Chapter 1 (from textbook)

CHE1401, Chapter 1
CHEMISTRY
T h e C e n T r a l S C i e n C e 13 Th ediTion
Theodore L. Brown
University of Illinois at Urbana-Champaign
H. Eugene LeMay, Jr.

University of Nevada, Reno
Bruce E. Bursten
University of Tennessee, Knoxville
Catherine J. Murphy
University of Illinois at Urbana-Champaign
Patrick M. Woodward
The Ohio State University
Matthew W. Stoltzfus
The Ohio State University
Boston Columbus indianapolis new York San Francisco Upper Saddle river
amsterdam Cape Town dubai london Madrid Milan Munich Paris Montréal Toronto
delhi Mexico City São Paulo Sydney hong Kong Seoul Singapore Taipei Tokyo
CONTENTS
Preface xx
2 Atoms, Molecules,
1 Introduction: Matter and Ions 40
and Measurement 2 2.1 The Atomic Theory of Matter 42
2.2 The discovery of Atomic Structure 43
1.1 The Study of Chemistry 2
Cathode Rays and Electrons 43
The Atomic and Molecular Perspective of Radioactivity 45 The Nuclear Model of the
Chemistry 4 Why Study Chemistry? 5 Atom 46
1.2 Classifications of Matter 6 2.3 The Modern View of Atomic Structure 47
States of Matter 7 Pure Substances 7 Atomic Numbers, Mass Numbers, and
Elements 7 Compounds 8 Mixtures 10 Isotopes 49
1.3 Properties of Matter 11 2.4 Atomic Weights 50
Physical and Chemical Changes 12 The Atomic Mass Scale 50 Atomic Weight 51
Separation of Mixtures 13
2.5 The Periodic Table 52
1.4 Units of Measurement 14
2.6 Molecules and Molecular
SI Units 15 Length and Mass 17
Compounds 56
Temperature 17 derived SI Units 19
Volume 19 density 19 Molecules and Chemical Formulas 56
Molecular and Empirical Formulas 56
1.5 Uncertainty in Measurement 22 Picturing Molecules 57
Precision and Accuracy 22 Significant
2.7 Ions and Ionic Compounds 58
Figures 22 Significant Figures in
Calculations 22 Predicting Ionic Charges 59 Ionic
Compounds 60
1.6 dimensional Analysis 27
2.8 Naming Inorganic Compounds 62
Using Two or More Conversion Factors 28
Conversions Involving Volume 29 Names and Formulas of Ionic Compounds 62
Names and Formulas of Acids 67 Names and
Chapter Summary and Key Terms 32
Formulas of Binary Molecular Compounds 68
Learning Outcomes 32
Key Equations 32 Exercises 32 Additional 2.9 Some Simple Organic Compounds 69
Exercises 37 Alkanes 69 Some derivatives of Alkanes 70
Chemistry Put to Work Chemistry and the Chapter Summary and Key Terms 72
Chemical Industry 6 Learning Outcomes 72 Key
Equations 73 Exercises 73
A Closer Look The Scientific Method 14
Additional Exercises 78
Chemistry Put to Work Chemistry in
the News 20 A Closer Look Basic Forces 49
Strategies in Chemistry Estimating Answers 28 A Closer Look The Mass Spectrometer 52
Strategies in Chemistry The Importance of A Closer Look What Are Coins Made Of? 54
Practice 31 Chemistry and Life Elements Required by Living
Strategies in Chemistry The Features of This Organisms 61
Book 32 Strategies in Chemistry How to Take a Test 71
vii
1
Introduction: Matter
and Measurement
In the title of this book we refer to chemistry as the central science. This
title reflects the fact that much of what goes on in the world around us
involves chemistry. The changes that produce the brilliant colors of tree leaves
in the fall, the electrical energy that powers a cell phone, the spoilage of foods
left standing at room temperature, and the many ways in which our bodies use
the foods we consume are all everyday examples of chemical processes.
Chemistry is the study of matter and the changes that matter undergoes. As you progress
in your study, you will come to see how chemical principles operate in all aspects of our
lives, from everyday activities like food preparation to more complex processes such as
those that operate in the environment. We use chemical principles to understand a host of
phenomena, from the role of salt in our diet to the workings of a lithium ion battery.
This first chapter provides an overview of what chemistry is about and what chem-
ists do. The “What’s Ahead” list gives an overview of the chapter organization and of
some of the ideas we will consider.
▶ THE BEAUTIFUL COLORS that develop
1.1 | The Study of Chemistry

in trees in the fall appear when the tree
ceases to produce chlorophyll, which
imparts the green color to the leaves during
Chemistry is at the heart of many changes we see in the world around us, and it ac- the summer. Some of the color we see has
counts for the myriad of different properties we see in matter. To understand how these been in the leaf all summer, and some
changes and properties arise, we need to look far beneath the surfaces of our everyday develops from the action of sunlight on the
observations. leaf as the chlorophyll disappears.
WHAT’S
AHEAD 1.3 ProPertIes of MAtter We then consider different
characteristics, or properties, used to characterize, identify, and
separate substances, distinguishing between chemical and
1.1 the stUdy of CheMIstry We begin with a brief physical properties.
description of what chemistry is, what chemists do, and why it is
useful to learn chemistry. 1.4 UnIts of MeAsUreMent We observe that many
properties rely on quantitative measurements involving
1.2 ClAssIfICAtIons of MAtter next, we examine some numbers and units. the units of measurement used throughout
fundamental ways to classify matter, distinguishing between science are those of the metric system.
pure substances and mixtures and between elements and
compounds.
1.5 UnCertAInty In MeAsUreMent We observe that the 1.6 dIMensIonAl AnAlysIs We recognize that units as well
uncertainty inherent in all measured quantities is expressed by as numbers are carried through calculations and that obtaining
the number of significant figures used to report the quantity. correct units for the result of a calculation is an important way to
significant figures are also used to express the uncertainty check whether the calculation is correct.
associated with calculations involving measured quantities.
4 ChaPter 1 Introduction: Matter and Measurement
The Atomic and Molecular Perspective of Chemistry

Chemistry is the study of the properties and behavior of matter. Matter is the physical
material of the universe; it is anything that has mass and occupies space. A property is
any characteristic that allows us to recognize a particular type of matter and to distinguish
it from other types. This book, your body, the air you are breathing, and the clothes you
are wearing are all samples of matter. We observe a tremendous variety of matter in our
world, but countless experiments have shown that all matter is comprised of combina-
tions of only about 100 substances called elements. One of our major goals will be to relate
the properties of matter to its composition, that is, to the particular elements it contains.
Chemistry also provides a background for understanding the properties of matter
in terms of atoms, the almost infinitesimally small building blocks of matter. Each ele-
ment is composed of a unique kind of atom. We will see that the properties of matter re-
late to both the kinds of atoms the matter contains (composition) and the arrangements
of these atoms (structure).
In molecules, two or more atoms are joined in specific shapes. Throughout this text
you will see molecules represented using colored spheres to show how the atoms are con-
nected (▼ Figure 1.1). The color provides a convenient way to distinguish between atoms
of different elements. For example, notice that the molecules of ethanol and ethylene gly-
col in Figure 1.1 have different compositions and structures. Ethanol contains one oxygen
atom, depicted by one red sphere. In contrast, ethylene glycol contains two oxygen atoms.
Even apparently minor differences in the composition or structure of molecules
can cause profound differences in properties. For example, let’s compare ethanol and
ethylene glycol, which appear in Figure 1.1 to be quite similar. Ethanol is the alcohol in
beverages such as beer and wine, whereas ethylene glycol is a viscous liquid used as au-
tomobile antifreeze. The properties of these two substances differ in many ways, as do
their biological activities. Ethanol is consumed throughout the world, but you should
never consume ethylene glycol because it is highly toxic. One of the challenges chemists
undertake is to alter the composition or structure of molecules in a controlled way, cre-
ating new substances with different properties. For example, the common drug aspirin,
shown in Figure 1.1, was first synthesized in 1897 in a successful attempt to improve on
a natural product extracted from willow bark that had long been used to alleviate pain.
Every change in the observable world—from boiling water to the changes that occur
as our bodies combat invading viruses—has its basis in the world of atoms and molecules.
GO FIGURE
Which of the molecules in the figure has the most carbon atoms? how many are there in that molecule?
=H =O =C
Oxygen
Water Ethanol
Carbon dioxide Ethylene glycol Aspirin

▲ Figure 1.1 Molecular models. The white, black, and red spheres represent atoms of
hydrogen, carbon, and oxygen, respectively.
seCtIon 1.1 the study of Chemistry 5
realms: the macroscopic realm of ordinary-sized objects 1macro = large2 and the submi-
Thus, as we proceed with our study of chemistry, we will find ourselves thinking in two
croscopic realm of atoms and molecules. We make our observations in the macroscopic
world, but to understand that world, we must visualize how atoms and molecules behave at
the submicroscopic level. Chemistry is the science that seeks to understand the properties
and behavior of matter by studying the properties and behavior of atoms and molecules.
Give It Some Thought

(a) approximately how many elements are there?
(b) What submicroscopic particles are the building blocks of matter?
Why Study Chemistry?

