expression for this reaction.
concentrations of the products vary with the concentration of the reactants. b) If [CO]=0.200M, [H2O]=0.500M, [H2]=0.32M and [CO2]=0.42M, find K.
If the value of K is greater than 1, the products in the reaction are favored; if the
value of K is less than 1, the reactants in the reaction are favored; if K is equal to 1, 17) Hydrogen sulfide decomposes according to the equation: 2H2S(g) ֖ 2H2(g) + S2(g).
neither reactants nor products are favored. a) Write an equilibrium expression for this reaction.
-
b) At equilibrium, the concentrations of each gas are as follows: [H2S]=7.06x10
Further Reading / Supplemental Links 3 -3 -3
M, [H2]=2.22x10 M and [S2]=1.11x10 M. What is Keq?
http://en.wikipedia.org/wiki/Chemical_equilibrium
Crabapples & Equlibrium: 18) Given the following system in equilibrium: 2SO2(g) + O2(g) ֖ 2SO3(g)
http://www.chem.ox.ac.uk/vrchemistry/ChemicalEquilibrium/HTML/page05.htm a) Write an equilibrium expression for the reaction.
Equilibrium Animation / Applet: Dots: http://chemconnections.org/Java/equilibrium/ b) If K=85.0, would you expect to find more reactants or products at equilibrium?
Why?
Vocabulary c) If [SO2]=0.0500 M and [O2]=0.0500M, what is the concentration of SO3
Equilibrium constant (K): A mathematical ratio that shows the concentrations of the at equilibrium?
products divided by concentration of the reactants.

7.8: Review Questions 7.9: The Effects of Applying Stress to Reactions at Equilibrium
1) Which phases of substances are not included in the equilibrium expression? Objectives
State Le Châtelier’s Principle.
Write an equilibrium expression for each reaction: Describe the effect of concentration on an equilibrium system.
2) 2 H2(g) + O2(g) ֖ 2 H2O(g) Describe the effect of temperature as a stress on an equilibrium system.
3) 2 NO(g) + Br2(g) ֖ 2 NOBr(g)
4) NO(g) + O3(g) ֖ O2(g) + NO2(g) Introduction
5) CH4(g) + H2O(g) ֖ CO(g) + 3 H2(g) When a reaction has reached equilibrium with a given set of conditions, if the
6) CO(g) + 2 H2(g) ֖ CH3OH(g) conditions are not changed, the reaction will remain at equilibrium forever. The forward
7) 2 C2H6(g) + 7 O2(g) ֖ 4 CO2(g) +6 H2O(g) and reverse reactions continue at the same equal and opposite rates and the macroscopic
8) C2H6(g) ֖ C2H4(g) + H2(g) properties remain constant.
9) Hg(g) + I2(g) ֖ HgI2(g) It is possible, however, to disturb that equilibrium by changing conditions. For
10) SnO2(s) + 2 CO(g) ֖ Sn(s) + 2 CO2(g) example, you could increase the concentration of one of the products, or decrease the
2+ - concentration of one of the reactants, or change the temperature. When a change of this type
11) Cu(OH)2(s) ֖ Cu (aq) + 2 OH (aq)
is made in a reaction at equilibrium, the reaction is no longer in equilibrium. When you alter
12) What does a large value for K imply? something in a reaction at equilibrium, chemists say that you put stress on the equilibrium.
13) What does a small value of K imply? When this occurs, the reaction will no longer be in equilibrium and the reaction itself will
begin changing the concentrations of reactants and products until the reaction comes to a
14) Consider the following equilibrium system: 2 NO(g) + Cl2(g) ֖ 2 NOCl(g). At a certain new position of equilibrium. How a reaction will change when a stress is applied can be
temperature, the equilibrium concentrations are as follows: [NO]=0.184 M, explained and predicted. That's the topic of this section.
[Cl2]=0.165 M, [NOCl]=0.060 M. What is the equilibrium constant for this reaction?
Le Châtelier’s Principle
170 In the late 1800’s, a chemist by the name of Henry-Louis Le Châtelier was studying
www.ck12.org stresses that were applied to chemical equilibria. He formulated a principle from this research

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and, of course, the principle is called Le Chatelier's Principle. Le Châtelier’s Principle states
that when a stress is applied to a system at equilibrium, the equilibrium will shift in a direction to
partially counteract the stress and once again reach
equilibrium.
Le Chatelier's principle is not an explanation of
anything else?
what happens on the molecular level to cause the
equilibrium shift, it is simply a quick way to determine Solution:
which way the reaction will run in response to a stress (a) The equilibrium will shift toward the products (forward).
applied to the system at equilibrium. *Thinking of a teeter-totter is a good (b) The equilibrium will shift toward the reactants (backward).
way to remember Le Chatlier’s
Principle. If a change is made, what The Effect of Changing Temperature on a System at Equilibrium
does the reaction need to do to get back
to equilibrium? Le Chatelier's principle also correctly predicts the equilibrium shift when systems at
CC Tracy Poulsen equilibrium are heated and cooled. An increase in temperature is the same as adding heat to
the system. Consider the following equilibrium:
2 SO2(g) + O2(g) ֖ 2 SO3(g) ΔH= - 191 kJ
We will learn more about this later, but ΔH has to do with the change in energy, usually heat,
Effect of Concentration Changes on a System at Equilibrium for this reaction. The negative sign (-) in the ΔH indicates that energy is being given off.
For instance, if a stress is applied by increasing the This equation can also be written as:
concentration of a reactant, the reaction will adjust in such 2 SO2(g) + O2(g) ֖ 2 SO3(g) + 191 kJ of heat
a way that the reactants and products can get back to What’s important to remember about increasing the temperature of an equilibrium
equilibrium. In this case, you made it so there is too much system, is the energy can be thought as just another product or reactant. In this example, you
reactant. The reaction will use up some of the reactant to can clearly see that the 191 kJ are a product. Therefore when the temperature of this system
make more product. We would say the reaction “shifts to is raised, heat is being added and the effect will be the same as increasing any other product.
the products” or “shifts to the right”. If you increase the Increasing a product causes the reaction will use up some of the products to make more
concentration of a product, you have the opposite effect. reactants. And, if the temperature for this equilibrium system is lowered, the equilibrium
The reaction will use up some of the product to make will shift to make up for this stress. When the temperature is decreased for this reaction, the
more reactant. The reaction “shifts to the reactants” or reaction will shift toward the products in an attempt to counteract the decreased temperature.
“shifts to the left”. Therefore, the [SO3] will increase and the [SO2] and [O2] will decrease.
What is we remove some reactant or product? If a In some reactions, though, heat is a reactant. These reactions are called endothermic
stress is applied by lowering a reactant concentration, the reactions. These reactions would have the opposite effect. If heat is a reactant, adding head
reaction will try to replace some of the missing reactant. It adds a reactant and the reaction will shift towards the products. If heat is removed (by
uses up some of the product to make more reactant, and lowering the temperature) from an endothermic reaction, a reactant is removed and the
the reaction “shifts to the reactants”. If a stress is applied reaction will shift to make more reactants.
by reducing the concentration of a product, the equilibrium
position will shift toward the products. The Effect of Temperature on an Endothermic and an Exothermic
Equilibrium System
Example: For the reaction: SiCl4(g) + O2(g) ֖ SiO2(s) + 2 Cl2(g), what would be the effect
Temperature Change Exothermic (-ΔH) Endothermic (+ΔH)
on the equilibrium system if:
Shifts left, favors Shifts right, favors
(a) [SiCl4] increases Increase Temperature
(b) [O2] increases reactants products
(c) [Cl2] increases Shifts right, favors Shifts left, favors
Decrease Temperature
Solution: products reactants
(a) [SiCl4] increases: The equilibrium would shift to the right
(b) [O2] increases: The equilibrium would shift to the right Example: Predict the effect on the equilibrium position if the temperature is increased in
(c) [Cl2] increases: The equilibrium would shift left each of the following.
(a) H2(g) + CO2(g) ֖ CO(g) + H2O(g) ΔH= + 40kJ/mol
(b) 2 SO2(g) + O2(g) ֖ 2 SO3(g) + energy
Example: Here's a reaction at equilibrium. A(aq) + B(aq) ֖ C(aq) + D(aq) Solution:
(a) Which way will the equilibrium shift if you add some A to the system without changing (a) The reaction is endothermic, because ΔH is positive, meaning heat is a reactant. We
anything else? would write: 40 kJ + H2(g) + CO2(g) ֖ CO(g) + H2O(g)
(b) Which way will the equilibrium shift if you add some C to the system without changing With an increase in temperature for an endothermic reaction, the reactions will shift right

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producing more products.

(b) The reaction is exothermic, meaning heat is a product. With an increase in temperature
for an exothermic reaction, the reactions will shift left producing more reactants.

The Haber Process Exothermic reaction: A reaction in which heat is released, or is a product of a
Let’s look at a particularly useful reaction and how chemists applied Le Chatlier’s reaction.
Principle to make more of a desires product. The reaction between nitrogen gas and hydrogen Endothermic reaction: A reaction in which heat is absorbed, or is a reactant of a
gas can produce ammonia, NH3. Under normal conditions, this reaction does not produce very reaction.
th
much ammonia. Early in the 20 century, the commercial use of this reaction was too Catalyst: A substance that increases the rate of a chemical reaction but is, itself, left
expensive because of the small yield of ammonia. The reaction is as follows: unchanged, at the end of the reaction.
N2(g) + 3 H2(g) ֖ 2 NH3(g) + energy
th
Fritz Haber, a German chemist working in the early years of the 20 century, applied 7.9: Review Questions
Le Châtelier’s principle to help solve this problem. Decreasing the concentration of 1) What is the effect on the equilibrium if the concentration of a reactant is increased?
ammonia, for instance, by immediately removing it from the reaction container will cause the 2) What is the effect on the equilibrium if the concentration of a reactant is decreased?
equilibrium to shift to the right and continue to produce more product. There were a number
of other ways that For the reaction: N2O5(s) ֖ NO2(g) + O2(g), what would be the effect on the equilibrium if:
One more factor that will affect this equilibrium system is the temperature. Since 3) [NO2] decreases
the forward reaction is exothermic (heat is released as a product), lowering the temperature 4) [NO2] increases
will once again shift the equilibrium system to the right and increase the ammonia that is 5) [O2] increases
produced. Specifically the conditions that were found to produce the greatest yield of
ammonia are 550°C (in commercial situations this is a “low” temperature) and 250 atm of For the reaction: C(s) + H2O(g) ֖ CO(g) + H2(g), what would be the effect on the
pressure. Once the equilibrium system is producing the ammonia, the product is removed, equilibrium system if:
cooled and dissolved in water. 6) [H2O] increases
7) [CO] increases
Lesson Summary 8) [H2] decreases
Increasing the concentration of a reactant causes the equilibrium to shift to the right
producing more products. Predict the effect on the equilibrium position if the temperature is increased in each of
Increasing the concentration of a product causes the equilibrium to shift to the left the following.
producing more reactants. 9) H2(g) + I2(g) ֖ 2 HI(g) ΔH= + 51.9 kJ
10) P4O10(s) + H2O(l) ֖ H3PO4(aq) + heat
Decreasing the concentration of a reactant causes the equilibrium to shift to the left + -
producing less products. 11) Ag (aq) + Cl (aq) ֖ AgCl(s) ΔH= - 112 kJ/mol.
12) 2 NOBr(g) ֖ 2 NO(g) + Br2(g) ΔH =
Decreasing the concentration of a product causes the equilibrium to shift to the right
+16.1kJ.
producing more products.
For a forward exothermic reaction, an increase in temperature shifts the equilibrium In the following reaction, what would be the effect of each of the following changes to
toward the reactant side whereas a decrease in temperature shifts the equilibrium the system at equilibrium? C(s) + O2(g) ֖ CO2(g) ΔH= -393.5 kJ/mol
toward the product side. 13) increase O2
14) increase the temperature
Further Reading / Supplemental Links
http://en.wikipedia.org/wiki Predict the effect on the equilibrium: H2O(g) + CO(g) ֖ H2(g) + CO2(g) ΔH= -42 kJ
Tutorial: Le Chatlier’s Principle: when each of the following changes are made to the equilibrium system.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf 15) Temperature is increased
16) [CO2] decreases
Vocabulary 17) [H2O] increases
Le Châtelier’s Principle: Applying a stress to an equilibrium system causes the 18) [H2] decreases
equilibrium position to shift to offset that stress and regain equilibrium.
Predict the effect on the chemical equilibrium 2 SO3(g) + heat ֖ 2 SO2(g) + O2(g), when
each of the following changes are made to the equilibrium system. What will the effect be
174 on the amount of product produced?
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20) [O2] decreases

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 Chapter 8: Describing Acids & Bases
Predict the effect on the chemical equilibrium: N2O4(g) + heat ֖ 2 NO2(g), when each of
the following changes are made to the equilibrium system. What will the effect be on the 8.1: Classifying Acids and Bases
amount of product produced?
Objectives
21) Temperature is decreased
List the properties of acids.
22) [N2O4] decreases
List the properties of bases.
Define an Arrhenius acid and list some substances that qualify as acids under this
definition.
Define an Arrhenius base and list some substances that qualify as bases under
this definition.