Chemistry lies near the heart of many matters of public concern, such as improvement
of health care, conservation of natural resources, protection of the environment, and the
supply of energy needed to keep society running. Using chemistry, we have discovered
and continually improved upon pharmaceuticals, fertilizers and pesticides, plastics, solar
panels, LEDs, and building materials. We have also discovered that some chemicals are
potentially harmful to our health or the environment. This means that we must be sure
that the materials with which we come into contact are safe. As a citizen and consumer,
it is in your best interest to understand the effects, both positive and negative, that chem-
icals can have, and to arrive at a balanced outlook regarding their uses.
You may be studying chemistry because it is an essential part of your curriculum.
Your major might be chemistry, or it could be biology, engineering, pharmacy, agricul-
ture, geology, or some other field. Chemistry is central to a fundamental understand-
ing of governing principles in many science-related fields. For example, our interactions
with the material world raise basic questions about the materials around us. ▼ Figure 1.2
illustrates how chemistry is central to several different realms of modern life.
Energy Biochemistry
Solar panels are composed The flash of the firefly results
of specially treated silicon. from a chemical reaction in
the insect.
Medicine
Technology Connectors and tubing for
Chemistry medical procedures such as
LED’s (light emitting diodes) intravenous injections are
are formed from elements made from plastics highly
such as gallium, arsenic and resistant to chemical attack.
phosphorus.
▲ Figure 1.2 Chemistry is central to our understanding of the world around us.
Chemistry Put to Work

Chemistry and the Chemical
Industry desired properties; (2) measure the properties of matter; and (3) develop
models that explain and/or predict the properties of matter. One chem-
Chemistry is all around us. Many people are familiar with household ist, for example, may work in the laboratory to discover new drugs. An-
chemicals, particularly kitchen chemicals such as those shown in other may concentrate on the development of new instrumentation to
▶ Figure 1.3. However, few realize the size and importance of the measure properties of matter at the atomic level. Other chemists may
chemical industry. Worldwide sales of chemicals and related prod- use existing materials and methods to understand how pollutants are
ucts manufactured in the United States total approximately $585 bil- transported in the environment or how drugs are processed in the
lion annually. Sales of pharmaceuticals total another $180 billion. The body. Yet another chemist will develop theory, write computer code,
chemical industry employs more than 10% of all scientists and engi- and run computer simulations to understand how molecules move and
neers and is a major contributor to the U.S. economy. react. The collective chemical enterprise is a rich mix of all of these
Vast amounts of industrial chemicals are produced each year. activities.
▼ Table 1.1 lists several of the chemicals produced in highest vol-
umes in the United States. Notice that they all serve as raw materi-
als for a variety of uses, including the manufacture and processing
of metals, plastics, fertilizers, and other goods.
Who are chemists, and what do they do? People who have
degrees in chemistry hold a variety of positions in industry, govern-
ment, and academia. Those in industry work as laboratory chem-
ists, developing new products (research and development); analyzing
materials (quality control); or assisting customers in using products
(sales and service). Those with more experience or training may
work as managers or company directors. Chemists are important
members of the scientific workforce in government (the National
Institutes of Health, Department of Energy, and Environmental
Protection Agency all employ chemists) and at universities. A chem-
istry degree is also good preparation for careers in teaching, medi-
cine, biomedical research, information science, environmental work,
technical sales, government regulatory agencies, and patent law.
Fundamentally, chemists do three things: (1) make new types
of matter: materials, substances, or combinations of substances with ▲ Figure 1.3 Common chemicals employed in home food production.
Table 1.1 Several of the Top Chemicals Produced by the U.S. Chemical Industry*
Annual Production
Chemical Formula (Billions of Pounds) Principal End Uses
Sulfuric acid H2SO4 70 Fertilizers, chemical manufacturing
Ethylene C2H4 50 Plastics, antifreeze
Lime CaO 45 Paper, cement, steel
Propylene C3H6 35 Plastics
Ammonia NH3 18 Fertilizers
Chlorine Cl2 21 Bleaches, plastics, water purification
Phosphoric acid H3PO4 20 Fertilizers
Sodium hydroxide NaOH 16 Aluminum production, soap
1.2 | Classifications of Matter

Let’s begin our study of chemistry by examining two fundamental ways in which mat-
ter is classified. Matter is typically characterized by (1) its physical state (gas, liquid,
or solid) and (2) its composition (whether it is an element, a compound, or a mixture).
*Data from Chemical & Engineering News, July 2, 2007, pp. 57, 60, American Chemical Society; data online
from U.S. Geological Survey.
seCtIon 1.2 Classifications of Matter 7
States of Matter
A sample of matter can be a gas, a liquid, or a solid. These
three forms, called the states of matter, differ in some of GO FIGURE
their observable properties. A gas (also known as vapor)
in which form of water are the water molecules farthest apart?
has no fixed volume or shape; rather, it uniformly fills its
container. A gas can be compressed to occupy a smaller Water vapor Ice
volume, or it can expand to occupy a larger one. A liq-
uid has a distinct volume independent of its container,
and assumes the shape of the portion of the container it Liquid water
occupies. A solid has both a definite shape and a definite
volume. Neither liquids nor solids can be compressed to
any appreciable extent.
The properties of the states of matter can be under-
stood on the molecular level (▶ Figure 1.4). In a gas the
molecules are far apart and moving at high speeds, col-
liding repeatedly with one another and with the walls of
the container. Compressing a gas decreases the amount
of space between molecules and increases the frequency
of collisions between molecules but does not alter the
size or shape of the molecules. In a liquid, the molecules
are packed closely together but still move rapidly. The
rapid movement allows the molecules to slide over one an-
other; thus, a liquid pours easily. In a solid the molecules
are held tightly together, usually in definite arrangements
in which the molecules can wiggle only slightly in their
otherwise fixed positions. Thus, the distances between
molecules are similar in the liquid and solid states, but
the two states differ in how free the molecules are to move
around. Changes in temperature and/or pressure can lead ▲ Figure 1.4 The three physical states of water—water vapor, liquid
water, and ice. We see the liquid and solid states but cannot see the
to conversion from one state of matter to another, illus- gas (vapor) state. The red arrows show that the three states of matter
trated by such familiar processes as ice melting or water interconvert.
vapor condensing.
Pure Substances
Most forms of matter we encounter—the air we breathe (a gas), the gasoline we burn in
our cars (a liquid), and the sidewalk we walk on (a solid)—are not chemically pure. We
can, however, separate these forms of matter into pure substances. A pure substance
(usually referred to simply as a substance) is matter that has distinct properties and a
composition that does not vary from sample to sample. Water and table salt (sodium
chloride) are examples of pure substances.
All substances are either elements or compounds. Elements are substances that cannot
be decomposed into simpler substances. On the molecular level, each element is composed
of only one kind of atom [Figure 1.5(a and b)]. Compounds are substances composed
of two or more elements; they contain two or more kinds of atoms [Figure 1.5(c)].
Water, for example, is a compound composed of two elements: hydrogen and oxygen.
Figure 1.5(d) shows a mixture of substances. Mixtures are combinations of two or more
substances in which each substance retains its chemical identity.
Elements
Currently, 118 elements are known, though they vary widely in abundance. Hydrogen
constitutes about 74% of the mass in the Milky Way galaxy, and helium constitutes
24%. Closer to home, only five elements—oxygen, silicon, aluminum, iron, and
calcium—account for over 90% of Earth’s crust (including oceans and atmosphere),
and only three—oxygen, carbon, and hydrogen—account for over 90% of the mass of
the human body (Figure 1.6).
GO FIGURE
how do the molecules of a compound differ from the molecules of an element?
(a) Atoms of an element (b) Molecules (c) Molecules (d) Mixture of elements
of an element of a compound and a compound
Only one kind of atom is in any element. Compounds must have at

least two kinds of atoms.
▲ Figure 1.5 Molecular comparison of elements, compounds, and mixtures.
▼ Table 1.2 lists some common elements, along with the chemical symbols
used to denote them. The symbol for each element consists of one or two letters,
GO FIGURE
with the first letter capitalized. These symbols are derived mostly from the Eng-
name two significant differences between
lish names of the elements, but sometimes they are derived from a foreign name
the elemental composition of earth’s crust
instead (last column in Table 1.2). You will need to know these symbols and learn
and the elemental composition of the
others as we encounter them in the text.
human body.
All of the known elements and their symbols are listed on the front
Aluminum inside cover of this text in a table known as the periodic table. In the periodic
Iron 7.5% Other
4.7% 9.2% table the elements are arranged in columns so that closely related elements
Calcium are grouped together. We describe the periodic table in more detail in
3.4% Section 2.5 and consider the periodically repeating properties of the elements
Silicon in Chapter 7.
Oxygen 25.7%
49.5% Compounds
Most elements can interact with other elements to form compounds. For
example, when hydrogen gas burns in oxygen gas, the elements hydrogen and
oxygen combine to form the compound water. Conversely, water can be decom-
Earth’s crust
posed into its elements by passing an electrical current through it (▶ Figure 1.7).
Other
7%
Hydrogen
Table 1.2 Some Common Elements and Their Symbols
10%
Carbon C Aluminum Al Copper Cu (from cuprum)
Oxygen Carbon
65% Fluorine F Bromine Br Iron Fe (from ferrum)
18%
Hydrogen H Calcium Ca Lead Pb (from plumbum)
Iodine I Chlorine Cl Mercury Hg (from hydrargyrum)
Nitrogen N Helium He Potassium K (from kalium)
Human body
Oxygen O Lithium Li Silver Ag (from argentum)
▲ Figure 1.6 Relative abundances of elements.*
Phosphorus P Magnesium Mg Sodium Na (from natrium)
elements in percent by mass in earth’s crust (including
oceans and atmosphere) and the human body. Sulfur S Silicon Si Tin Sn (from stannum)
*U.S. Geological Survey Circular 285, U.S Department of the Interior.

seCtIon 1.2 Classifications of Matter 9
GO FIGURE
how are the relative gas volumes collected in the two tubes related to the relative number of gas molecules in the tubes?
Oxygen gas, O2
Water, H2O Hydrogen gas, H2

▲ Figure 1.7 Electrolysis of water. Water decomposes into its component elements, hydrogen
and oxygen, when an electrical current is passed through it. The volume of hydrogen, collected
in the right test tube, is twice the volume of oxygen.
Pure water, regardless of its source, consists of 11% hydrogen and 89% oxygen by mass.
This macroscopic composition corresponds to the molecular composition, which
consists of two hydrogen atoms combined with one oxygen atom:
Hydrogen atom Oxygen atom Water molecule

(written H) (written O) (written H2O)
The elements hydrogen and oxygen themselves exist naturally as diatomic (two-
atom) molecules:
Oxygen molecule (written O2)
Hydrogen molecule (written H2)
As seen in ▼ Table 1.3, the properties of water bear no resemblance to the proper-
ties of its component elements. Hydrogen, oxygen, and water are each a unique sub-
stance, a consequence of the uniqueness of their respective molecules.
Table 1.3 Comparison of Water, Hydrogen, and Oxygen

Water Hydrogen Oxygen
Statea Liquid Gas Gas
Normal boiling point 100 °C - 253 °C -183 °C
Densitya 1000 g/L 0.084 g/L 1.33 g/L
Flammable No Yes No
a
At room temperature and atmospheric pressure.
The observation that the elemental composition of a compound is always the same
is known as the law of constant composition (or the law of definite proportions).
French chemist Joseph Louis Proust (1754–1826) first stated the law in about 1800.
Although this law has been known for 200 years, the belief persists among some peo-
ple that a fundamental difference exists between compounds prepared in the labora-
tory and the corresponding compounds found in nature. However, a pure compound
has the same composition and properties under the same conditions regardless of its
source. Both chemists and nature must use the same elements and operate under the
same natural laws. When two materials differ in composition or properties, either they
are composed of different compounds or they differ in purity.

hydrogen, oxygen, and water are all composed of molecules. What is it about a
molecule of water that makes it a compound, whereas hydrogen and oxygen are
elements?
Mixtures
Most of the matter we encounter consists of mixtures of different substances. Each sub-
stance in a mixture retains its chemical identity and properties. In contrast to a pure
substance, which by definition has a fixed composition, the composition of a mixture
can vary. A cup of sweetened coffee, for example, can contain either a little sugar or a
lot. The substances making up a mixture are called components of the mixture.
Some mixtures do not have the same composition, properties, and appearance
throughout. Rocks and wood, for example, vary in texture and appearance in any
typical sample. Such mixtures are heterogeneous [▼ Figure 1.8(a)]. Mixtures that are
uniform throughout are homogeneous. Air is a homogeneous mixture of nitrogen,
oxygen, and smaller amounts of other gases. The nitrogen in air has all the proper-
ties of pure nitrogen because both the pure substance and the mixture contain the
same nitrogen molecules. Salt, sugar, and many other substances dissolve in water to
form homogeneous mixtures [Figure 1.8(b)]. Homogeneous mixtures are also called
solutions. Although the term solution conjures an image of a liquid, solutions can be
solids, liquids, or gases.
▶ Figure 1.9 summarizes the classification of matter into elements, compounds,
and mixtures.
(a) (b)
▲ Figure 1.8 Mixtures. (a) Many common materials, including rocks, are heterogeneous mixtures.
This photograph of granite shows a heterogeneous mixture of silicon dioxide and other metal
oxides. (b) homogeneous mixtures are called solutions. Many substances, including the blue solid
shown here [copper(ii) sulfate], dissolve in water to form solutions.
seCtIon 1.3 Properties of Matter 11
Matter
NO Is it uniform YES
throughout?
Heterogeneous
Homogeneous
mixture
NO Does it have a YES

variable
composition?
Homogeneous
Pure substance
mixture
(solution)
NO Does it contain YES

more than one
kind of atom?
Element Compound
▲ Figure 1.9 Classification of matter. all pure matter is classified ultimately as either an element
or a compound.
SAMPLE
EXERCISE 1.1 Distinguishing among Elements, Compounds, and Mixtures
“White gold” contains gold and a “white” metal, such as palladium. Two samples of white gold
differ in the relative amounts of gold and palladium they contain. Both samples are uniform in
composition throughout. Use Figure 1.9 to classify white gold.
SOLUTION
Because the material is uniform throughout, it is homogeneous. (c) It consists of a heterogeneous mixture of compounds.
Because its composition differs for the two samples, it cannot be a (d) It consists of a heterogeneous mixture of elements and
compound. Instead, it must be a homogeneous mixture. compounds.
(e) It consists of a single compound in different states.
Practice Exercise 1
Which of the following is the correct description of a cube of Practice Exercise 2
material cut from the inside of an apple? Aspirin is composed of 60.0% carbon, 4.5% hydrogen, and 35.5%
(a) It is a pure compound. oxygen by mass, regardless of its source. Use Figure 1.9 to classify
(b) It consists of a homogenous mixture of compounds. aspirin.
1.3 | Properties of Matter

Every substance has unique properties. For example, the properties listed in Table 1.3
allow us to distinguish hydrogen, oxygen, and water from one another. The properties
of matter can be categorized as physical or chemical. Physical properties can be ob-
served without changing the identity and composition of the substance. These proper-
ties include color, odor, density, melting point, boiling point, and hardness. Chemical
properties describe the way a substance may change, or react, to form other substances.
A common chemical property is flammability, the ability of a substance to burn in the
presence of oxygen.
Some properties, such as temperature and melting point, are intensive properties.
Intensive properties do not depend on the amount of sample being examined and are
particularly useful in chemistry because many intensive properties can be used to identify
substances. Extensive properties depend on the amount of sample, with two examples
being mass and volume. Extensive properties relate to the amount of substance present.

When we say that lead is a denser metal than aluminum, are we talking about an
extensive or intensive property?
Physical and Chemical Changes

The changes substances undergo are either physical or chemical. During a physical
change, a substance changes its physical appearance but not its composition. (That is, it
is the same substance before and after the change.) The evaporation of water is a physi-
cal change. When water evaporates, it changes from the liquid state to the gas state, but
it is still composed of water molecules, as depicted in Figure 1.4. All changes of state
(for example, from liquid to gas or from liquid to solid) are physical changes.
In a chemical change (also called a chemical reaction), a substance is transformed
into a chemically different substance. When hydrogen burns in air, for example, it under-
goes a chemical change because it combines with oxygen to form water (▼ Figure 1.10).
H2 O2
Burn
H2 O2 H2O
▲ Figure 1.10 A chemical reaction.
Chemical changes can be dramatic. In the account that follows, Ira Remsen, author
of a popular chemistry text published in 1901, describes his first experiences with
chemical reactions. The chemical reaction that he observed is shown in ▼ Figure 1.11.
▲ Figure 1.11 The chemical reaction between a copper penny and nitric acid. The dissolved copper produces the blue-green solution;
the reddish brown gas produced is nitrogen dioxide.
seCtIon 1.3 Properties of Matter 13
While reading a textbook of chemistry, I came upon the statement “nitric acid acts upon
copper,” and I determined to see what this meant. Having located some nitric acid, I had
only to learn what the words “act upon” meant. In the interest of knowledge I was even
willing to sacrifice one of the few copper cents then in my possession. I put one of them
on the table, opened a bottle labeled “nitric acid,” poured some of the liquid on the cop-
per, and prepared to make an observation. But what was this wonderful thing which I
beheld? The cent was already changed, and it was no small change either. A greenish-blue
liquid foamed and fumed over the cent and over the table. The air became colored dark
red. How could I stop this? I tried by picking the cent up and throwing it out the window.
I learned another fact: nitric acid acts upon fingers. The pain led to another unpremedi-
tated experiment. I drew my fingers across my trousers and discovered nitric acid acts
upon trousers. That was the most impressive experiment I have ever performed. I tell of it
even now with interest. It was a revelation to me. Plainly the only way to learn about such
remarkable kinds of action is to see the results, to experiment, to work in the laboratory.*

Which of these changes are physical and which are chemical? explain.
(a) Plants make sugar from carbon dioxide and water.
(b) Water vapor in the air forms frost.
(c) a goldsmith melts a nugget of gold and pulls it into a wire.
Separation of Mixtures
We can separate a mixture into its components by taking advantage of differences in
their properties. For example, a heterogeneous mixture of iron filings and gold filings
could be sorted by color into iron and gold. A less tedious approach would be to use a
magnet to attract the iron filings, leaving the gold ones behind. We can also take ad-
vantage of an important chemical difference between these two metals: Many acids dis-
solve iron but not gold. Thus, if we put our mixture into an appropriate acid, the acid
would dissolve the iron and the solid gold would be left behind. The two could then be
separated by filtration (▶ Figure 1.12). We would have to use other chemical reactions,
which we will learn about later, to transform the dissolved iron back into metal.
An important method of separating the components of a homogeneous mixture
is distillation, a process that depends on the different abilities of substances to form
gases. For example, if we boil a solution of salt and water, the water evaporates, forming
a gas, and the salt is left behind. The gaseous water can be converted back to a liquid on
the walls of a condenser, as shown in ▼ Figure 1.13.
2 Water is condensed,
1 Boiling the solution and then collected in
vaporizes the water the receiving flask
Condenser
Salt water
Cold water
out
Cold water
in
3 After water has boiled away, Pure water ▲ Figure 1.12 Separation by filtration.
pure sodium chloride remains in receiving flask a mixture of a solid and a liquid is poured
through filter paper. The liquid passes
▲ Figure 1.13 Distillation. apparatus for separating a sodium chloride solution (salt through the paper while the solid remains
water) into its components. on the paper.
*Remsen, Ira, The Principles of Theoretical Chemistry, 1887.