Introduction
We may not realize how much acids and bases affect our lives. Have you ever
thought of drinking a can of soda pop and actually drinking acid? Have you looked at bottles
of household cleaners and noticed what the main ingredients were? Have you ever heard a
shampoo commercial and heard them say that the shampoo was “pH balanced” and
wondered what this means and why it is so important for hair? Thanks to the beginning
th
work of scientists in the latter part of the 19 century, we started to learn about acids and
bases; our study continued and is constantly growing. Let’s begin our study of this
wonderful branch of chemistry.

Bases cause red litmus paper to turn Properties of Acids

blue
Acids are a special group of compounds with a set of common properties. This helps
to distinguish them from other compounds. Thus, if you had a number of compounds and
you were wondering whether these were acids or otherwise, you could identify them by their
properties. But what exactly are the properties? Think
about the last time you tasted lemons. Did they taste
sour, sweet, or bitter? Lemons taste sour. This is a
property of acids. Another property of acids is that
they turn blue litmus paper red. Litmus paper is an
indicator, which is a substance that changes color
depending on how acidic or basic something is. If blue
litmus paper turns red when it is dipped into a
solution, then the solution is an acid. Another property
of acids that many people are familiar with is their Acids cause blue litmus paper to turn
ability to cause burns to skin. This is why it is a bad red
idea to play with battery acid or other acids.
Acids react with many metals to produce hydrogen gas. For some examples, look
at the reactions below:
Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
What do you notice that is the same for all three equations? In each case, the reactants are a
metal (Zn or Mg) and an acid (HCl). They all produce hydrogen gas, H2. This is another
property of acids. Acids react with most metals to produce hydrogen gas.
Think about the last time you took an aspirin or a vitamin C tablet. Aspirin is
+
acetylsalicylic acid while vitamin C is ascorbic acid; both are acids that can produce H ions

176 177
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when dissolved in water. Acetic acid (HC2H3O2) is a component of vinegar, hydrochloric
acid (HCl) is stomach acid, phosphoric acid (H3PO4) is commonly found in dark soda pop,
sulfuric acid H2SO4 is used in car batteries and formic acid HCO2H is what causes the sting
in ant bites. For all of these acids, the chemical formula of an acid begins with one or more Arrhenius Acids
+
hydrogen atoms. Acids dissolve in water to make H ions. Because they make ions (charged Take a look at all of the following chemical equations. What do you notice about
particles) when they are dissolved, acids will also conduct electricity when they are them? What is common for each of the equations below?
dissolved in water. Hydrochloric acid: HCl(aq) → H+(aq) + Cl-(aq)
We interact with acids on a daily basis so some knowledge of their properties and
interactions is essential. Acids are present in our everyday lives. + -
Perchloric acid: HClO4(aq) → H (aq) + ClO4 (aq)
+
One of the distinguishable features about acids is the fact that acids produce H ions in
Properties of Bases
solution. If you notice in all of the above chemical equations, all of the compounds
There is one common base that some may have had the opportunity to taste: milk of +
dissociated to produce H ions. This is the one main, distinguishable characteristic of
magnesia, which is a slightly soluble solution of magnesium hydroxide. This substance is acids and the basis for the Arrhenius definition of acids. An Arrhenius acid is a substance
used for acid indigestion. Flavorings have been added +
that produces H ions in solution.
to improve the taste, otherwise it would have a bitter
taste when you drink it. Other common bases include Arrhenius Bases
substances like Windex, Drano, oven cleaner, soaps -
In contrast, an Arrhenius base is a substance that releases OH ions in solution.
and many cleaning other products. Please note: do not -
Many bases are ionic substances made up of a cation and the anion hydroxide, OH . The
taste any of these substances. A bitter taste is one dissolving equation for the base sodium hydroxide, NaOH, is shown below:
property you will have to take for granted. Bases also + -
NaOH(s) → Na (aq) + OH (aq)
tend to have a slippery feel. This matches what you Barium hydroxide produces a similar reaction when dissociating in water:
have experienced with soaps and detergents. 2+ -
Ba(OH)2(s) → Ba (aq) + 2 OH (aq)
As with acids, bases have properties that allow -
The production of OH ions is the definition of bases according to the Arrhenius.
us to distinguish them from other substances. We have
learned that acids turn blue litmus paper red. Bases turn red litmus paper blue. Notice that the Lesson Summary
effect of the indicator is the opposite of that of acids. Acids turn blue litmus paper red, taste sour, and react with metals to produce
Most acids have formulas that start with H. On the other hand, most of the bases we hydrogen gases.
will be using in this course have formulas that end with –OH. These bases contain the
polyatomic ion called hydroxide. When bases dissolve in water, they produce hydroxide Common acids include vinegar (HC2H3O2), phosphoric acid in soda pop (H3PO4)
- and stomach acid HCl.
(OH ) ions. Because they dissolve into charged particles, bases will also conduct electricity Bases turn red litmus paper blue, have a bitter taste, and are slippery to the
when they are dissolved. touch. Common bases include Drano (NaOH), soaps and detergents, milk of
Although many people have already heard of the danger of acids at causing burns, magnesia (Mg(OH)2) and Windex (NH4OH).
many bases are equally dangerous and can also cause burns. It is important to be very careful +
Arrhenius defined an acid as a substance that donates H ions when dissociating
and to follow correct safety procedures when dealing with both acids and bases. in solution.
-
An Arrhenius base is a substance that releases OH ions in solution.
Acids & Bases Defined
Although scientists have been able to classify acids and bases based on their Vocabulary
properties for some time, it took a while to come up with a theory explaining why some +
Arrhenius acid: a substance that produces H ions in solution
substances were acidic and others were basic. Svante Arrhenius set the groundwork for our
current understanding of acid-base theory. We will focus on his famous acid-base
th
definitions. This was quite an accomplishment for a scientist in the late 19 century with Further Reading / Supplemental Links
very little technology, but with the combination of knowledge and intellect available at the Strong & Weak Acids animation:
time Arrhenius led the way to our understanding of how acids and bases differed, their http://www.mhhe.com/physsci/chemistry/chang7/esp/folder_structure/ac/m2/s1/acm2
properties, and their reactions. Keep in mind that Arrhenius came up with these theories in s1_1.htm
the late 1800’s so his definitions came with some limitations. For now we will focus on his
definitions. Tutorial: Acids & Bases
http://visionlearning.com/library/module_viewer.php?mid=58&l=&c3=

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8.1: Review Questions

Indicate whether each of the following is a property of acids, bases, or both acids and bases.
1) Have a sour taste
2) Taste bitter + - - -7
3) Turns litmus paper red If [H ] < [OH ], the solution is basic. This means that [OH ] > 1x10 M.
4) Feels slippery + -
5) React with metals We can use this equation to calculate the concentrations of H and OH . Consider the
6) Turns litmus paper blue following example.
7) What is the Arrhenius definition of an acid? + -4
Example: Suppose acid is added to some water, and [H ] is measured to be 1x10 M. What
-
would [OH ] be?
8.2: pH Solution: substitute what we know into the equilibrium expression:
-14 + -
Objectives Kw=1x10 =[H ] [OH ]
-14 -4 -
+ - 1x10 =[1x10 ][OH ]
State the [H ], [OH ], and Kw values for the self-ionization of water. Define and - -4
+ To isolate [OH ], divide by sides by 1x10 .
describe the pH scale and describe how logarithmic scales work Calculate [H ], - -10
This leaves, [OH ]=1x10 M
-
[OH ], and pH given the value of any one of the other values in a water solution + -
Note that because [H ] > [OH ], the solution must be acidic.
at 25°C.
Explain the relationship between the acidity or basicity of a solution and the Suppose, on the other hand, something is added to the solution that reduces the
+ -
hydrogen ion concentration, [H ], and the hydroxide ion concentration, [OH ], of the
solution. Predict whether an aqueous solution is acidic, basic, or neutral from the hydrogen ion concentration, a base.
+ -
[H ], [OH ], or the pH. -12
Example: If the final hydrogen ion concentration is 1x10 M, we can calculate the
Introduction final hydroxide ion concentration.
We have been discussing what makes an acid or a base and what properties acids and Solution:
-14 + -
bases have. It is frequently useful to compare how acidic or basic a solution is in comparison Kw=1x10 =[H ] [OH ]
-14 -12 -
+ - 1x10 =[1x10 ] [OH ]
to other solutions. A couple of ways to do this is to compare [H ] to [OH ] or to find the pH - -12
To isolate [OH ], divide by sides by 1x10 .
of a solution. - -2
This leaves, [OH ]=1x10 M
+ - + -
Relationship Between [H ] and [OH ] Note that because [H ] < [OH ], the solution must be basic.
+
We have learned that acids and bases are related to hydrogen ions [H ] and + -
- Using the Kw expression, anytime we know either the [H ] or the [OH ]in a water
hydroxide ions [OH ]. Both of these ions are present in both acids and bases. However, they
are also present in pure water. Water self-ionizes according to the following reaction: solution, we can always calculate the other one.
+ -
H2O(l) ֖ H (aq) + OH (aq) + --11 Example: What would be the [H ] for a grapefruit found to have a [OH ] of 1.26x10 ?
The equilibrium expression for this reaction would be:
+ -
Kw=[H ][OH ]
Solution:
The equilibrium constant for this particular equilibrium is K w, meaning the equilibrium -14 + -
constant for water. From experimentation, chemists have determined that in pure water, Kw=1x10 =[H ] [OH ]
+ -7 - -7
[H ]=1x10 M and [OH ]=1x10 M. If you substitute these values into the equilibrium + -11
-14 To isolate [H ], divide by sides by 1.26x10 .
expression, you find that Kw=1x10 . Any solution which contains water, even if other + -4
things are added, will shift to establish this equilibrium. Therefore, for any solution, the This leaves, [H ]=7.94x10 M
+ -
following relationship will always be true: Also, the solution must be acidic because [H ] > [OH ]
-14 + -
Kw=1x10 =[H ] [OH ] @ 25°C
We can describe whether a solution is acidic, basic, or neutral according to the concentrations pH Scale
in this equilibrium. A few very concentrated acid and base solutions are used in industrial chemistry and
+ - inorganic laboratory situations. For the most part, however, acid and base solutions that occur
If [H ] = [OH ], the solution is neutral (such as in pure water)
+ - + -7 in nature, those used in cleaning, and those used in organic or biochemistry applications are
If [H ] > [OH ], the solution is acidic. This means that [H ] > 1x10 M.
relatively dilute. Most of the acids and bases dealt with in laboratory situations have
-14
180 hydrogen ion concentrations between 1.0 M and 1.0x10 M. Expressing hydrogen ion
concentrations in exponential numbers becomes tedious and is difficult for those not trained
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in chemistry. A Danish chemist named Søren Sørensen developed a shorter method for
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expressing acid strength or hydrogen ion concentration with a non-exponential number. He
named his method pH. The p from pH comes from the German word potenz meaning
“power or the exponent of”. Sørensen’s idea that the pH would be a simpler number to deal
+ Have you ever
with in terms of discussing acidity level led him to a formula that relates pH and [H ]:
+ cut an onion and had
pH = - log [H ]
-14 your eyes water up? This
If the hydrogen ion concentration is between 1.0 M and 1.0x10 , the value of the pH will is because of a
be between 0 and 14.
compound with the
+ formula C3H6OS that is
Example: Calculate the pH of a solution given that [H ]=0.01 M. found in onions. When
Solution: you cut the onion, a
pH = - log (0.01) variety of reactions
pH = 2 occur that release a gas.
This gas can diffuse into pH Scale for Common Substances.
Sometimes you will need to use a calculator.
the air and mix with the water found in your eyes to produce a dilute solution of sulfuric
+ -6 acid. This is what irritates your eyes and causes them to water. There are many common
Example: Calculate the pH of saliva with [H ]=1.58x10 M. examples of acids and bases in our everyday lives. Look at the pH scale to see how these
Solution: common examples relate in terms of their pH.
-6
pH = - log (1.58x10 )
pH = 5.8 Example: Compare lemon juice (pH=2.5) to milk (pH=6.5). Answer each of the following:
a) Label each as acidic, basic, or neutral
If you are given [OH-] it is still possible to find the pH, but it requires one more step. You +
b) Which has a higher concentration of H ions?
must first find [H+] and then use the pH equation. +
c) How many times more H does that solution have?
- -4 Solution:
Example: Calculate the pH of a solution with [OH ]=7.2x10 M. a) Both lemon juice and milk are acidic, because their pH’s are less than 7. (*Note: milk is
Solution: In order to find pH, we need [H+]. -14 + - only very slightly acidic as its pH is very close to 7)
Kw=1x10 =[H ] [OH ] +
-14 + -4 b) The lower the pH, the higher the concentration of H ions. Therefore, lemon juice has
1x10 =[H ] [7.2x10 ] more H .
+
+ -4
To isolate [H ], divide by sides by 7.2x10 . +
+ -11 c) Each step down on the pH scale increases the H concentration by 10 times. It is 4 steps
This leaves, [H ]=1.39x10 M down on the pH scale to go from 6.5 to 2.5. Therefore, lemon juice has 10x10x10x10 or
We can now find the pH +
10,000 times more H ions than milk.
-11
pH = - log (1.39x10 )
pH = 10.9 Lesson Summary
+ -
Water ionizes slightly according to the equation H2O(l) ֖ H (aq) + OH (aq) The
The pH scale developed by Sørensen is a logarithmic scale, which means that a -14 + -
equilibrium constant for the dissociation of water is: Kw=1x10 =[H ] [OH ]
difference of 1in pH units indicates a difference of a factor of 10 in the hydrogen ion +
pH= - log [H ].
concentrations. A difference of 2 in pH units indicates a difference of a factor of 100 in the
hydrogen ion concentrations. Not only is the pH scale a logarithmic scale but by defining the 8.2: Review Questions
pH as the negative log of the hydrogen ion concentration, the numbers on the scale get + -13
1) In saturated limewater, [H ]=3.98x10 M.
smaller as the hydrogen ion concentration gets larger. For example, pH=1is a stronger acid -
than pH=2 and, it is stronger by a factor of 10 (the difference between the pH’s is 1). a) Find [OH]
b) What is the pH?
The closer the pH is to 0 the
+ c) Is the solution acidic, basic, or neutral?
greater the concentration of [H ] ions + -7
2) In butter, [H ]=6.0x10 M.
and therefore the more acidic the -
a) Find [OH]
solution. The closer the pH is to 14, the
-
b) What is the pH?
higher the concentration of OH ions and The pH Scale. c) Is the solution acidic, basic, or neutral?
- -11
3) In peaches, [OH ]=3.16x10 M
+
a) Find [H ]
182 b) What is the pH?
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c) Is the solution acidic, basic, or neutral?