GO FIGURE
is the separation of a, b, and c in Figure 1.14 a physical or chemical process?
I II III
Solvent
flow of solvent
Mixture of a a
compounds b+c
(a + b + c) b
c
Adsorbent
(stationary
phase)
Glass
Compounds a, b, and c wool Stopcock
are adsorbed to different
degrees on the solid
stationary phase
▲ Figure 1.14 Separation of three substances using column chromatography.
The differing abilities of substances to adhere to the surfaces of solids can also
be used to separate mixtures. This ability is the basis of chromatography, a technique
shown in ▲ Figure 1.14.
1.4 | Units of Measurement

Many properties of matter are quantitative, that is, associated with numbers. When a
number represents a measured quantity, the units of that quantity must be specified.
To say that the length of a pencil is 17.5 is meaningless. Expressing the number with its
units, 17.5 centimeters (cm), properly specifies the length. The units used for scientific
▲ Figure 1.15 Metric units. Metric measurements are those of the metric system.
measurements are increasingly common in
The metric system, developed in France during the late eighteenth century, is used
the United States, as exemplified by the
volume printed on this soda can in both as the system of measurement in most countries. The United States has traditionally
english units (fluid ounces, fl oz) and metric used the English system, although use of the metric system has become more common
units (milliliters, ml). (◀ Figure 1.15).
A Closer Look
The Scientific Method known as a hypothesis, to explain the observations. Initially the hy-
pothesis is likely to be pretty tentative. There could be more than one
Where does scientific knowledge come from? How is it acquired? How reasonable hypothesis. If a hypothesis is correct, then certain results
do we know it is reliable? How do scientists add to it, or modify it? and observations should follow from it. In this way hypotheses can
There is nothing mysterious about how scientists work. The first stimulate the design of experiments to learn more about the system
idea to keep in mind is that scientific knowledge is gained through being studied. Scientific creativity comes into play in thinking of hy-
observations of the natural world. A principal aim of the scientist is potheses that are fruitful in suggesting good experiments to do, ones
to organize these observations, by identifying patterns and regularity, that will shed new light on the nature of the system.
making measurements, and associating one set of observations with As more information is gathered, the initial hypotheses get
another. The next step is to ask why nature behaves in the manner we winnowed down. Eventually just one may stand out as most consis-
observe. To answer this question, the scientist constructs a model, tent with a body of accumulated evidence. We then begin to call this
seCtIon 1.4 Units of Measurement 15
hypothesis a theory, a model that has predictive powers, and that ac- happens, to the best of our knowledge. A theory, on the other hand,
counts for all the available observations. A theory also generally is is an explanation for what happens. If we discover some law fails to
consistent with other, perhaps larger and more general theories. For hold true, then we must assume the theory underlying that law is
example, a theory of what goes on inside a volcano has to be consistent wrong in some way.
with more general theories regarding heat transfer, chemistry at high
Related Exercises: 1.60, 1.82
temperature, and so forth.
We will be encountering many theories as we proceed through
this book. Some of them have been found over and over again to
be consistent with observations. However, no theory can be proven
to be absolutely true. We can treat it as though it is, but there Collect information via
always remains a possibility that there is some respect in which a observations of natural
theory is wrong. A famous example is Einstein’s theory of relativ- phenomena and experiments
ity. Isaac Newton’s theory of mechanics yielded such precise results
for the mechanical behavior of matter that no exceptions to it were
found before the twentieth century. But Albert Einstein showed
that Newton’s theory of the nature of space and time is incorrect.
Formulate one or more
Einstein’s theory of relativity represented a fundamental shift in explanatory hypotheses
how we think of space and time. He predicted where the exceptions
to predictions based on Newton’s theory might be found. Although
only small departures from Newton’s theory were predicted, they
were observed. Einstein’s theory of relativity became accepted
as the correct model. However, for most uses, Newton’s laws of Perform experiments to
test the hypotheses
motion are quite accurate enough.
The overall process we have just considered, illustrated in
▶ Figure 1.16 , is often referred to as the scientific method. But
there is no single scientific method. Many factors play a role in
advancing scientific knowledge. The one unvarying requirement is Use the most successful
hypotheses to formulate
that our explanations be consistent with observations, and that they
a theory
depend solely on natural phenomena.
When nature behaves in a certain way over and over again,
under all sorts of different conditions, we can summarize that
behavior in a scientific law. For example, it has been repeatedly
observed that in a chemical reaction there is no change in the Repeatedly test theory.
total mass of the materials reacting as compared with the materi- Modify as needed to match
experimental results, or reject.
als that are formed; we call this observation the Law of Conserva-
tion of Mass. It is important to make a distinction between a theory
and a scientific law. The latter simply is a statement of what always ▲ Figure 1.16 The scientific method.
SI Units
In 1960 an international agreement was reached specifying a particular choice of metric
units for use in scientific measurements. These preferred units are called SI units, after
the French Système International d’Unités. This system has seven base units from which
all other units are derived (▼ Table 1.4). In this chapter we will consider the base units
for length, mass, and temperature.
Table 1.4 SI Base Units

Physical Quantity Name of Unit Abbreviation
Mass Kilogram kg
Length Meter m
Time Second s or sec
Temperature Kelvin K
Amount of substance Mole mol
Electric current Ampere A or amp
Luminous intensity Candela cd

The package of a fluorescent bulb for a table lamp lists the light output in terms
of lumens, lm. Which of the seven Si units would you expect to be part of the
definition of a lumen?
With SI units, prefixes are used to indicate decimal fractions or multiples of vari-
ous units. For example, the prefix milli- represents a 10-3 fraction, one-thousandth, of
a unit: A milligram (mg) is 10-3 gram (g), a millimeter (mm) is 10-3 meter (m), and so
forth. ▼ Table 1.5 presents the prefixes commonly encountered in chemistry. In using
SI units and in working problems throughout this text, you must be comfortable using
exponential notation. If you are unfamiliar with exponential notation or want to review
it, refer to Appendix A.1.
Although non–SI units are being phased out, some are still commonly used by sci-
entists. Whenever we first encounter a non–SI unit in the text, the SI unit will also be
given. The relations between the non–SI and SI units we will use most frequently in this
text appear on the back inside cover. We will discuss how to convert from one to the
other in Section 1.6.
Table 1.5 Prefixes Used in the Metric System and with SI Units
Prefix Abbreviation Meaning Example
Peta P 1015 1 petawatt (PW) = 1 * 1015 wattsa
Tera T 1012 1 terawatt (TW) = 1 * 1012 watts
Giga G 109 1 gigawatt (GW) = 1 * 109 watts
Mega M 106 1 megawatt (MW) = 1 * 106 watts
Kilo k 103 1 kilowatt (kW) = 1 * 103 watts
Deci d 10-1 1 deciwatt (dW) = 1 * 10-1 watt
Centi c 10-2 1 centiwatt (cW) = 1 * 10-2 watt
Milli m 10-3 1 milliwatt (mW) = 1 * 10-3 watt
Micro mb 10-6 1 microwatt 1mW2 = 1 * 10-6 watt
Nano n 10-9 1 nanowatt (nW) = 1 * 10-9 watt
Pico p 10-12 1 picowatt (pW) = 1 * 10-12 watt
Femto f 10-15 1 femtowatt (fW) = 1 * 10-15 watt
Atto a 10-18 1 attowatt (aW) = 1 * 10-18 watt
Zepto z 10-21 1 zeptowatt (zW) = 1 * 10-21 watt
or consumed. The SI unit of energy is the joule (J); 1 J = 1 kg # m2 >s2 and 1 W = 1 J>s.
a
The watt (W) is the SI unit of power, which is the rate at which energy is either generated
b
Greek letter mu, pronounced “mew.”

how many mg are there in 1 mg?
Length and Mass

The SI base unit of length is the meter, a distance slightly longer than a yard. Mass* is a
measure of the amount of material in an object. The SI base unit of mass is the kilogram (kg),
which is equal to about 2.2 pounds (lb). This base unit is unusual because it uses a pre-
fix, kilo-, instead of the word gram alone. We obtain other units for mass by adding
prefixes to the word gram.
SAMPLE
EXERCISE 1.2 Using SI Prefixes
What is the name of the unit that equals (a) 10-9 gram, (b) 10-6 second, (c) 10-3 meter?
SOLUTION
We can find the prefix related to each power of ten in Table 1.5: (a) nanogram, ng; (b) microsec-
ond, ms; (c) millimeter, mm.
Practice Exercise 1
Which of the following weights would you expect to be suitable for weighing on an ordinary
bathroom scale?
(a) 2.0 * 107 mg, (b) 2500 mg, (c) 5 * 10-4 kg, (d) 4 * 106 cg, (e) 5.5 * 108 dg.
Practice Exercise 2
(a) How many picometers are there in 1 m? (b) Express 6.0 * 103 m using a prefix to replace
the power of ten. (c) Use exponential notation to express 4.22 mg in grams. (d) Use decimal
notation to express 4.22 mg in grams.
Temperature
Temperature, a measure of the hotness or coldness of an object, is a physical property
that determines the direction of heat flow. Heat always flows spontaneously from a sub-
stance at higher temperature to one at lower temperature. Thus, the influx of heat we
feel when we touch a hot object tells us that the object is at a higher temperature than
our hand.
The temperature scales commonly employed in science are the Celsius
and Kelvin scales. The Celsius scale was originally based on the assignment of
0 °C to the freezing point of water and 100 °C to its boiling point at sea level
( Figure 1.17).
*Mass and weight are often incorrectly thought to be the same. The weight of an object is the force
that is exerted on its mass by gravity. In space, where gravitational forces are very weak, an astronaut
can be weightless, but he or she cannot be massless. The astronaut’s mass in space is the same as it is
on Earth.
GO FIGURE
True or false: The “size” of a degree on the Celsius scale is the same as the “size” of a degree on the Kelvin scale.
373 K 100 °C 212 °F Water boils

100 degree-intervals
310 K 37.0 °C 98.6 °F Normal body temperature
273 K 0 °C 32 °F Water freezes
Kelvin scale Celsius scale Fahrenheit scale

▲ Figure 1.17 Comparison of the Kelvin, Celsius, and Fahrenheit temperature scales.
The Kelvin scale is the SI temperature scale, and the SI unit of temperature is the
kelvin (K). Zero on the Kelvin scale is the lowest attainable temperature, referred to
as absolute zero. On the Celsius scale, absolute zero has the value, -273.15 °C. The
Celsius and Kelvin scales have equal-sized units—that is, a kelvin is the same size as a
degree Celsius. Thus, the Kelvin and Celsius scales are related according to
K = °C + 273.15 [1.1]
degree sign 1°2 with temperatures on the Kelvin scale.

The freezing point of water, 0 °C, is 273.15 K (Figure 1.17). Notice that we do not use a
The common temperature scale in the United States is the Fahrenheit scale, which
is not generally used in science. Water freezes at 32 °F and boils at 212 °F. The Fahren-
heit and Celsius scales are related according to
1°F - 322 or °F = 1°C2 + 32

5 9
°C = [1.2]
9 5
SAMPLE
EXERCISE 1.3 Converting Units of Temperature
A weather forecaster predicts the temperature will reach 31 °C. What is this temperature (a) in K,
(b) in °F?
SOLUTION
(a) Using Equation 1.1, we have K = 31 + 273 = 304 K. liquid at 525 K (assume samples are protected from air):
(b) Using Equation 1.2, we have (a) bismuth, Bi; (b) platinum, Pt; (c) selenium, Se; (d) calcium, Ca;
(e) copper, Cu.
1312 + 32 = 56 + 32 = 88 °F.
9
°F =
5
Practice Exercise 2
Practice Exercise 1 Ethylene glycol, the major ingredient in antifreeze, freezes at
Using Wolfram Alpha (http://www.wolframalpha.com/) or -11.5 °C. What is the freezing point in (a) K, (b) °F?
some other reference, determine which of these elements would be
Derived SI Units
The SI base units are used to formulate derived units. A derived unit is obtained by
multiplication or division of one or more of the base units. We begin with the defin- GO FIGURE
ing equation for a quantity and, then substitute the appropriate base units. For exam-
how many 1-l bottles are required
ple, speed is defined as the ratio of distance traveled to elapsed time. Thus, the SI unit
to contain 1 m3 of liquid?
for speed—m/s, read “meters per second”—is a derived unit, the SI unit for distance
(length), m, divided by the SI unit for time, s. Two common derived units in chemistry 1m 1m
are those for volume and density.
Volume
1m
The volume of a cube is its length cubed, length3. Thus, the derived SI unit of volume is
the SI unit of length, m, raised to the third power. The cubic meter, m3, is the volume
of a cube that is 1 m on each edge (▶ Figure 1.18). Smaller units, such as cubic cen-
timeters, cm3 (sometimes written cc), are frequently used in chemistry. Another vol-
1 dm3 = 1 L
ume unit used in chemistry is the liter (L), which equals a cubic decimeter, dm3, and is
slightly larger than a quart. (The liter is the first metric unit we have encountered that is
not an SI unit.) There are 1000 milliliters (mL) in a liter, and 1 mL is the same volume
1 cm3 = 1 mL
as 1 cm3: 1 mL = 1 cm3. The devices used most frequently in chemistry to measure vol-
ume are illustrated in ▼ Figure 1.19.
Syringes, burettes, and pipettes deliver amounts of liquids with more precision
1 cm
than graduated cylinders. Volumetric flasks are used to contain specific volumes
of liquid. 1 cm 1 cm
▲ Figure 1.18 Volume relationships.
The volume occupied by a cube 1 m
Give It Some Thought on each edge is one cubic meter, 1 m3.
Which of the following quantities represents volume measurement: each cubic meter contains 1000 dm3.
15 m2; 2.5 * 102 m3; 5.77 l>s? how do you know? one liter is the same volume as one
cubic decimeter, 1 l = 1 dm3. each
cubic decimeter contains 1000 cubic
Density centimeters, 1 dm3 = 1000 cm3. one
cubic centimeter equals one milliliter,
Density is defined as the amount of mass in a unit volume of a substance: 1 cm3 = 1 ml.
mass
Density = [1.3]
volume
These deliver variable volumes Pipette delivers a Volumetric flask contains

specific volume a specific volume
mL 0
1
2
mL 100 3
90 4
5
80
70 45
60 46
50 47
40 48
30 49
50
20 Stopcock,
10 a valve to
control the
liquid flow
Graduated cylinder Syringe Burette Pipette Volumetric flask
▲ Figure 1.19 Common volumetric glassware.

cubic centimeter 1g>cm32 or grams per milliliter 1g>mL2. The densities of some com-
The densities of solids and liquids are commonly expressed in either grams per
Table 1.6 Densities of Selected
Substances at 25 °C mon substances are listed in ◀ Table 1.6. It is no coincidence that the density of water
Substance Density 1g , cm3 2 is 1.00 g>mL; the gram was originally defined as the mass of 1 mL of water at a specific
temperature. Because most substances change volume when they are heated or cooled,
Air 0.001
densities are temperature dependent, and so temperature should be specified when re-
Balsa wood 0.16 porting densities. If no temperature is reported, we assume 25 °C, close to normal room
Ethanol 0.79 temperature.
The terms density and weight are sometimes confused. A person who says that iron
Water 1.00
weighs more than air generally means that iron has a higher density than air—1 kg of
Ethylene glycol 1.09 air has the same mass as 1 kg of iron, but the iron occupies a smaller volume, thereby
Table sugar 1.59 giving it a higher density. If we combine two liquids that do not mix, the less dense liq-
uid will float on the denser liquid.
Table salt 2.16
Iron 7.9
Gold 19.32
SAMPLE
EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass
(a) Calculate the density of mercury if 1.00 * 102 g occupies a volume of 7.36 cm3.
(b) Calculate the volume of 65.0 g of liquid methanol (wood alcohol) if its density is 0.791 g>mL.
(c) What is the mass in grams of a cube of gold 1density = 19.32 g>cm32 if the length of the cube
is 2.00 cm?
SOLUTION
(a) We are given mass and volume, so Equation 1.3 yields Practice Exercise 1
mass 1.00 * 102 g Platinum, Pt, is one of the rarest of the metals. Worldwide annual
Density = = = 13.6 g>cm3 production is only about 130 tons. (a) Platinum has a density of
volume 7.36 cm3
21.4 g>cm3. If thieves were to steal platinum from a bank using a
(b) Solving Equation 1.3 for volume and then using the given mass small truck with a maximum payload of 900 lb, how many 1 L bars
mass 65.0 g of the metal could they make off with? (a) 19 bars, (b) 2 bars,
and density gives Volume = = = 82.2 mL (c) 42 bars, (d) 1 bar, (e) 47 bars.
density 0.791 g>mL
(c) We can calculate the mass from the volume of the cube and its
density. The volume of a cube is given by its length cubed: Practice Exercise 2
Volume = 12.00 cm23 = 12.0023 cm3 = 8.00 cm3

(a) Calculate the density of a 374.5-g sample of copper if it has a
volume of 41.8 cm3. (b) A student needs 15.0 g of ethanol for an
Solving Equation 1.3 for mass and substituting the volume and experiment. If the density of ethanol is 0.789 g>mL, how many
25.0 mL of mercury 1density = 13.6 g>mL2?

density of the cube, we have milliliters of ethanol are needed? (c) What is the mass, in grams, of
Mass = volume * density = 18.00 cm32119.32 g>cm32 = 155 g
Chemistry Put to Work
Chemistry in the News especially valuable in powering data centers which consume large amounts
of electrical power that must be absolutely reliable. For example, failure
Because chemistry is so central to our lives, reports on matters of chem- of electrical power at a major data center for a company such as Amazon,
ical significance appear in the news nearly every day. Some reports tell eBay, or Apple could be calamitous for the company and its customers.
of breakthroughs in the development of new pharmaceuticals, materi- eBay recently contracted to build the next phase of its major data cen-
als, and processes. Others deal with energy, environmental, and public ter in Utah, utilizing solid–state fuel cells as the source of electrical power.
safety issues. As you study chemistry, you will develop the skills to better The fuel cells, manufactured by Bloom Energy, a Silicon Valley startup, are
understand the importance of chemistry in your life. Here are summa- large industrial devices about the size of a refrigerator (▶ Figure 1.20).
ries of a few recent stories in which chemistry plays an important role. The eBay installation utilizes biogas, which consists of methane and other
Clean energy from fuel cells. In fuel cells, the energy of a chemical fuel gases derived from landfills and farms. The fuel is combined with
reaction is converted directly into electrical energy. Although fuel cells oxygen, and the mixture run through a special solid–state device to pro-
have long been known as potentially valuable sources of electrical energy, duce electricity. Because the electricity is being produced close to the data
their costs have kept them from widespread use. However, recent advanc- center, transmission of the electrical power from source to consumption
es in technology have brought fuel cells to the fore as sources of reliable is more efficient. In contrast to electrical backup systems employed in the
and clean electrical power in certain critical situations. They are past, the new power source will be the primary source of power, operating
ably painful. The word anesthesia was sug-