4) During the course of the day, human saliva varies between being acidic and basic. If
- -8
[OH ]=3.16x10 M,
+
a) Find [H ] Example: Complete the following neutralization reactions.
b) What is the pH?
c) Is the solution acidic, basic, or neutral? (a) H2SO4 + Ba(OH)2 →
-8
5) A solution contains 4.33x10 M hydroxide ions. What is the pH of the solution? (b) HCOOH + Ca(OH)2 →
(c) HCl + NaOH →
-9
6) A solution contains a hydrogen ion concentration of 6.43x10 M. What is the pH of Solution:
+ -
the solution? (a) The H in H2SO4 will combine with the OH part of Ba(OH)2 to make water (H2O or
2+
HOH). The salt produced is what is formed when Ba (the cation from the base) combines
+ 2-
7) If the pH of one solution is 5 less than another solution, how does the amount of H with SO4 (the anion from the acid). These have charges of +2 and -2, so the formula for this
+
in each solution compare? Which has more H ? How many times more? compound is BaSO4.
Before it is balanced, the reaction is:
H2SO4 + Ba(OH)2 → BaSO4 + H2O
After balancing, we get:
8.3: Neutralization H2SO4 + Ba(OH)2 → BaSO4 + 2 H2O
Objectives + -
(b) The H in HCOOH will combine with the OH part of Ca(OH)2 to make water (H2O or
Explain what is meant by a neutralization reaction 2+
HOH). The salt produced is what is formed when Ca (the cation from the base) combines
Write the balanced equation for the reaction that occurs when an acid reacts with a -
with COOH the anion from the acid). These have charges of +2 and -1, so the formula for
base. this compound is Ca(COOH)2.
Before it is balanced, the reaction is:
Introduction
HCOOH + Ca(OH)2 → Ca(COOH)2 + H2O
Neutralization is a reaction between an acid and a base that produces water and a salt. After balancing, we get:
The general reaction for the neutralization reaction is shown below. 2 HCOOH + Ca(OH)2 → Ca(COOH)2 + 2 H2O
acid + base → salt + water + -
(c) The H in HCl will combine with the OH part of NaOH to make water (H 2O or HOH).
In this section, we will be writing the products of neutralization reactions. +
The salt produced is what is formed when Na (the cation from the base) combines with Cl
-
the anion from the acid). These have charges of +1 and -1, so the formula for this compound
Neutralization Reactions is NaCl.
+
Acids are a combination of hydrogen ions (H ) and an anion. Examples include HCl, Before it is balanced, the reaction is:
HNO3, and HC2H3O2. Bases can be a combination of metal cations and hydroxide ions, HCl + NaOH → NaCl + H2O
-
OH . Examples include NaOH, KOH, and Mg(OH)2. According to the Arrhenius definitions The reaction is already balanced, so we are done.
+ -
of acids and bases, the acid will contribute the H ion that will react to neutralize the OH
ion, contributed by the base, to produce neutral water molecules. Lesson Summary
All acid-base reactions produce salts. The anion from the acid will combine with A neutralization reaction between an acid and a base will produce a salt and water.
the cation from the base to form the ionic salt. Look at the following equations. What do
they have in common? Vocabulary
HClO4 + NaOH → NaClO4 + HOH Neutralization: a reaction between an acid and a base that produces water and a salt
H2SO4 + 2 KOH → K2SO4 + 2 HOH
(Note: HOH is the same as H2O) 8.3: Review Questions
No matter what the acid or the base may be, the products of this type of reaction will Write a balanced reaction for each of the following neutralization reactions:
+ -
always be a salt and water. The H ion from the acid will neutralize the OH ion from the 1) HNO3 + KOH →
base to form water. The other product is a salt formed when the cation of the base combines 2) HClO4 + NH4OH →
with the anion of the acid. Remember, the total charge on the salt MUST be zero. You must 3) H2SO4 + NaOH →
have the correct number of cations and anions to cancel out the charges of each. 4) HNO3 + NH4OH →
5) HF + NH4OH →
6) HC2H3O2 + KOH →
7) HCl + KOH →
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185
8) M k
il
 of magnesia, Mg(OH)2 is a common over- the - counter antacid that has, as its main
ingredient, magnesium hydroxide. It is used by the public to relieve acid indigestion.
Acid indigestion is caused by excess stomach acid, HCl, being present. a hydrogen ion attached as HIn and we represent the
9) Hydrochloric acid (HCl) reacts with barium hydroxide. indicator ion without the hydrogen attached as In .
-
10) Sodium hydroxide reacts with perchloric acid (HClO4). -
For the example above, HIn red and In
yellow. If we add hydrogen ion to the solution, the
8.4: Titration equilibrium will be driven toward the reactants and The color of an indicator depends on
Objectives the solution will turn red. If we add base to the the pH of the solution.
Explain what an acid/base indicator is. solution (reduce hydrogen ion concentration), the
Explain how a titration is performed equilibrium will shift
Calculate the concentration of unknown acid or base when given the concentration of toward the products and
the other and the volume needed to reach the equilibrium point in a titration. the solution will turn
yellow. It is important to
note that if this indicator
changes color at pH=5,
Introduction then at all pH values less than 5, the solution will be red and at all pH values greater than 5,
For acid-base neutralization reactions, the typical laboratory procedure for the solution will be yellow. Therefore, putting this indicator into a solution and having the
determining the stoichiometric amounts of acid and/or base in the reaction is to complete a solution turn yellow does NOT tell you the pH of the solution . . . it only tells you that the
titration. As we go through this section, we will use some of the prior knowledge we have pH is greater than 5 . . . it could be 6, 7, 8, 9, etc. There are many indicators that are available
obtained about acids and bases, chemical reactions, and molarity calculations, to apply them to be used to help determine the pH of solutions.
to the concept of titrations.
The Titration Process
Indicators One of the properties of acids and bases is that they neutralize each other to form
An indicator is a substance that changes color at a specific pH and is used to water and a salt. In the laboratory setting, an experimental procedure where an acid is
indicate the pH of the solution. Litmus paper is a paper that has been dipped in an indicator. neutralized by a base (or vice versa) is known as titration. Titration, by definition, is the
The litmus paper is called an indicator because it is used to indicate whether the solution is addition of a known concentration of base (or acid) to a solution of acid (or base) of
an acid or a base. If the red litmus paper turns blue, the solution is basic (pH > 7), if the blue unknown concentration. Since both volumes of the acid and base are known, the
litmus turns red the solution is acidic (pH < 7). concentration
The juice from red cabbage can be used of the unknown solution is then
to prepare an indicator paper. It contains the mathematically determined.
chemical anthrocyanin, which is the active So what does one do in a titration?
ingredient in the indicator. Red beets, When doing a titration, you need to have a few
blueberries, and cranberries are other great pieces of equipment. A buret is used to
examples of a naturally occurring indicators. accurately dispense the volume of the solution
Another example of a natural indicator is of known concentration (either the base or the
flowers. Hydrangea is a common garden plant acid). A flask is used to hold a known,
with flowers that come in many colors measured volume of the unknown
depending on the pH of the soil. If you are concentration of the other solution (either the
travelling around and see a hydrangea plant acid or the base).
with blue flowers, the soil is acidic, the creamy If the basic solution was in the buret,
white flowers indicate the soil is neutral, and the you would first read the volume of base in the The set up of a titration
pink flowers mean the soil is basic. buret at the beginning. You would add the
There are two requirements for a substance to function as an acid-base indicator; 1) the base to the flask containing the acid until all of the acid has reacted and then read the
substance must have an equilibrium affected by hydrogen ion concentration, and 2) the two volume of base in the buret again. To see how much was added, you would subtract the
forms of the compound on opposite sides of the equilibrium must have different colors. Most initial volume from the final volume.
indicators function in the same general manner and can be presented by a generic indicator In a titration, just enough base is added to completely react with all of the acid,
equation. In the equation shown in the figure, we represent in the indicator ion with without extra base being added. This is called the equivalence point because you have
added equal moles of acid and base. For most acids and bases, this point is difficult to see,

186 187
because the acid and base reactants as well as the salt and water products have no color. This
is where indicators come in. An indicator is used to determine the equivalence of the
titration. A few drops of the indicator are added to the flask before you begin the titration. If
an appropriate indicator has been chosen, the indicator will only react and change color (and
Lesson Summary
stay color changed) when all of the other acid has reacted. Therefore, the indicator will
change color immediately after enough base was added to completely react with all of the An indicator is a substance that changes color at a specific pH and is used to
acid (the equivalence point). indicate the pH of the solution.
A titration is the addition of a known concentration of base (or acid) to a solution
Some laboratories have pH meters that measures this point more accurately than the
of acid (or base) of unknown concentration.
indicator, although an indicator is much more visual. The main purpose of a pH meter is to
The equivalence point is the point in the titration where the number of moles of acid
measure the changes in pH as the titration goes from start to finish. It is also possible to
equals the number of moles of base, and, if you chose an appropriate indicator, where
determine the equivalence point using the pH meter as the pH will change dramatically once
the indicator changes color.
all of the acid and base have been neutralized. + -
For titrations where the stoichiometric ratio of mol H : mol OH is 1:1, the
The Mathematics of Titration formula (Ma)(Va)=(Mb)(Vb) can be used to calculate concentrations or volumes
For the calculations involved here, we will restrict our acid and base examples where for the unknown acid or base.
+ -
the stoichiometric ratio of H and OH is 1:1. The formula for these 1:1 reactions, in which Vocabulary
1 mole of acid is needed to react with 1 mole of base, has the structure:
(Ma)(Va)=(Mb)(Vb) Titration: the lab process in which a known concentration of base (or acid) is added to
Where a solution of acid (or base) of unknown concentration
Ma is the molarity of the acid Indicator: a substance that changes color at a specific pH and is used to indicate the
Va is the volume of the acid pH of the solution
Mb is the molarity of the base Equivalence point: the point in the titration where the number of moles of acid equals
Vb is the volume of the base the number of moles of base
This equation works because the left side calculates the number of moles of acid which react
and the right side calculates the number of moles of base. To reach the equivalence point, 8.4: Review Questions
mol acid = mol base. 1) What is an indicator? What is it used for?
2) What is an equivalence point?
Example: When 10.0 mL of a 0.125 M solution of hydrochloric acid, HCl, is titrated with a 3) If 22.50 mL of a sodium hydroxide is necessary to neutralize 18.50 mL of a 0.1430
0.100 M solution of potassium hydroxide, KOH, what the volume of the hydroxide solution M HNO3 solution, what is the concentration of NaOH?
is required to neutralize the acid? 4) Calculate the concentration of hypochlorous acid if 25.00 mL of HClO is used in a
Solution: titration with 32.34 mL of a 0.1320 M solution of sodium hydroxide.
Step 1: Write the balanced ionic chemical equation. Check that the acid:base ratio is 1:1. 5) What volume of 0.45 M hydrochloric acid must be added to 15.0 mL of .997
HCl + KOH → H2O + KCl M potassium hydroxide to neutralize the base? (HCl + KOH → H2O + KCl)
Since 1 HCl is needed for each KOH, the reaction is 1:1. 6) What volume of .20 M HI is needed to neutralize 25 mL of .50 M KOH?
Step 2: Use the formula and fill in all of the given information. The acid is HCl and the base 7) What is the molarity of sodium hydroxide if .174L of the solution is neutralized by .20L
is KOH. of 1.2 M HCl? (HCl + NaOH → H2O + NaCl)
Ma=0.125 M 8) Suppose we used .150L of 0.500M NaOH and .250L of vinegar (acetic acid solution) of
Va=10.0 mL an unknown concentration. What is the molarity of the vinegar? (Balanced reaction is:
Mb=0.100 M NaOH(aq) + HC2H3O2 (aq) → NaC2H3O2 (aq) + H2O(l))
Vb=?

(Ma)(Va)=(Mb)(Vb)
(0.125 M)(10.0 mL)=(0.100 M)(Vb)
Vb=12.5 mL
Therefore, for this weak acid-strong base titration, the volume of base required for the
titration is 12.5 mL.

188 189
 Chapter 9: Energy of Chemical Changes
different amounts of potential energy because they are made up of different atoms, and those
9.1: Energy atoms have different positions relative to one another.
Since different chemicals have different amounts of potential energy, scientists will
Objectives: sometimes say potential energy depends on not only position but also composition.
Distinguish between kinetic and potential energy and give examples of each. Composition affects potential energy because it determines which molecules and atoms
end up next to each other. For example, the total potential energy in a cup of pure water is
Introduction different than the total potential energy in a cup of apple juice because the cup of water and
Just like matter, energy is a term that we are all familiar with and use on a daily basis. the cup of apple juice are composed of different amounts of different chemicals.
Before you go on a long hike, you eat an energy bar; every month, the energy bill is paid; on
The Law of Conservation of Matter and Energy
TV, politicians argue about the energy crisis. But what is energy? If you stop to think about
it, energy is very complicated. When you plug a lamp into an electric socket, you see energy While it’s important to understand the difference between kinetic energy and
in the form of light, but when you plug a heating pad into that same socket, you only feel potential energy, the truth is energy is constantly changing. Kinetic energy is constantly
warmth. Without energy, we couldn’t turn on lights, we couldn’t brush our teeth, we couldn’t being turned into potential energy, and potential energy is constantly being turned into
make our lunch, and we couldn’t travel to school. In fact, without energy, we couldn’t even kinetic energy. Even though energy can change form, it must still follow the fundamental
wake up because our bodies require energy to function. We use energy for every single thing law: energy cannot be created or destroyed, it can only be changed from one form to another.
that we do, whether we're awake or asleep. This law is known as the law of conservation of energy.

Types of Energy: Kinetic and Potential 9.2: Endothermic and Exothermic Changes
Kinetic energy is energy associated with motion. When an object is moving, it has Objectives
kinetic energy, and when the object stops moving, it has no kinetic energy. Although all Define potential energy and kinetic energy.
moving objects have kinetic energy, not all moving objects have the same amount of Define endothermic and exothermic reactions.
kinetic energy. The amount of kinetic energy possessed by an object is determined by its Describe how heat is transferred in endothermic and exothermic reactions.
mass and its speed. The heavier an object is and the faster it is moving, the more kinetic Determine whether a reaction is endothermic or exothermic through
energy it has. Kinetic energy is very common and is easy to spot in the world around you. observations, temperature changes, or an energy diagram.
Sometimes we even capture kinetic energy and use it to power things like our home
appliances. Forms of kinetic energy include heat, light, sound, and electricity. All Chemical Reactions Involve Energy
Potential energy is stored energy that remains available until we choose to use it. Remember that all chemical reactions involve a change in the bonds of the reactants.
Think of a battery in a flashlight. If you leave a flashlight on, the battery will run out of The bonds in the reactants are broken and the bonds of the products are formed. Chemical
energy within a couple of hours. If, instead, you only use the flashlight when you need it bonds have potential energy or "stored energy". Because we are changing the bonding, this
and turn it off when you don’t, the battery will last for days or even months. Because the means we are also changing how much of this “stored energy” there is in a reaction.
battery stores potential energy, you can choose to use the energy all at once, or you can save When chemical reactions occur, the new bonds formed never have exactly the same
it and use a small amount at a time. amount of potential energy as the bonds that were
Any stored energy is potential energy and has the “potential” to be used at a later broken. Therefore, all chemical reactions involve
time. Unfortunately, there are a lot of different ways in which energy can be stored, making energy changes. Energy is either given off by the
potential energy very difficult to recognize. Generally speaking, an object has potential reaction or energy is taken in by the reaction. There
energy due to its position relative to another object. are many types of energy that can be involved in
For some examples of potential energy, though, it’s harder to see how “position” is these changes. Different types of energy include:
involved. In chemistry, we are often interested in what is called chemical potential energy. Heat
Chemical potential energy is energy stored in the atoms, molecules, and chemical bonds Electricity
that make up matter. How does this depend on position? The world and all of the chemicals Light
in it are made up of atoms. These atoms store potential energy that is dependent on their Chemical potential energy
positions relative to one another. Although we cannot see atoms, scientists know a lot about In endothermic changes, kinetic energy
Sometimes the products have more energy (such as heat) is absorbed and changed
the ways in which atoms interact. This allows them to figure out how much potential energy stored in their bonds than the reactants had to start into chemical potential energy.
is stored in a specific quantity of a particular chemical. Different chemicals have with. This means that the reaction started with less In exothermic changes, chemical
hidden energy than we had at the end. Where did this potential energy is released as heat or
extra energy come from? In these reactions, heat or other kinetic energy.

190 191
*If more reactant is added or a product
is removed, the reaction must shift to
make more products to get back to
equilibrium.

*If more product is added or a reactant

is removed, the reaction must shift to
make more reactants to get back to
equilibrium.
CC Tracy Poulsen
 The color change of an indicator occurs over a very short range.

Hydrangeas are also a natural indicator. The

petals will change colors based on the pH of
the soil.
was called the torr in honor of Torricelli. 760 torr is exactly the same as 760 mmHg. For
certain work, it became convenient to express gas pressure in terms of multiples of
normal atmospheric pressure at sea level and so the unit atmosphere (atm) was
calculations you do dealing with the kinetic energy of molecules is done with
introduced. The conversion you need to know between various pressure units are:
Kelvin temperatures.
1.00 atm = 760. mmHg = 760. torr Some important principles can be derived from this relationship:
1. All gases at the same temperature have the same average kinetic energy.
Example: Convert 425 mmHg to atm. 2. Heavier gases must move more slowly in order to have the same kinetic energy as
Solution lighter gases.
The conversion factor is 760. mmHg = 1.00 atm
Example: If molecules of H2, O2, and N2 are all placed in the same container at the same
temperature, which molecules will have the greatest velocity?
Solution: Because they are at the same temperature, they will have the same energy.
This example shows how to perform this conversion using dimensional analysis. If However, lighter particles must move faster in order to have the same kinetic energy. We
you are the memorizing type, you can just memorize that to convert from mmHg to atm must, therefore, look at their masses. Use your periodic table: Mass of H2 = 2(1.008
you must divide by 760. g/mol) = 2.016 g/mol
Mass of O2 = 2(16.00 g/mol) = 32.00 g/mol
Gas Temperature and Kinetic Energy Mass of N2 = 2(14.01 g/mol) = 28.02 g/mol
Kinetic energy is the energy of motion and therefore, all moving objects have kinetic
energy. The mathematical formula for calculating the kinetic energy of an object is KE=1/2 Because H2 is the lightest, it must have the greatest velocity in order to have the same
2
mv , where m is the mass and v is the velocity of the object or particle. This physics formula energy (the same temperature) as the other gases.
applies to all objects in exactly the same way whether we are talking about the moon moving
in its orbit, a baseball flying toward home plate, or a gas molecule banging around in a Section Summary
bottle. All of these objects have kinetic energy and their kinetic energies can all be calculated The collisions between molecules are perfectly elastic. The phrase “perfectly elastic
with the same formula. The kinetic energy of a molecule would be calculated in exactly this collision” comes from physics and means that kinetic energy is conserved in
same way. You should note that if the mass of an object is doubled while its velocity remains collisions.
the same, the kinetic energy of the object would also be doubled. If, on the other hand, the
The molecules of an ideal gas have no attraction or repulsion for each other.
velocity is doubled while the mass remains the same, the kinetic energy would be quadrupled
because of the square in the formula. At any given moment, the molecules of a gas have different kinetic energies. We
When you measure the temperature of a group of molecules, what you are actually deal with this variation by considering the average kinetic energy of the molecules.
measuring is their average kinetic energy. They are the same thing but expressed in The average kinetic energy of a group of molecules is measured by temperature.
different units. The formula for this relationship is KEave=3/2RT where R is the gas Molecules of a gas are so far apart, on average, that the volume of the molecules
constant and T is the absolute temperature, measured in Kelvin. When a substance is heated, themselves in negligible compared to the volume of the gas.
the average kinetic energy of the molecules is increased. Since the mass of the molecules Molecular collisions with container walls cause the gas to exert pressure.
cannot be increased by heating, it is clear that the velocity of the molecules is increasing. Because of the molecular motion of molecules, they possess kinetic energy at all
Remember, the motion of molecules is related to their temperature. If you think of temperatures above absolute zero.
the average kinetic energy of a group of molecules and temperature measured in degrees Temperature is directly proportional to the average kinetic energy of gas molecules.
Kelvin, the relationship is a direct proportion. That means that if the temperature, in Kelvin, Lighter gases will have higher velocities than heavier gases, at the same
is doubled the kinetic energy of the particles is also doubled. It is absolutely vital that you temperature and pressure.
keep in mind that the mathematical relationship between the temperature and the average In the Kelvin scale, 0 K means the particles have no kinetic energy. Doubling the
kinetic energy of molecules only exists when the temperature is expressed in the Kelvin temperature in Kelvin doubles the kinetic energy of particles.
scale. In order for the direct proportion to exist, the molecules must have zero kinetic energy
Real gases tend to deviate from ideal gases at high pressures and low temperatures, as
when the temperature is zero. The temperature at which molecular motion stops is 0 K (-273
the attractive forces between molecules and the volume of gas molecules becomes
C). It is surely apparent to you that molecules do NOT have zero kinetic energy at 0 C.
significant
Balloons and automobile tires do not go flat when the outside temperature reaches 0 C. If
temperature is measured in Kelvin degrees, then the average kinetic energy of a substance at Vocabulary
100 K is exactly double the average kinetic energy of a substance at 50 K. Make sure all the
Kelvin temperature: The absolute temperature scale where 0 K is the theoretical
absence of all thermal energy (no molecular motion).
Kinetic energy: Kinetic energy is the energy a body possesses due to it
2
motion, KE=1/2mv .