gested by Oliver Wendell Holmes, Sr. in 1846
to describe the state in which a person lacks
awareness, either total or of a particular part of
the body. In time chemists were able to iden-
tify certain organic compounds that produced
anesthesia without being severely toxic.
More than 40 million patients in North
America each year undergo medical proce-
dures that call for anesthesia. The anesthet-
ics used today are most often injected into
the blood stream rather than inhaled as a gas.
Several organic substances have been identi-
fied as effective anesthetics. While modern
anesthetics are generally quite safe, they must
be administered with care, because they can
affect breathing, blood pressure, and heart
function. Every drug has a therapeutic index,
the ratio of the smallest dose that would be fa-
tal to the smallest dose that gives the desired
therapeutic effect. Naturally, one wants the
therapeutic index for any drug to be as large as
possible. Anesthetics have generally low thera-
peutic indices, which means that they must
▲ Figure 1.20 Solid-State fuel cells manufactured by Bloom Energy. be administered carefully and with constant
monitoring. The death of the entertainer Mi-
24 hours per day, every day of the year. The eBay facility in Utah is the chael Jackson in June 2009 from an overdose of propofol, a widely used
largest nonelectric utility fuel cell installation in the nation. It generates anesthetic (▼ Figure 1.21), illustrates how dangerous such drugs can
6 megawatts of power, enough to power about 6000 homes. be when not properly administered. Propofol very quickly renders a pa-
tient unconscious and affects breathing. Hence its use must be carefully
Regulation of greenhouse gases. In 2009 the Environmental Pro-
monitored by a person trained in anesthesiology.
tection Agency (EPA) took the position that, under the provisions of
Despite a great deal of research, it is still not clear how anesthet-
the Clean Air Act, it should regulate emissions of “greenhouse” gases.
ics actually work. It is a near-universal characteristic of life that spe-
Greenhouse gases are substances that have the potential to alter the
cies ranging from tadpoles to humans can be reversibly immobilized.
global climate because of their ability to trap long–wavelength radia-
The search for the mechanisms by which this occurs is important, be-
Greenhouse gases include carbon dioxide 1CO22, methane 1CH42, and
tion at Earth’s surface. (This subject is covered in detail in Section 18.2.)
cause it may lead us not only to safer anesthetics, but also to deeper
nitrous oxide 1N2O2, as well as other substances. The EPA decision was
understanding of what we mean by consciousness itself.
challenged in the courts by several states, industry organizations, and
conservative groups. In a major victory for the EPA, the federal court
of appeals of the District of Columbia in July 2012 upheld the agen-
cy’s position. This case is interesting in part because of the grounds on
which the EPA policy was challenged, and the way the court responded.
The plaintiffs argued that the EPA improperly based its decision on as-
sessments from the Intergovernmental Panel on Climate Change, the
U.S. Global Climate Change program, and reports from the Nation-
al Research Council, rather than on citing the findings of individual
research programs in the published literature. The court replied that “it
makes no difference that much of the scientific evidence in large part
consisted of ‘syntheses’ of individual studies and research. This is how
science works. EPA is not required to re-prove the existence of the atom
every time it approaches a scientific question.”*
This is an important example of an interaction between science
and social policy in our complex, modern society. When other than
purely scientific interests are involved, questions about science’s reli-
ability and objectivity are bound to arise.
Anesthesia. In the period around the 1840s it became recognized
that certain substances, notably ether, chloroform, and nitrous oxide,
could induce a state in which the patient had no awareness of bodily
pain. You can imagine how joyfully these new discoveries were received
by people who had to undergo surgery that would otherwise be unbear- ▲ Figure 1.21 Propofol, an anesthetic.
*U.S. Court of Appeals for the District of Columbia , Case No. 09-1322.
1.5 | Uncertainty in Measurement

GO FIGURE
how would the darts be positioned Two kinds of numbers are encountered in scientific work: exact numbers (those whose
on the target for the case of “good values are known exactly) and inexact numbers (those whose values have some uncer-
accuracy, poor precision”? tainty). Most of the exact numbers we will encounter in this book have defined values.
For example, there are exactly 12 eggs in a dozen, exactly 1000 g in a kilogram, and ex-
actly 2.54 cm in an inch. The number 1 in any conversion factor, such as 1 m = 100 cm
or 1 kg = 2.2046 lb, is an exact number. Exact numbers can also result from counting
objects. For example, we can count the exact number of marbles in a jar or the exact
number of people in a classroom.
Numbers obtained by measurement are always inexact. The equipment used to
measure quantities always has inherent limitations (equipment errors), and there are
differences in how different people make the same measurement (human errors). Sup-
pose ten students with ten balances are to determine the mass of the same dime. The
Good accuracy ten measurements will probably vary slightly for various reasons. The balances might
Good precision be calibrated slightly differently, and there might be differences in how each student
reads the mass from the balance. Remember: Uncertainties always exist in measured
quantities.

Which of the following is an inexact quantity?
(a) the number of people in your chemistry class
(b) the mass of a penny
(c) the number of grams in a kilogram
Poor accuracy
Good precision
Precision and Accuracy
The terms precision and accuracy are often used in discussing the uncertainties of mea-
sured values. Precision is a measure of how closely individual measurements agree
with one another. Accuracy refers to how closely individual measurements agree with
the correct, or “true,” value. The dart analogy in ◀ Figure 1.22 illustrates the difference
between these two concepts.
In the laboratory we often perform several “trials” of an experiment and aver-
age the results. The precision of the measurements is often expressed in terms of
the standard deviation (Appendix A.5), which reflects how much the individual
measurements differ from the average. We gain confidence in our measurements
Poor accuracy if we obtain nearly the same value each time—that is, when the standard deviation
Poor precision
is small. Figure 1.22 reminds us, however, that precise measurements can be inac-
▲ Figure 1.22 Precision and accuracy. curate. For example, if a very sensitive balance is poorly calibrated, the masses we
measure will be consistently either high or low. They will be inaccurate even if they
are precise.
Significant Figures
Suppose you determine the mass of a dime on a balance capable of measuring to
the nearest 0.0001 g. You could report the mass as 2.2405 { 0.0001 g. The { no-
tation (read “plus or minus”) expresses the magnitude of the uncertainty of your
measurement. In much scientific work we drop the { notation with the under-
standing that there is always some uncertainty in the last digit reported for any mea-
sured quantity.
▶ Figure 1.23 shows a thermometer with its liquid column between two scale
marks. We can read the certain digits from the scale and estimate the uncertain one.
Seeing that the liquid is between the 25° and 30 °C marks, we estimate the temperature
high precision can be achieved on a scale to be 27 °C, being uncertain of the second digit of our measurement. By uncertain we
like this one, which has 0.1 milligram mean that the temperature is reliably 27 °C and not 28° or 26 °C, but we can’t say that
accuracy. it is exactly 27 °C.
seCtIon 1.5 Uncertainty in Measurement 23
◀ Figure 1.23 Uncertainty and significant

figures in a measurement.
100 °C
80 °C
60 °C 30 °C Second digit in 27 °C is
estimated and therefore
uncertain
40 °C
27 °C
20 °C 25 °C
0 °C
All digits of a measured quantity, including the uncertain one, are called signifi-
cant figures. A measured mass reported as 2.2 g has two significant figures, whereas
one reported as 2.2405 g has five significant figures. The greater the number of signifi-
cant figures, the greater the precision implied for the measurement.
SAMPLE
EXERCISE 1.5 Relating Significant Figures to the Uncertainty of a Measurement
What difference exists between the measured values 4.0 and 4.00 g?
SOLUTION
The value 4.0 has two significant figures, whereas 4.00 has three. Practice Exercise 1
This difference implies that 4.0 has more uncertainty. A mass Mo Farah won the 10,000 meter race in the 2012 Olympics with
reported as 4.0 g indicates that the uncertainty is in the first decimal an official time of 27 minutes, 30.42 s. To the correct number of
place. Thus, the mass is closer to 4.0 than to 3.9 or 4.1 g. We can rep- significant figures, what was Farah’s average speed in m/sec?
resent this uncertainty by writing the mass as 4.0 { 0.1 g. A mass (a) 0. 6059 m/s, (b) 1.65042 m/s, (c) 6.059064 m/s, (d) 0.165042 m/s,
reported as 4.00 g indicates that the uncertainty is in the second
(e) 6.626192 m/s.
decimal place. In this case the mass is closer to 4.00 than 3.99 or
4.01 g, and we can represent it as 4.00 { 0.01 g. (Without further Practice Exercise 2
information, we cannot be sure whether the difference in uncertain-
ties of the two measurements reflects the precision or the accuracy A sample that has a mass of about 25 g is weighed on a balance that
of the measurement.) has a precision of {0.001 g. How many significant figures should
be reported for this measurement?

a digital bathroom scale gives you the following four readings in a row: 155.2,
154.8, 154.9, 154.8 lbs. how would you record your weight?
To determine the number of significant figures in a reported measurement, read

the number from left to right, counting the digits starting with the first digit that is not
zero. In any measurement that is properly reported, all nonzero digits are significant.
Because zeros can be used either as part of the measured value or merely to locate the
decimal point, they may or may not be significant:
1. Zeros between nonzero digits are always significant—1005 kg (four significant
figures); 7.03 cm (three significant figures).
2. Zeros at the beginning of a number are never significant; they merely indicate the
position of the decimal point—0.02 g (one significant figure); 0.0026 cm (two sig-
nificant figures).
3. Zeros at the end of a number are significant if the number contains a decimal
point—0.0200 g (three significant figures); 3.0 cm (two significant figures).
A problem arises when a number ends with zeros but contains no decimal point.
In such cases, it is normally assumed that the zeros are not significant. Exponential
notation (Appendix A.1) can be used to indicate whether end zeros are significant. For
example, a mass of 10,300 g can be written to show three, four, or five significant fig-
ures depending on how the measurement is obtained:
1.03 * 104 g (three significant figures)

4 (four significant figures)
1.030 * 10 g
4 (five significant figures)
1.0300 * 10 g
In these numbers all the zeros to the right of the decimal point are significant (rules
1 and 3). (The exponential term 104 does not add to the number of significant
figures.)
SAMPLE
EXERCISE 1.6 Assigning Appropriate Significant Figures
The state of Colorado is listed in a road atlas as having a population of 4,301,261 and an area of
104,091 square miles. Do the numbers of significant figures in these two quantities seem reason-
able? If not, what seems to be wrong with them?
SOLUTION
The population of Colorado must vary from day to day as people Practice Exercise 1
move in or out, are born, or die. Thus, the reported number suggests Which of the following numbers in your personal life are exact
a much higher degree of accuracy than is possible. Secondly, it would numbers?
not be feasible to actually count every individual resident in the state (a) Your cell phone number, (b) your weight, (c) your IQ,
at any given time. Thus, the reported number suggests far greater (d) your driver’s license number, (e) the distance you walked
precision than is possible. A reported number of 4,300,000 would yesterday.
better reflect the actual state of knowledge.
The area of Colorado does not normally vary from time to time, so
the question here is whether the accuracy of the measurements is Practice Exercise 2
good to six significant figures. It would be possible to achieve such The back inside cover of the book tells us that there are 5280 ft in
accuracy using satellite technology, provided the legal boundaries are 1 mile. Does this make the mile an exact distance?
known with sufficient accuracy.
SAMPLE
EXERCISE 1.7 Determining the Number of Significant Figures in a Measurement
How many significant figures are in each of the following numbers (assume that each number is
a measured quantity)? (a) 4.003, (b) 6.023 * 1023, (c) 5000.
SOLUTION
(a) Four; the zeros are significant figures. (b) Four; the exponential thermometer placed under her tongue and gets a value of 102.8 °F.
term does not add to the number of significant figures. (c) One; we How many significant figures are in this measurement?
assume that the zeros are not significant when there is no decimal (a) Three, the number of degrees to the left of the decimal point;
point shown. If the number has more significant figures, a decimal (b) four, the number of digits in the measured reading; (c) two, the
point should be employed or the number written in exponential number of digits in the difference between her current reading and
notation. Thus, 5000. has four significant figures, whereas her normal body temperature; (d) three, the number of digits in
5.00 * 103 has three. her normal body temperature; (e) one, the number of digits to the
right of the decimal point in the measured value.
Practice Exercise 1 Practice Exercise 2