224 225
 Kinetic theory: used to explain how properties of gases the cylinder can be measured. The amount of gas inside the cylinder cannot change and
Pressure: a measure of the force with which gas particles collide with the walls the temperature of the gas is not allowed to change.
of their containers In the picture on the right, the volume of the gas is 4.0 L and the pressure exerted by
Temperature: a measurement of the kinetic energy of particles the gas is 2.0 atm. If the piston is pushed down to
decrease the volume of the gas to 2.0 L, the pressure of
the gas is found to increase to 4.0 atm. The piston can be
11.2: Gas Laws moved up and down to positions for several different
volumes and the pressure of the gas read at each of the
Objectives
volumes.
Predict effect on pressure, volume, or temperature if one of the other variables are We might note from casual observation of the
changed. data that doubling volume is associated with the
Solve problems using the combined gas law pressure being reduced to half and if we move the piston
Volume and temperature (in K) are
?maybe do combined first and get other gas laws after directly related. to cause the pressure to double, the volume is halved.
Introduction The data show that the relationship is an inverse The relationship between volume and
pressure is an inverse relationship.
Gases are often characterized by their volume, temperature, and pressure. These relationship, meaning that as volume increases the
characteristics, however, are not independent of each other. Gas pressure is dependent on the pressure decreases. The opposite is also true.
force exerted by the molecular collisions and the area over which the force is exerted. The
force exerted by the molecular collisions is dependent on the absolute temperature and so Boyle’s Law can be summarized in the following equation:
forth. The relationships between these characteristics can be determined both experimentally
and logically from their mathematical definitions. Where:
The gas laws are mathematical relationships that exist for gases between the
P1=the initial pressure
volume, pressure, temperature, and quantity of gas present. They were determined
V1=the initial volume
experimentally over a period of 100 years. They are logically derivable from our present
day definitions of pressure, volume, and temperature. P2=the final pressure
V2=the final volume
Boyle’s Law: Pressure vs. Volume For this equation, the units used for pressure are unimportant, as long as both pressures have the
The relationship between the pressure same unit (either mmHg or atm) and each volume has the same unit (either mL or L).
and volume of a gas was first determined
Charles’s Law: Temperature and Volume
experimentally by an Irish chemist named Robert
Boyle (1627-1691). The relationship between the The relationship between the volume and temperature of a gas was investigated by a
pressure and volume of a gas is commonly French physicist, Jacques Charles (1746-1823). (As a piece of trivia, Charles was also the
referred to as Boyle’s Law. first person to fill a large balloon with hydrogen gas and take a solo balloon flight.) The
When we wish to observe the relationship relationship
between two variables, it is absolutely necessary to between the volume and temperature of a gas is
keep all other variables constant so that the change PV Data often referred to as Charles’s Law.
in one variable can be directly related to the TrialVolumePressure An apparatus that can be used to study the
change in the other. Therefore, when the relationship between the temperature and volume of
1 8.0 L 1.0 atm
relationship between gas volume and gas pressure a gas is shown in the picture to the right. Once
2 4.0 L 2.0 atm
is investigated, the quantity of gas and its again, we have a sample of gas trapped inside a
3 2.0 L 4.0 atm cylinder so no gas can get in or out. Thus we have a
temperature must be held constant so these factors
do not contribute to any observed changes. 4 1.0 L 8.0 atm constant mass of gas. We also have a mass set on
You may have noticed that when you try to Volume and pressure data for a gas sample. top of a moveable piston to keep a constant force
squeeze a balloon, the resistance to squeezing is pushing against the gas. This guarantees that the gas
greater as the balloon becomes smaller. That is, the pressure inside the balloon becomes greater pressure in the cylinder will be constant because if The picture on the left shows the volume of a
sample of gas at 250. K and the picture on
when the volume is reduced. This phenomenon can be studied more carefully with an apparatus the pressure inside increases, the piston will be the right shows the volume when the
like that in Figure 9. This is a cylinder tightly fitted with a piston that can be raised or lowered. pushed up expanding inside volume until the inside temperature has been raised to 500 K.
There is also a pressure gauge fitted to the cylinder so that the gas pressure inside pressure becomes equal to outside pressure again.

226 227
Similarly, if the inside pressure decreases, the outside pressure will push the cylinder down,
decreasing volume, until the two pressures again
become the same. This system guarantees constant Standard Temperature and Pressure (STP)
gas pressure inside the cylinder.
It should be apparent by now that expressing a quantity of gas simply by stating its
This relationship is a direction relationship. If
volume is totally inadequate. Ten liters of oxygen gas could contain any mass of oxygen
the temperature, in Kelvin, doubles, so does the volume.
from 4000g to 0.50g depending on the temperature and pressure of the gas. Chemists have
This relationship would also be expected when we
found it useful to have a standard temperature and pressure with which to express gas
recognize that we are increasing the total force of
volume. The standard conditions of temperature and pressure (STP) were chosen to be 0 C
molecular collisions with the walls by raising the
(273 K) and 1.00 atm (760 mmHg). You will commonly see gas volumes expressed as 1.5L
temperature and the only way to keep the pressure from
at STP. Once you know the temperature and pressure conditions of a volume of gas, you can
increasing is to increase the area over which that larger
calculate the volume at other conditions and you can also calculate the mass of the gas if
force is exerted. This mathematical relationship is
you know the formula.
known as a direct proportionality. When one variable is
increased, the other variable also increases by exactly The Combined Gas Law
the same factor. An equation to show how these values are related is given by:
Boyle’s Law shows how the volume of a gas changes when its pressure is changed
(temperature held constant) and Charles’s Law shows how the volume of a gas changes
when the temperature is changed (pressure held constant). Is there a formula we can use to
calculate the change in volume of a gas if both pressure and temperature change? The answer
This relationship is ONLY true if the temperature is measured in Kelvin. However, the units is “yes”, we can use a formula that combines Boyle’s Law and Charles’s Law.
of volume are irrelevant, as long as the two volumes are measured in the same units. This equation is most commonly written in the from shown below and is known
as the Combined Gas Law.
Gay-Lussac’s Law: Temperature and Pressure
The relationship between temperature and
pressure was investigated by the French chemist, Pressure vs. Temperature Data
Joseph Gay-Lussac (1778-1850). In an apparatus Trial Temperature Pressure As in the other laws, when solving problems with the combined gas law, temperatures must
used for this investigation, the cylinder does not have 1 200. K 600. mmHg always be in Kelvin. The units for pressure and volume may be any appropriate units but the
a moveable piston because it is necessary to hold the 2 300. K 900 mmHg units for each value of pressure must be the same and the units for each value of volume must
volume constant as well as the quantity of gas. This 3 400. K 1200 mmHg be the same.
apparatus allows us to alter the temperature of a gas 4 500. K 1500 mmHg Another interesting point about the combined gas law is that all the other gas laws
and measure the pressure exerted by the gas at each Temperature and pressure data. Note (Charles’, Gay-Lussac’s, and Boyle’s) can be derived from this equation. To do this, you
temperature. that if the temperature doubles from simply cancel out the variable that was held constant in the reaction. For example,
After a series of temperatures and pressures 200. K to 400. K, the pressure also temperature is constant in Boyle’s Law. If you cancel the temperature’s out of Boyle’s Law,
have been measured, a data table like the others can doubles. you get:
be produced.
Temperature and pressure are also directly Although the other equations are not as obvious, the same method can be used to derive the
related, meaning that if the temperature, in Kelvin, other equations. If you are able to derive the other equations, you will not have to memorize
doubles, so does the pressure. This relationship is also them.
logical since by increasing temperature, we are
increasing the force of molecular collision and keeping Example: A sample of gas has a volume of 400. liters when its temperature is 20. C and its
the area over which the force is exerted constant requires pressure is 300. mmHg. What volume will the gas occupy at STP?
that the pressure increases. Solution:
Step 1: Identify the given information & check units. Temperature must be in Kelvin.
Pressure and temperature (in K) are Volume units must match and pressure units must match.
directly related in Gay-Lussac’s P1=300 mmHg
Law V1=400. L
T1=293 K (remember, ALL temperatures must be in Kelvin)
P2=760 mmHg (standard pressure)
V2=?

228 229

T2=273 K
Step 2: Solve the combined gas law for the unknown variable.

that equal volumes of gas under the same conditions of temperature and pressure contain
the same number of molecules.
V2=147 L This relationship is important for a couple of reasons. It means that all gases under
the same conditions behave the same way: all of these equations work equally well for
Example: A sample of gas occupies 1.00 under standard conditions. What temperature carbon dioxide, helium, or a mixture of gases. Furthermore, we will be able to use this
relationship again when we deal with balanced reactions. The volume of two gases at the
would be required for this sample of gas to occupy 1.50 L and exert a pressure of 2.00 atm?
same temperature and pressure are directly related to the number of molecules (or moles)
Solution:
of the gases involved in a chemical reaction.
Step 1: Identify the given information & check units. Temperature must be in Kelvin.
Volume units must match and pressure units must match. Section Summary
P1=1.00 atm (standard pressure)
V1=1.00 L For a fixed sample of ideal gas at constant temperature, volume is inversely
T1=273 K (standard temperature, remember, ALL temperatures must be in Kelvin) proportional to pressure.
P2=2.00 atm For a fixed sample of ideal gas at constant pressure, volume in directly proportional
V2=1.50 L to temperature.
T2=? For a fixed sample of ideal gas at constant volume, pressure is directly proportional to
Step 2: Solve the combined gas law for the unknown variable. temperature.
The volume of a mass of gas is dependent on the temperature and pressure. Therefore,
these conditions must be given along with the volume of a gas.
T2=819 K
Standard conditions of temperature and pressure are 0 C and 1.0 atm.
Avogadro's Law: Equal volumes of gases under the same conditions of temperature
Example: A sample of gas has a volume of 500.mL under a pressure of 500.mmHg. What
and pressure contain equal numbers of molecules.
will be the new volume of the gas if the pressure is reduced to 300.mmHg at constant
temperature? Further Reading / Supplemental Links
Solution:
Section 7-6 is on the Combined Gas
Step 1: Identify the given information & check units. Temperature must be in Kelvin.
Law. http://www.fordhamprep.org/gcurran/sho/sho/Sections/Section31.htm
Volume units must match and pressure units must match.
http://en.wikipedia.org/wiki/Kinetic_theory;
P1=500. mmHg
http://www.chm.davidson.edu/chemistryapplets/kineticmoleculartheory/basicconcept
V1=500. mL
P2=300. mmHg s.html
V2=?
Temperature is constant, so it cancels out of the combined gas law.
Step 2: Solve the combined gas law for the unknown variable. (Or, recognize this is 11.3: Ideal Gas Law
Boyle’s Law and start with that equation.) Objectives
Solve problems using the ideal gas law, PV=nRT.

V2=833 mL Introduction
The individual gas laws and the combined gas law all require that the quantity of gas
Avogadro’s Law remain constant. The Universal Gas Law (also sometimes called the Ideal Gas Law) allows
Avogadro’s Law was known as Avogadro’s hypothesis for the first century of its us to make calculations on different quantities of gas as well.
existence. Since Avogadro's hypothesis can now be demonstrated mathematically, it was
decided that it should be called a law instead of a hypothesis. Avogadro’s Law postulates The Universal Gas Law Constant
We have considered four laws that describe the behavior of gases: Boyle’s Law,
Charles’s Law, Avogadro’s Law, and Gay-Lussac’s Law. These three relationships, which
show how the volume of a gas depends on pressure, temperature, and the number of moles of
gas, can be combined to form the ideal gas law:

230 231
Where each variable and its units are:
P=pressure (atm) Section Summary
V=volume (L) The Universal Gas Law: PV=nRT
n=number of moles of gas (mol) At STP, one mole of any gas occupies 22.4
T=temperature (K) The universal gas law is often used along with laboratory data to find the molar
R=ideal gas constant = 0.0821 atm‫ڄ‬L/mol‫ڄ‬K mass of an unknown substance.
Up to this point in gas law calculations, we haven’t worried too much about which All images, unless otherwise stated, are created by the CK-12 Foundation and are under
unit you use for pressure and volume as long as the units matched. Notice that the gas the Creative Commons license CC-BY-NC-SA.
constant, R, has specific units. Your units of pressure and volume must be in atm and L,
respectively, because they must match the appropriate units in the constant, R. Moles, of
course, always have the unit moles and temperature must always be Kelvin. You can convert
the value of R into values for any set of units for pressure and volume, if you wanted, but the
numerical value of R would also change.

Example: A sample of nitrogen gas, N2, has a volume of 5.56 L at 0 C and 1.50
atm pressure. How many moles of nitrogen are present in this sample?
Solution:
Step 1: Identify the given information & check units. Temperature must be in Kelvin.
Volume and pressure units must match R.
P=1.50 atm
V=5.56 L
n=?
T=273 K (must be in K)
Step 2: Solve the ideal gas law for the unknown variable.

n=0.372 mol

Example: 2.00 mol of methane gas, CH4, are placed in a rigid 500. mL container and
heated to 100. C. What pressure will be exerted by the methane?
Solution:
Step 1: Identify the given information & check units. Temperature must be in Kelvin.
Volume and pressure units must match R.
P=?
V=500 mL = 0.500 L
n=2.00 mol
T=100 C = 373 K
Step 2: Solve the ideal gas law for the unknown variable.

P=122 atm

232 233
Answers to Selected Problems
Section 2.3
1) T
2) F
Section 1.1 3) F
1) B 4) T
2) A 5) B
3) D 6) C
4) D 7) A
5) B 8) T
6) Whether or not the plants 9) T
received fertilizer 10) T
7) Growth (height) of plants b) 11) F
8) Amount of sun, amount of water, c) 2.6 mmHg/°C 12) F
type of plant (corn) b) d) About 830 mmHg 13) Ru
c) 11.3 g/mL e) About 1000 mmHg 14) Zn
Section 1.2 d) 27 g 15) Kr
1) D e) 5.2 mL Section 2.1 16) B
2) A 4) Exact graphs and answers may 1) C 17) E
3) D vary, but should look similar to the 2) Dalton had experimental evidence 18) A
4) A following to support his claims. Democritus 19) D
5) A a) Independent variable: concentration did not. 20) C
58
6) D of ammonia (mol/L). Dependent 3) No! Inaccurate theories give scientists 21) 26Fe
19
7) D variable: time (s) an idea to build from and a way to 22) 9F
+ 0 -
8) F test other ideas and develop 23) p =2, n =2, e =2
+ 0 -
experiments. Most current ideas are 24) p =11, n =12, e =11
+ 0 -
Section 1.3 adaptations of previous ideas. 25) p =1, n =0, e =1
+ 0 -
1) The independent variable is the label 26) p =26, n =29, e =26
+ 0 -
of the first column. The dependent Section 2.2 27) p =17, n =20, e =17
+ 0 -
variable is the label of the last 1) F 28) p =5, n =6, e =5
+ 0 -
column(s). 2) T 29) p =92, n =146, e =92
2) The independent variable is the label 3) T + 0 -
30) p =92, n =143, e =92
of the x-axis or the key. The 4) F
dependent variable is the label of the 5) D Section 2.4
y-axis. 6) A 1) 63.55 amu
3) Exact graphs and answers may b) 7) E 2) 35.49 amu
vary, but should look similar to the c) About 47 seconds 8) C
following d) About 1.45 mol/L 9) B Section 2.5
a) Independent variable: volume of 5) Exact graphs and answers may 10) Nucleus 1) Red
solution (mL); dependent: mass vary, but should look similar to the 11) Repelled by…attracted to 2) As the energy of a wave increases,
of solution (g) following 12) Location of positive mass in atoms frequency increases. As the energy of
a) Independent variable: temperature 13) C a wave increases, the wavelength
(°C); dependent variable: pressure 14) If the particles hit the positive decreases.
(mmHg) central mass they would bounce off. 3) The wave on the left has more energy,
If they missed the central positive because it has a shorter wavelength.
part, they would go straight through. 4) Radio, infrared, visible, UV, gamma

234 235
1) Quantized means to have specific
5) Red, orange, yellow, green, blue, violet Section 2.6 amounts of energy. Bohr said
 electrons can have only specific
amounts of energy and are,
therefore, quantized. 2 2 3
11) 1s 2s 2p 6) Each element has a different number
2) Because each element has a different 2 2 6 2 6 2 10 6 2 1
12) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d of protons, and elements are now Section 3.4
spectrum, you can use it to identify 2 2 6 2 6 2 10 6 2 10
which elements are present 13) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d arranged in order of increasing 1) C
2
3) The sun gives off the specific pattern 5p atomic number instead of increasing 2) B
2 2 6 2 6 2 10 6
of light unique to helium. No other 14) 1s 2s 2p 3s 3p 4s 3d 4p atomic mass. 3) C
2 2 6 2 6 2 10 6 2 10
element produces that pattern of light. 15) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 7) There are more elements, the table is 4) Halogen
6 2 14 10 6
4) Electrons gain energy and move to 5p 6s 4f 3d 6p in order of increasing atomic number 5) Noble gas
2 2 6 2 6 2 10 6 2 10
16) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d instead of mass, the family of noble 6) Alkaline earth metal
higher energy levels. When electrons 6 2 14 10 6 2 14
lose energy they move to lower energy 5p 6s 4f 3d 6p 7s 5f gases has been added, the table has 7) Alkali metal
2 5
levels. 17) [He] 2s 2p been turned sideways from its 8) Halogens
2 1
5) The electrons give off the extra energy 18) [Ne] 3s 3p original form. 9) Halogens
2 2
as light. 19) [Ar] 4s 3d 10) Noble gas
2 10 3 Section 3.2
6) Each element has different possible 20) [Ar]4s 3d 4p 11) Transition metals
1
energy levels which its electrons can 21) [Kr]5s 1) Nonmetal 12) Alkaline earth metals
2 2
occupy, so there are different possible 22) [He] 2s 2p 2) Nonmetal 13) Alkali metals
“drops” electrons can make, producing 3) Nonmetal
Section 3.1 4) Metal Section 3.5
different spectra.
7) Electrons start at ground state (lowest 1) Mendeleev first put the elements in 5) Metal 1) B
energy level possible). When they are order from lightest mass to heaviest 6) Nonmetal 2) D
given energy as light, heat, or mass. Then he put elements with 7) Metal 3) A
similar properties in the same group. 8) Metalloid 4) Ga
electricity, the electrons may move up
to a higher energy level (excited state). 2) Mendeleev left room for undiscovered 9) Nonmetal 5) K
The electrons will drop back to lower elements, he didn’t force elements into 10) Metal 6) Ti
energy levels releasing the extra groups which didn’t have similar 11) Nonmetal 7) Iodine has electrons in a higher energy
energy as a photon (or piece of light). properties even if the mass didn’t 12) Metal level further from the nucleus than
follow his original pattern 13) Metalloid bromine.
Section 2.7 3) Yes, unlike other methods of 14) Mercury has properties of metals (such 8) B, Al, Ga, In, Tl
1) 2 organization. As new elements were as being malleable, lustrous, 9) C, Ga, Sn
2) p discovered there was room for them conductivity, ductile, etc) and bromine 10) The energy required to remove the
3) d in Mendeleev’s table. has properties of nonmetals (such as electron furthest from the nucleus
4) 10 4) By looking how the properties such as brittle, insulator, etc). Even though 11) Na, Mg, S, Ar
5) 2p melting point, density, etc, changed as they are both liquids, their other 12) The relative attraction for electrons
6) 3 he went down a group/family or properties place them as metal for in a bond (how hard an atom pulls on
7) Al across a row of his periodic table, he mercury and nonmetal for bromine. electrons in a bond)
8) 4d predicted what numbers would fit the 13) Li
pattern. Section 3.3 14) Na
9) 4s
2 2 6 2 5) Sometimes the next heaviest element 1) 5 15) K
10) 1s 2s 2p 3s
didn’t fit according to properties in the 2) 3 16) Mg
next available place. He either traded 3) 5
the order of the neighboring element 4) 5 Section 4.1
(such as what he did for I and Te) or 5) 1 1) Compound
he left blank spaces to put elements in 6) 1 2) Mixture
the appropriate group (such as leaving 7) 3 3) Element
holes where Ga and Ge currently are 8) 3 4) Compound
placed). 9) 8 5) Mixture
10) 8 6) Mixture

236 237
7) Compound 16) The names of transition ions include 26) FeCl2
8) Element the charge of the ion, because they can 27) Cu(NO3)2
9) Mixture form ions with more than one charge. 28) MgO
10) Mixture 17) Copper(II) 29) CaO
18) Cobalt(II) 30) CuBr
Section 4.2 19) Cobalt(III) 31) Al2S3
1) Elements are most stable with eight 20) Copper(I) 32) H2CO3
valence electrons. 21) Nickel(II) 33) KMnO4
2) Nonmetals can form covalent bonds. 22) Chromium(III) 34) Cu2Cr2O7
Metals cannot. 23) Iron(II) 35) FeCl3
3) Covalent 24) Iron(III) 36) FeSO4 8)
4) Ionic 25) A group of atoms which together hold
5) Covalent a charge Section 4.5 9)
6) Ionic 26) Nitrate 1) Potassium chloride
7) Ionic 27) Acetate 2) Magnesium oxide
10)
8) Covalent 28) Hydroxide 3) Copper(II) sulfate
9) Ionic 29) Phosphate 4) Sodium chloride
10) Metallic 30) Sulfate 5) Cobalt(II) bromide
11) Covalent 31) Carbonate 6) Magnesium fluoride 11)
12) Covalent 7) Nickel(II) hydroxide
13) Covalent Section 4.4 8) Sodium acetate
14) Metallic 1) NaNO3 9) Copper(II) oxide 12)
15) Ionic 2) Ca(NO3)2 10) Iron(II) chloride
16) Ionic 3) Fe(NO3)3 11) Lithium chloride 13)
17) Metallic 4) Na2SO4 12) Magnesium bromide
18) 1=sucrose; 2=sodium chloride, 3=zinc 5) CaSO4 13) Ammonium hydroxide 14)
6) Fe2(SO4)3 14) Copper(I) oxide
Section 4.3 7) NaCl 15) Calcium fluoride
1) An atom or group of atoms with a 8) CaCl2 16) Potassium carbonate 15)
charge 9) FeCl3 17) Sodium chloride
2) Positive, metal atoms lose electrons to 10) Na3PO4 18) Lead(II) oxide
form positive ions 11) Ca3(PO4)2 19) Calcium nitrate
3) Negative, nonmetals will gain 12) FePO4 20) Magnesium hydroxide 16)
electrons to form negative ions 13) NaOH 21) Tin(IV) oxide
4) -1, chloride 14) Ca(OH)2 Section 4.7
5) -1, bromide 15) Fe(OH)3 Section 4.6 1) Tetrahedron
6) -3, nitride 16) Na2CO3 1) Covalent 2) Tetrahedron
7) -2, oxide 17) CaCO3 2) Ionic 3) Tetrahedron
8) +2, calcium 18) Fe2(CO3)3 3) Ionic 4) Bent
9) -1, fluoride 19) MgS 4) Covalent 5) Trigonal pyramid
10) +2, magnesium 20) Pb(NO3)2 5) Ionic 6) Tetrahedron
11) +1, lithium 21) Na2O 6) Ionic
12) -1, iodide 22) Ca(OH)2 Section 4.8
13) +1, sodium 23) K2CO3 7) 1) In a nonpolar covalent bond,
14) +1, potassium 24) AlBr3 electrons are evenly shared between
atoms. In a polar covalent bond
15) +3 aluminum 25) Fe(NO3)3 electrons are not evenly shared
resulting in a partial positive and
partial negative side. In an ionic
238 bond, electrons are shared at all but
one atom loses electrons to another
atom forming particles with full
charges.
2) P-Cl is more polar than S-Cl, because there is a bigger difference in
239