Sylvia feels as though she may have a fever. Her normal body How many significant figures are in each of the following mea-
temperature is 98.7 °F. She measures her body temperature with a surements? (a) 3.549 g, (b) 2.3 * 104 cm, (c) 0.00134 m3.
seCtIon 1.5 Uncertainty in Measurement 25
Significant Figures in Calculations

When carrying measured quantities through calculations, the least certain measure-
ment limits the certainty of the calculated quantity and thereby determines the number
of significant figures in the final answer. The final answer should be reported with only
one uncertain digit. To keep track of significant figures in calculations, we will make
frequent use of two rules: one for addition and subtraction, and another for multiplica-
tion and division.
1. For addition and subtraction, the result has the same number of decimal places as
the measurement with the fewest decimal places. When the result contains more
than the correct number of significant figures, it must be rounded off. Consider the
following example in which the uncertain digits appear in color:
This number limits 20.42 — two decimal places

the number of significant 1.322 — three decimal places
figures in the result ¡ 83.1 — one decimal place
104.842 — round off to one decimal place (104.8)
We report the result as 104.8 because 83.1 has only one decimal place.
2. For multiplication and division, the result contains the same number of sig-
nificant figures as the measurement with the fewest significant figures. When
the result contains more than the correct number of significant figures, it must
be rounded off. For example, the area of a rectangle whose measured edge lengths
are 6.221 and 5.2 cm should be reported with two significant figures, 32 cm2, even
though a calculator shows the product to have more digits:
Area = 16.221 cm215.2 cm2 = 32.3492 cm2 1 round off to 32 cm2
because 5.2 has two significant figures.
Notice that for addition and subtraction, decimal places are counted in determining how
many digits to report in an answer, whereas for multiplication and division, significant
figures are counted in determining how many digits to report.
In determining the final answer for a calculated quantity, exact numbers are as-
sumed to have an infinite number of significant figures. Thus, when we say, “There are
12 inches in 1 foot,” the number 12 is exact, and we need not worry about the number
of significant figures in it.
In rounding off numbers, look at the leftmost digit to be removed:
• If the leftmost digit removed is less than 5, the preceding number is left unchanged.
Thus, rounding off 7.248 to two significant figures gives 7.2.
• If the leftmost digit removed is 5 or greater, the preceding number is increased by 1.
Rounding off 4.735 to three significant figures gives 4.74, and rounding 2.376 to
two significant figures gives 2.4.*

a rectangular garden plot is measured to be 25.8 m by 18 m. Which of these
dimensions needs to be measured to greater accuracy to provide a more accurate
estimate of the area of the plot?
*Your instructor may want you to use a slight variation on the rule when the leftmost digit to be removed is
exactly 5, with no following digits or only zeros following. One common practice is to round up to the next
higher number if that number will be even and down to the next lower number otherwise. Thus, 4.7350
would be rounded to 4.74, and 4.7450 would also be rounded to 4.74.
SAMPLE
EXERCISE 1.8 Determining the Number of Significant Figures in a Calculated Quantity
The width, length, and height of a small box are 15.5, 27.3, and 5.4 cm, respectively. Calculate the
volume of the box, using the correct number of significant figures in your answer.
SOLUTION
In reporting the volume, we can show only as many significant figures Practice Exercise 1
as given in the dimension with the fewest significant figures, which is Ellen recently purchased a new hybrid car and wants to check her
that for the height (two significant figures): gas mileage. At an odometer setting of 651.1 mi, she fills the tank.
At 1314.4 mi she requires 16.1 gal to refill the tank. Assuming that
Volume = width * length * height
= 115.5 cm2127.3 cm215.4 cm2
the tank is filled to the same level both times, how is the gas mile-
age best expressed? (a) 40 mi/gal, (b) 41 mi/gal, (c) 41.2 mi/gal,
(d) 41.20 mi/gal.
= 2285.01 cm3 1 2.3 * 103 cm3
Practice Exercise 2
A calculator used for this calculation shows 2285.01, which we must It takes 10.5 s for a sprinter to run 100.00 m. Calculate her average
round off to two significant figures. Because the resulting number is speed in meters per second and express the result to the correct
2300, it is best reported in exponential notation, 2.3 * 103, to clearly number of significant figures.
indicate two significant figures.
SAMPLE
EXERCISE 1.9 Determining the Number of Significant Figures in a Calculated Quantity
A vessel containing a gas at 25 °C is weighed, emptied, and then reweighed as depicted in
▼ Figure 1.24. From the data provided, calculate the density of the gas at 25 °C.
SOLUTION
To calculate the density, we must know both the mass and the volume case each quantity has two decimal places. Thus, the mass of the gas,
of the gas. The mass of the gas is just the difference in the masses of 1.38 g, has two decimal places.
the full and empty container: Using the volume given in the question, 1.05 * 103 cm3, and the defi-
1837.63 - 836.252 g = 1.38 g nition of density, we have
In subtracting numbers, we determine the number of significant fig- mass 1.38 g
ures in our result by counting decimal places in each quantity. In this Density = =
volume 1.05 * 103 cm3
= 1.31 * 10-3 g>cm3 = 0.00131 g>cm3
Pump out gas
In dividing numbers, we determine the number of significant fig-
ures our result should contain by counting the number of significant
figures in each quantity. There are three significant figures in our
answer, corresponding to the number of significant figures in the two
numbers that form the ratio. Notice that in this example, following the
rules for determining significant figures gives an answer containing
only three significant figures, even though the measured masses con-
tain five significant figures.
Practice Exercise 1
Which of the following numbers is correctly rounded to three
(b) 4.5671 * 10-9 34.567 * 10-94, (c) 3.00072 [3.001], (d) 0.006739
significant figures, as shown in brackets? (a) 12,556 [12,500],
[0.00674], (e) 5.4589 * 105 35.459 * 1054.

Practice Exercise 2
Volume: 1.05 × 103 cm3 Mass: 836.25 g If the mass of the container in the sample exercise (Figure 1.24)
Mass: 837.63 g were measured to three decimal places before and after pumping
out the gas, could the density of the gas then be calculated to four
▲ Figure 1.24 Uncertainty and significant figures in a significant figures?
measurement.
When a calculation involves two or more steps and you write answers for intermedi-
ate steps, retain at least one nonsignificant digit for the intermediate answers. This pro-
cedure ensures that small errors from rounding at each step do not combine to affect
the final result. When using a calculator, you may enter the numbers one after another,
seCtIon 1.6 Dimensional analysis 27
rounding only the final answer. Accumulated rounding-off errors may account for
small differences among results you obtain and answers given in the text for numerical
problems.
1.6 | Dimensional Analysis

Because measured quantities have units associated with them, it is important to keep
track of units as well as numerical values when using the quantities in calculations.
Throughout the text we use dimensional analysis in solving problems. In dimen-
sional analysis, units are multiplied together or divided into each other along with
the numerical values. Equivalent units cancel each other. Using dimensional analysis
helps ensure that solutions to problems yield the proper units. Moreover, it provides
a systematic way of solving many numerical problems and of checking solutions for
possible errors.
The key to using dimensional analysis is the correct use of conversion factors to
change one unit into another. A conversion factor is a fraction whose numerator and
denominator are the same quantity expressed in different units. For example, 2.54 cm
and 1 in. are the same length: 2.54 cm = 1 in. This relationship allows us to write two
conversion factors:
2.54 cm 1 in.
and
1 in. 2.54 cm
We use the first factor to convert inches to centimeters. For example, the length in
centimeters of an object that is 8.50 in. long is
Desired unit
2.54 cm
Number of centimeters = (8.50 in.) = 21.6 cm
1 in.
Given unit
The unit inches in the denominator of the conversion factor cancels the unit
inches in the given data (8.50 inches), so that the centimeters unit in the numera-
tor of the conversion factor becomes the unit of the final answer. Because the
numerator and denominator of a conversion factor are equal, multiplying any
quantity by a conversion factor is equivalent to multiplying by the number 1 and
so does not change the intrinsic value of the quantity. The length 8.50 in. is the
same as the length 21.6 cm.
In general, we begin any conversion by examining the units of the given data
and the units we desire. We then ask ourselves what conversion factors we have
available to take us from the units of the given quantity to those of the desired one.
When we multiply a quantity by a conversion factor, the units multiply and divide
as follows:
desired unit
Given unit * = desired unit
given unit
If the desired units are not obtained in a calculation, an error must have been made
somewhere. Careful inspection of units often reveals the source of the error.
SAMPLE
EXERCISE 1.10 Converting Units
If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units
given on the back inside cover of the text.)
SOLUTION
Because we want to change from pounds to grams, we look for a relationship between these units
of mass. The conversion factor table found on the back inside cover tells us that 1 lb = 453.6 g.
To cancel pounds and leave grams, we write the conversion factor with grams in the numerator
Given: lb and pounds in the denominator:
Mass in grams = 1115 lb2a b = 5.22 * 104 g