240 6) 23000
7) 0.0009156
8) .0072
electronegativities. P is less 9) 8,255,000
electronegative than S, so Cl is able to Section 5.2 10) 7.3(EE)14
pull the electrons further from P than S 1) W 11) 6.01(EE)(-)6
making it more polar. h 12) 7.98(EE)5
3) Electrons are not evenly shared e 13) 6.0x10
7
4) N, O, or F n -2
14) 6.67x10 or 0.067
5) Hydrogen bonding is a stronger w -2
15) 1.4x10 or 0.014
attraction between molecules with or 16) 9.13x10
-5
partial charges than the attraction ki
between polar molecules. n Section 5.3
6) Hydrogen bonding is a strong g 1) 22.9 cm
attraction between neighboring w 2) 48 min
molecules in which H is bonded to it 3) 296 g
N, O, or F. h 4) $5.48
7) Polar re 5) An
8) Nonpolar al sw
9) Polar ly ers
10) Polar and hydrogen bonding la var
11) Polar and hydrogen bonding rg y.
12) Nonpolar e A
13) Ammonia or 12
re 0
14) Water lb
15) Ammonia al
per
16) Ammonia ly so
s n
Section 5.1 m has
1) Based on the decimal (10) system; al a
used internationally; units are based on l ma
physical constants n ss
2) 1kg u of
3) 10 g m 5.4
b 5
4) 100cg x1
5) 1000 er 4
6) cL s 0
a g.
7) 100
n 15
8) Scientists need to use the same unit of 0
measurement so they can share d
m lb
information, data, and calculations is
more effectively. ea 4 4
6.82x10 g. 175 lbs is 7.95x10 g.
9) Meter (m) su 6) 7.5 miles
10) Kilogram (kg) re 7) 249 min
11) Liter (L); the liter is the volume of a m 6
8) 3.64x10 g
10cm x 10cm x 10cm container or 1dm e 9) 15.7 km/L
x 1 dm x 1 dm nt 4
10) 3.2x10 miles/hr
12) NO! This is only 2°C (almost as cold s
-5
as ice water). 2) 4.79x10 Section 5.4
3
3) 4.26x10 23
1) 1.5x10 molecules H2O
9
4) 2.51x10 21
2) 2.71x10 molecules Al2(CO3)3
-3
5) 2.06x10
 -4
3) 1.66x10 mol H2O 4) Chemical
-15
4) 8.3x10 mol C 5) Chemical
5) 18.02 g/mol H2O 6) Chemical
6) 40.0 g/mol NaOH 11) 1.5 mol NaOH 5) 1.35 M
7) Chemical
7) 53.49 g/mol NH4Cl 12) 0.058 mol H2SO4 6) 273 g 8) Physical
8) 98.08 g/mol H2SO4 13) 0.051 mol NH4Cl 7) 5.68 m 9) Chemical
9) 234.0 g/mol Al2(CO3)3 14) 0.042 mol PbO2 8) 0.26 M 10) Chemical
10) 239.2 g/mol PbO2 15) 4.40 g CO2 9) 0.17 M
16) 48.05 g (NH4)2CO3 10) 0.35 m Section 7.2
17) 17.48 g NaOH 11) 0.37 m 1) A,B,C
18) 54.06 g H2O
21 Section 6.4
19) 5.68x10 molecules Na2CO3
25
20) 3.34x102 molecules H2O 1) A, D
21) 0.67 g H2O 2) The salt lowers the freezing point
22) 169.85 g NaCl of water, making it so it must be
23) 5.71 g NaOH colder before the water will freeze
into ice.
Section 6.1 3) The salt raises the boiling
1) Solutions have extremely small particles that allow light to go through. temperature of the water, cooking
Colloids have larger particles which scatter light. Suspensions separate into the spaghetti at a higher
layers upon standing. temperature.
2) Solutions will allow light to go through and colloids will not. 4) Ionic compounds split into separate
3) C ions when they dissolve, but
4) A covalent compounds stay as whole
formulas.
Section 6.2 5) Ionic, 2
1) Polar (or ionic) compounds will dissolve in polar compounds. Nonpolar 6) Covalent, 1
compounds will dissolve in nonpolar compounds. 7) Covalent, 1
2) LiCl is ionic and CCl4 is nonpolar. Ionic compounds do not dissolve in 8) Ionic, 3
nonpolar compounds, because the ionic compound has charged particles 9) Ionic, 4
which are not attracted to the nonpolar solvent. 10) Ionic, 2
3) A 11) 0.2 m CaCl2
4) The oil is nonpolar and does not mix with the polar water. 12) 0.1 m KI
13) 0.2m NaCl
Section 6.3 14) 0.076°C
1) Ppm and ppb are convenient for very, very small concentrations. 15) -1.86°C
2) 0.60 M 16) -3.01°C
3) 0.22 m
4) 0.18 M Section 7.1
1) Physical
2) Chemical
241 3) Physical
particles. The greater the frequency of effective collisions, the faster the 6) F2 + NaOH → NaF + O2 + 7) H2S + 2KOH→2HOH +K2S
reaction. H2O 8) XeF6 + 3 H2O → XeO3 + 6 HF
10) Slower 7) Fe + CuNO3 → Fe(NO3)2 + Cu 9) Cu + 2 AgNO3 → 2 Ag + Cu(NO3)2
2) A
11) Faster 10) 4 Fe + 3 O2 → 2 Fe2O3
3) A 12) Slower Section 7.4 11) 2 Al(OH)3 + Mg3(PO4)2 → 2 AlPO4
4) B 1) 2 Cu + O2 → 2 CuO + 3 Mg(OH)2
5) C Section 7.3 2) 2H2O→2H2 + O2 12) 2 Al + 3 H2SO4 → 3 H2 +
6) Crushed ice 1) Ca +H2O → Ca(OH)2 + H2 3) 2 Fe + 3 H2O → 3 H2 + Fe2O3 Al2(SO4)3
7) Sugar crystals 2) NaOH + Cl2 → NaCl + NaClO + H2O 4) 2 NaCl → 2 Na + Cl2 13) H3PO4 + 3 NH4OH → 3 HOH +
8) Wood shavings 3) Fe + S → FeS 5) 2 AsCl3 + 3 H2S → As2S3 + 6 (NH4)3PO4
9) If the surface area is higher, there are 4) Al + H2SO4 → H2 + Al2(SO4)3 HCl 14) C3H8 +5O2 → 3CO2 +4H2O
more collisions between reacting 5) Al + Fe2O3 → Al2O3 + Fe 6) CaCO3 → CaO + CO2 15) 4 Al + 3 O2 → 2 Al2O3
16) CH4 +2O2→CO2 +2H2O
17) 5
18) 9 11) 66.5 g Cs 13) Reactants are favored over products;
19) Changing the subscripts changes 12) 297 L O2 there is a greater concentration of
22 reactants than products at equilibrium.
which compounds are involved in the 13) 4.3x10 molecules CO
chemical reaction, while changing the 14) 12.7 g O2 14) 0.64
coefficients only changes how many 15) 180.2 g H2O 15) 2.89 M
of a specific substances are involved in 16) ; K=1.34
the reaction. Section 7.7
-4
1) PCl5 is reacting to form PCl3 and Cl2 17) ; K=1.1x10
Section 7.5 at the same time and at the same speed
1) Synthesis that PCl3 and Cl2 are recombining to 18) ; b)more products are
2) Decomposition form PCl5. present at equilibrium because K>1;
3) Single replacement 2) Between 2.0 and 2.5 minutes the c)[SO3]=0.103 M
4) Double replacement reaction reaches equilibrium, because
5) Single replacement the concentration is no longer Section 7.9
6) Decomposition changing after this time. 1) The equilibrium shifts toward the
7) Double replacement 3) The rate of the forward reaction is products.
8) Decomposition equal to the rate of the reverse 2) The equilibrium shifts toward the
9) Single replacement reaction. reactants.
10) Combustion 4) F 3) More products are formed
11) Synthesis combines two or more 5) T 4) More reactants are formed
substances into one product, whereas 6) T 5) More reactants are formed
decomposition splits one reactant into 7) F 6) More products are formed
more than one product. 8) T 7) More reactants are formed
12) CO2 + H2O 8) More products are formed
13) Single replacement; 2 Fe + 3 H2O → Section 7.8 9) More products are formed
3 H2 + Fe2O3 1) Solids and liquids 10) More reactants are formed
14) Double replacement; H2O (or HOH) 2) 11) More reactants are formed
+ (NH4)3PO4 12) More products are formed
15) Combustion; CO2 + H2O 3) 13) More products are formed
16) Synthesis; Al2O3 14) More reactants are formed
4)
17) Double replacement; BaSO4 + NaCl 15) More reactants are formed
18) Single replacement; CaCl2 + H2 5) 16) More products are formed
19) Double replacement; FeCl2 + H2S 17) More reactants are formed
20) Single Replacement; NaBr + I2 6) 18) More products are formed
19) More products are formed
Section 7.6 7)
20) More products are formed
1) 3 mol H2O 8) 21) More reactants are formed
2) 0.31 mol Bi2O3 22) More reactants are formed
3) 15 mol LiCl 9)
4) 1.05 mol SiO2 Section 8.1
5) 0.15 mol Ca3(PO4)2 10)
1) Acids
6) 0.36 mol Fe2O3 2) Bases
11)
7) 3.28 mol FeS 3) Acids
12) Products are favored over reactants;
8) 2.79 g HNO3 there is a greater concentration of 4) Bases
9) 2.12 g I2 products than reactants at equilibrium. 5) Acids
10) 0.78 g LiOH 6) Bases

242 243
+
7) Acids react to form H ions in water
Section 8.2
-2
1) 2.5x10 M; 12.4; basic
-8
2) 1.7x10 M; 6.22; slightly acidic 3) Endothermic and the total number of particles in the 2) The process by which two small nuclei
-4
3) 3.16x10 M; 3.5; acidic 4) Exothermic nucleus decreases by 4 combine to make one larger nuclei
-7
4) 3.16x10 M; 6.5; slightly acidic 5) Exothermic 7) Gamma 3) Nuclear changes involve much more
5) 6.6
6) Exothermic 8) Alpha energy then chemical changes
6) 8.2
7) The solution with the lower pH has 7) Endothermic 9) (frequently about 1 million times
100,000 times greater concentration 8) Exothermic more energy per atom)
10)
+ 9) Endothermic 4) The manner in which the heat is
of H ions. 11)
10) Exothermic; the temperature rises produced that heats the water to
Section 8.3 initially meaning that heat was given 12) turn the turbine
11) HNO3 + KOH → H2O + KNO3 off to the surroundings. 13) 5) Control the speed at which the
14) fission reaction occurs by absorbing
12) HClO4 + NH4OH → H2O + Section 9.2
NH4ClO4 many of the free neutrons which start
15)
1) Oxidized: Cu; Reduced: H fission reactions
13) H2SO4 + 2 NaOH → 2 H2O + Na2SO4 16)
2) Oxidized: H; Reduced: O 6) No, the power plants in the US contain
14) HNO3 + NH4OH → H2O + NH4NO3
3) Oxidized: Al; Reduced: H less than the critical mass of the
15) HF + NH4OH → H2O + NH4F Section 10.3
4) Oxidized: Zn; Reduced: H fissionable isotopes so are unable to
16) HC2H3O2 + KOH → H2O + KC2H3O2 1) 0.25g
The energy that comes from the sun and 5) Oxidized: Al; Reduced: Cu cause a nuclear explosion
17) HCl + KOH → H2O + KCl 2) 1.0 g
other stars is produced by fusion. (Source: 6) Energy of moving electrons 7) Fission
18) Mg(OH)2 + 2 HCl → 2 H2O + MgCl2 http://commons.wikimedia.org/wiki/File:Su 3) 8.0 years 8) Nuclear decay
19) 2 HCl + Ba(OH)2 → 2 H2O + BaCl2 n-in-X-ray. Public Domain) 4) 17,100 years 9) Fusion
20) NaOH + HClO4 → H2O + NaClO4 7) The anode is where electrons are lost 5) 10 years 10) Nuclear decay
(where oxidation occurs)
Section 8.4 11) Fission and fusion
8) The cathode is where electrons are Section 10.4
1) Indicators are weak acids that gained (where reduction occurs) 1) The process by which a large nucleus
change color when they react with a 9) Zinc is split into two or more smaller nuclei
base. They are used in a titration to 10) The anode is the zinc electrode,
show when all of the acid or base because the zinc is being oxidized
has reacted. 11) The anode is the copper electrode,
2) When the number of moles of acid is because copper ions are being reduced
equal to the number of moles of base 12) Electrons will flow from the anode
3) 0.1176 M NaOH (zinc electrode) to the cathode (copper
4) 0.1708 M HClO electrode)
5) 33.2 mL HCl
6) 62.5 mL HI Section 10.2
7) 1.4 M NaOH 1) C
8) 0.30 M HC2H3O2 2) B
3) A
Section 9.1 4) Alpha particles will move toward the
1) Endothermic reactions absorb (take negative plate; beta particles will
in) energy; exothermic reactions move toward the positive plate;
release energy. gamma particles will continue in a
2) Exothermic straight path.
5) The number of protons increases by 1,
the number of neutrons decreases by 1,
and the total number of particles in t he
nucleus does not change.
6) The number of protons decreases by 2,
the number of neutrons decreases by 2,

244 245
Glossary
A substances
Absolute Zero: the temperature at which molecules stop moving and therefore, have zero kinetic Coefficient: a small whole number that appears in front of a formula in a balanced chemical
energy equation
Alkali earth metals: group 2A of the periodic table Colligative property: a property that is due only to the number of particles in solution and not the type
Alkali metals: group 1A of the periodic table of the solute
Alpha decay: Alpha decay is a common mode of radioactive decay in which a nucleus emits an Colloid: type of mixture in which the size of the particles is between 1x10 3 pm and 1x108
alpha particle (a helium-4 nucleus). pm
Alpha particle: An alpha particle is a helium-4 nucleus, composed of 2 protons and 2 neutrons Combustion reaction: a reaction in which oxygen reacts with another substance to produce carbon
dioxide and water.
Anion: negative ion; formed by gaining electrons
Compound: a substance that is made up of more than one type of atom bonded together
Anode: The electrode at which oxidation occurs.
Concentrated: a solution in which there is a large amount of solute in a given amount of solvent
Arrhenius acid: a substance that produces H+ ions in solution
Concentration: the measure of how much of a given substance is mixed with another substance
Arrhenius base: a substance that produces OH- ions in a solution
Control rods : made of chemical elements capable of absorbing many neutrons and are used to
Atom: Democritus’ word for the tiny, indivisible, solid objects that he believed made up
control the rate of a fission chain reaction in a nuclear reactor.
all matter in the universe
Controlled experiment: An experiment that compares the results of an experimental sample to a control
Atomic mass unit a unit of mass equal to one-twelfth the mass of a carbon-twelve atom
sample
(amu):
Atomic mass: the weighted average of the masses of the isotopes of an element Conversion factor: a ratio used to convert one unit of measurement into another.
Cosmic background energy in the form of radiation leftover from the early big bang
Atomic number: the number of protons in the nucleus of an atom
radiation:
Avogadro's number: The number of objects in a mole; equal to 6.02x1023. Covalent bond: A type of bond in which electrons are shared by atoms.
B Covalent compound: two or more atoms (typically nonmetals) forming a molecule in which electrons
Background radiation: Radiation that comes from environment sources including the earth's crust, the are being shared between atoms.
atmosphere, cosmic rays, and radioisotopes. These natural sources of radiation Critical mass: The smallest mass of a fissionable material that will sustain a nuclear chain
account for the largest amount of radiation received by most people. reaction at a constant level.
Balanced chemical a chemical equation in which the number of each type of atom is equal on the two D
equation: sides of the equation Dalton’s Atomic the first scientific theory to relate chemical changes to the structure, properties,
Battery: A group of two or more cells that produces an electric current. Theory: and behavior of the atom
Beta decay: Beta decay is a common mode of radioactive decay in which a nucleus emits beta Decomposition a reaction in which one reactant breaks down to form two or more products
particles. The daughter nucleus will have a higher atomic number than the original reaction:
nucleus. Dilute: a solution in which there is a small amount of solute in a given amount of solvent
Beta particle: A beta particle is a high speed electron, specifically an electron of nuclear origin. Double replacement a reaction in which two reactants form products by having the cations exchange
Big Bang Theory: the idea that the universe was originally extremely hot and dense at some reaction: places with the anions
finite time in the past and has since cooled by expanding to the present state and Ductile: can be drawn out into thin wires
continues to expand today
Boiling point elevation: the amount the boiling point of a solution increases from the boiling point of a
E
Electrochemical cell: An arrangement of electrodes and ionic solutions in which a redox reaction is used
pure solvent
to make electricity; a battery
C Electrolysis: A chemical reaction brought about by an electric current.
Catalyst: A substance that increases the rate of a chemical reaction but is, itself, left
Electron configuration: a list that represents the arrangement of electrons of an atom.
unchanged, at the end of the reaction; lowers activation energy
Cathode: electrode at which reduction occurs. Electron: a negatively charged subatomic particle, responsible for chemical bonding
Cation: positive ion; formed by losing electrons Electronegativity: The tendency of an atom in a molecule to attract shared electrons to itself.
Chain reaction: A multi-stage nuclear reaction that sustains itself in a series of fissions in which Electronegativity: the ability of an atom in a molecule to attract shared electrons
the release of neutrons from the splitting of one atom leads to the splitting of Electroplating: A process in which electrolysis is used as a means of coating an object with a
others. layer of metal.
Chemical changes: changes that occur when one substance is turned into another substance; different electrostatic attraction: The force of attraction between opposite electric charges.
types of molecules are present at the beginning and end of the change. electrostatic attraction: the attraction of oppositely charged particles
Chemical reaction: the process in which one or more substances are changed into one or more new Element: a substance that is made up of only one type of atom Subatomic particles: particles
that are smaller than the atom
Endothermic: reactions in which energy is absorbed, heat can be considered as a reactant

246 247

Equilibrium constant A mathematical ratio that shows the concentrations of the products divided by
(K): concentration of the reactants.
Equilibrium: A state that occurs when the rate of forward reaction is equal to the rate of the
reverse reaction.
Equivalence point: the point in the titration where the number of moles of acid equals the number of L
moles of base Le Châtelier’s Applying a stress to an equilibrium system causes the equilibrium position to shift
Exothermic reaction: A reaction in which heat is released, or is a product of a reaction. Principle: to offset that stress and regain equilibrium.
Experiment: A controlled method of testing a hypothesis. M
Extrapolation: the process of creating data points beyond the end of the graph line, using the
malleable: can be hammered into thin sheets
basic shape of the curve as a guide
Mass number: the total number of protons and neutrons in the nucleus of an atom
F
Fission: A nuclear reaction in which a heavy nucleus splits into two or more smaller Mass: a measure of the amount of matter in an object
fragments, releasing large amounts of energy. Mendeleev: the Russian chemist credited with organizing the periodic table in the form we use
Formula unit: the empirical formula of an ionic compound; shows the lowest possible ratio today.
Freezing point the amount the freezing point of a solution decreases from the freezing point of a Metric system: international decimal-based system of measurement.
depression: pure solvent Miscible: liquids that have the ability to dissolve in each other
Fusion: A nuclear reaction in which nuclei combine to form more massive nuclei with the Mixture: a combination of two or more elements or compounds which have not reacted to
simultaneous release of energy. bond together; each part in the mixture retains its own properties
G Molality:
Gamma ray: Gamma radiation is the highest energy on the spectrum of electromagnetic Molar Mass: The mass, in grams, of 1 mole of a substance. This can be found by adding up the
radiation. masses on the periodic table.
Graph: a pictorial representation of patterns using a coordinate system Molarity: the number of moles of solute per liter of solution
Group (family): a vertical column in the periodic table, have similar chemcial properties Mole ratio: the ratio of the moles of one reactant or product to the moles of another reactant
or product according to the coefficients in the balanced chemical equation
H
Mole: An Avogadro’s number of objects; 6.02 x 1023 particles
Half-life: the time interval required for a quantity of material to decay to half its original
value. Molecular geometry: The specific three-dimensional arrangement of atoms in molecules.
Halogens: group 7A of the periodic table; reactive non-metals N
Hydrocarbon: an organic substance consisting of only hydrogen and carbon Neutralization: a reaction between an acid and a base that produces water and a salt
Hypothesis: A tentative explanation that can be tested by further investigation. Neutron: a subatomic particle with no charge
I Noble gases: group 8A of the periodic table; extremely non-reactive
Immiscible: liquids that do not have the ability to dissolve in each other Nuclear charge: the number of protons in the nucleus
Indicator: a substance that changes color at a specific pH and is used to indicate the pH of Nucleus: the small, dense center of the atom
the solution
International System of the internationally agreed upon standard metric system O
Units (Le Système Octet rule: the tendency for atoms gain or lose the appropriate number of electrons so that
International d’ Unites): the resulting ion has either completely filled or completely empty outer energy
Interpolation: the process of estimating values between measured values levels, or 8 valance electrons.
Ion: An atom or group of atoms with an excess positive or negative charge, lost or Oxidation: a loss of electrons, resulting in an increased charge or oxidation number
gained electrons P
ionic bond: A bond between ions resulting from the transfer of electrons from one of the periodic law: states that the properties of the elements recur periodically as their atomic
bonding atoms to the other and the resulting electrostatic attraction between the numbers increase
ions. Periodic table: a tabular arrangement of the chemical elements according to atomic number.
Ionic compound: a positively charged particle (typically a metal) bonded to a negatively charged
Physical changes: changes that do not alter the identity of a substance, the same types of molecules
particle (typically a nonmetal) held together by electrostatic attraction
are present at the beginning and end of the change.
Ionic Formula: includes the symbols and number of each ion (atom) present in a compound in the
Polar covalent bond: A covalent bond in which the electrons are not shared equally because one atom
lowest whole number ratio
attracts them more strongly that the other.
Ionization energy: the energy required to remove the most loosely held electron from a gaseous atom
or ion Potential energy: The energy of position or stored energy, including bond energy.
Isotopes: atoms of the same element that have the same number of protons but different Products: materials present at the end of a reaction, shown on the right of the arrow in a
numbers of neutrons, same atomic number but different mass number chemical equation
Proton: a positively charged subatomic particle
Q
Quark: particles that form one of the two basic constituents of matter. Various species of

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 quarks combine in specific ways to form protons and neutrons, in each case taking
exactly three quarks to make the composite particle.
R
Reactants: the starting materials in a reaction, shown left of the arrow in a chemical equation
Reduction: gaining electrons, resulting in a decreased charge or oxidation number
S
Scientific notation: a shorthand method of writing very large and very small numbers in terms of a
decimal number between 1 and 10 multiplied by 10 to a power.
Shielding effect: the inner electrons help “shield” the outer electrons and the nucleus from each
other.
Single replacement a reaction in which an element reacts with a compound to form products
reaction:
Slope: the ratio of the change in one variable with respect to the other variable.
Solute: the substance in a solution present in the least amount, dissolved by the solute
Solution: a homogeneous mixture of substances
Solvent: the substance in a solution present in the greatest amount
Stoichiometry: the calculation of quantitative relationships of the reactants and products in a
balanced chemical equation
Subscripts: part of the chemical formulas of the reactants and products that indicate the
number of atoms of the preceding element
Surface area to volume The comparison of the volume inside a solid to the area exposed on the surface.
ratio:
Suspension: type of mixture in which the particles settle to the bottom of the container and can
be separated by filtration
Synthesis reaction: a reaction in which two or more reactants combine to make one product
T
Temperature: the average kinetic energy of the particles that make up a material
Theory: A well-established explanation based on extensive experimental data
Titration: the lab process in which a known concentration of base (or acid) is added to a
solution of acid (or base) of unknown concentration
Transition elements: groups 3 to 12 of the periodic table
V
Valence electrons: the electrons in the outermost energy level of an atom.
VSEPR model: A model whose main postulate is that the structure around a given atom in a
molecule is determined by minimizing electron-pair repulsion.
W
Weight: the force of attraction between the object and the earth (or whatever large body it
is resting on)

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