453.6 g
453.6 g 1 lb
Use
1 lb
The answer can be given to only three significant figures, the number of significant figures in
115 lb. The process we have used is diagrammed in the margin.
Find: g
Practice Exercise 1
At a particular instant in time the Earth is judged to be 92,955,000 miles from the Sun.
What is the distance in kilometers to four significant figures? (See back inside cover
for conversion factor). (a) 5763 * 104 km, (b) 1.496 * 108 km, (c) 1.49596 * 108 km,
(d) 1.483 * 104 km, (e) 57,759,000 km.
Practice Exercise 2
By using a conversion factor from the back inside cover, determine the length in kilometers of
a 500.0-mi automobile race.
Strategies in Chemistry
Estimating Answers done without a calculator. Even though this approach does not give an
exact answer, it gives one that is roughly the correct size. By using di-
Calculators are wonderful devices; they enable you to get to the mensional analysis and by estimating answers, you can readily check
wrong answer very quickly. Of course, that’s not the destination the reasonableness of your calculations.
you want. You can take certain steps to avoid putting that wrong You can get better at making estimates by practicing in every-
answer into your homework set or on an exam. One is to keep track day life. How far is it from your dorm room to the chemistry lecture
of the units in a calculation and use the correct conversion factors. hall? How much do your parents pay for gasoline per year? How many
Second, you can do a quick mental check to be sure that your an- bikes are there on campus? If you respond “I have no idea” to these
swer is reasonable: you can try to make a “ballpark” estimate. questions, you’re giving up too easily. Try estimating familiar quanti-
A ballpark estimate involves making a rough calculation using ties and you’ll get better at making estimates in science and in other
numbers that are rounded off in such a way that the arithmetic can be aspects of your life where a misjudgment can be costly.

how do we determine how many digits to use in conversion factors, such as the
one between pounds and grams in Sample exercise 1.10?
Using Two or More Conversion Factors

It is often necessary to use several conversion factors in solving a problem. As an ex-
ample, let’s convert the length of an 8.00-m rod to inches. The table on the back inside
the relationship between centimeters and inches 11 in. = 2.54 cm2. From our knowl-
cover does not give the relationship between meters and inches. It does, however, give
edge of SI prefixes, we know that 1 cm = 10-2 m. Thus, we can convert step by step,
first from meters to centimeters and then from centimeters to inches:
Given: Find:
Use Use
m cm in.
1 cm 1 in.
10−2 m 2.54 cm
Combining the given quantity (8.00 m) and the two conversion factors, we have
Number of inches = 18.00 m2a ba b = 315 in.

1 cm 1 in.
10-2 m 2.54 cm
The first conversion factor is used to cancel meters and convert the length to centime-
ters. Thus, meters are written in the denominator and centimeters in the numerator.
The second conversion factor is used to cancel centimeters and convert the length to
inches, so it has centimeters in the denominator and inches, the desired unit, in the
numerator.
Note that you could have used 100 cm = 1 m as a conversion factor as well in the
second parentheses. As long as you keep track of your given units and cancel them
properly to obtain the desired units, you are likely to be successful in your calculations.
SAMPLE
EXERCISE 1.11 Converting Units Using Two or More Conversion Factors
The average speed of a nitrogen molecule in air at 25 °C is 515 m>s. Convert this speed to miles per hour.
SOLUTION
To go from the given units, m/s, to the desired units, mi/hr, we must Thus, we can convert m to km and then convert km to mi. From our
convert meters to miles and seconds to hours. From our knowledge of knowledge of time we know that 60 s = 1 min and 60 min = 1 hr.
SI prefixes we know that 1 km = 103 m. From the relationships given Thus, we can convert s to min and then convert min to hr. The overall
on the back inside cover of the book, we find that 1 mi = 1.6093 km. process is
Given: Find:
Use Use Use Use
m/s km/s mi/s mi/min mi/hr
1 km 1 mi 60 s 60 min
103 m 1.6093 km 1 min 1 hr
Applying first the conversions for distance and then those for time, we
can set up one long equation in which unwanted units are canceled:
Speed in mi>hr = a515 ba 3 ba ba ba b

m 1 km 1 mi 60 s 60 min
s 10 m 1.6093 km 1 min 1 hr
= 1.15 * 103 mi>hr
Our answer has the desired units. We can check our calculation, us-
ing the estimating procedure described in the “Strategies in Chem- Practice Exercise 1
istry” box. The given speed is about 500 m>s. Dividing by 1000 Fabiola, who lives in Mexico City, fills her car with gas, paying 357
converts m to km, giving 0.5 km>s. Because 1 mi is about 1.6 km, pesos for 40.0 L. What is her fuel cost in dollars per gallon, if 1
this speed corresponds to 0.5>1.6 = 0.3 mi>s. Multiplying by 60 peso = 0.0759 dollars? (a) $1.18/gal, (b) $3.03/gal, (c) $1.47/gal,
gives about 0.3 * 60 = 20 mi>min. Multiplying again by 60 gives (d) $9.68/gal, (e) $2.56/gal.
20 * 60 = 1200 mi>hr. The approximate solution (about 1200 mi/hr)
and the detailed solution (1150 mi/hr) are reasonably close. The Practice Exercise 2
answer to the detailed solution has three significant figures, cor- A car travels 28 mi per gallon of gasoline. What is the mileage in
responding to the number of significant figures in the given speed kilometers per liter?
in m/s.
Conversions Involving Volume

The conversion factors previously noted convert from one unit of a given measure to
another unit of the same measure, such as from length to length. We also have conver-
sion factors that convert from one measure to a different one. The density of a sub-
Suppose we want to know the mass in grams of 2 cubic inches 12.00 in.32 of gold, which
stance, for example, can be treated as a conversion factor between mass and volume.
has a density of 19.3 g>cm3. The density gives us the conversion factors:
19.3 g 1 cm3
and
1 cm3 19.3 g
Because we want a mass in grams, we use the first factor, which has mass in grams in
the numerator. To use this factor, however, we must first convert cubic inches to cubic
centimeters. The relationship between in.3 and cm3 is not given on the back inside
(exactly). Cubing both sides of this equation gives 11 in.23 = 12.54 cm23, from which
cover, but the relationship between inches and centimeters is given: 1 in. = 2.54 cm
we write the desired conversion factor:
12.54 cm23 12.5423 cm3 16.39 cm3

11 in.2 112 in.
3 = 3 3 =
1 in.3
number, we can retain as many digits of 12.5423 as we need. We have used four, one
Notice that both the numbers and the units are cubed. Also, because 2.54 is an exact
more than the number of digits in the density 119.3 g>cm32. Applying our conversion
factors, we can now solve the problem:
Mass in grams = 12.00 in.32a ba b = 633 g

16.39 cm3 19.3 g
1 in.3 1 cm3
The procedure is diagrammed here. The final answer is reported to three significant
figures, the same number of significant figures as in 2.00 in.3 and 19.3 g.
Given: Find:
Use Use
in.3 cm3 g
3 19.3 g
2.54 cm
1 in. 1 cm3
SAMPLE
EXERCISE 1.12 Converting Volume Units
Earth’s oceans contain approximately 1.36 * 109 km3 of water. Calculate the volume in
liters.
SOLUTION
From the back inside cover, we find 1 L = 10-3 m3, but there is no relationship listed in-
volving km3. From our knowledge of SI prefixes, however, we know 1 km = 103 m and we
can use this relationship between lengths to write the desired conversion factor between
volumes:
a b =
103 m 3 109 m3
1 km 1 km3
How many liters of water do Earth’s
oceans contain?
Thus, converting from km3 to m3 to L, we have
Volume in liters = 11.36 * 109 km32a b a -3 3 b = 1.36 * 1021 L

109 m3 1L
3
1 km 10 m
Practice Exercise 1
A barrel of oil as measured on the oil market is equal to 1.333 U.S. barrels. A U.S. barrel is
equal to 31.5 gal. If oil is on the market at $94.0 per barrel, what is the price in dollars per
gallon? (a) $2.24/gal, (b) $3.98/gal, (c) $2.98/gal, (d) $1.05/gal, (e) $8.42/gal.
Practice Exercise 2
The surface area of Earth is 510 * 106 km2, and 71% of this is ocean. Using the
data from the sample exercise, calculate the average depth of the world’s oceans
in feet.
The Importance of Practice sure to more fully master the content of the chapters by reading them
through at least twice, even more for passages that present you with
If you have ever played a musical instrument or participated in ath- difficulties in understanding.
letics, you know that the keys to success are practice and discipline. Throughout the book, we have provided sample exercises in
You cannot learn to play a piano merely by listening to music, and which the solutions are shown in detail. For practice exercises, we sup-
you cannot learn how to play basketball merely by watching games ply only the answer, at the back of the book. It is important that you
on television. Likewise, you cannot learn chemistry by merely watch- use these exercises to test yourself.
ing your instructor give lectures. Simply reading this book, listening to The practice exercises in this text and the homework assignments
lectures, or reviewing notes will not usually be sufficient when exam given by your instructor provide the minimal practice that you will
time comes around. Your task is to master chemical concepts and need to succeed in your chemistry course. Only by working all the as-
practices to a degree that you can put them to use in solving problems signed problems will you face the full range of difficulty and coverage
and answering questions. Solving problems correctly takes practice— that your instructor expects you to master for exams. There is no sub-
actually, a fair amount of it. You will do well in your chemistry course stitute for a determined and perhaps lengthy effort to work problems
if you embrace the idea that you need to master the materials pre- on your own. If you are stuck on a problem, however, ask for help
sented, and then learn how to apply them in solving problems. Even from your instructor, a teaching assistant, a tutor, or a fellow student.
if you’re a brilliant student, this will take time; it’s what being a stu- Spending an inordinate amount of time on a single exercise is rarely
dent is all about. Almost no one fully absorbs new material on a first effective unless you know that it is particularly challenging and is ex-
reading, especially when new concepts are being presented. You are pected to require extensive thought and effort.
SAMPLE
EXERCISE 1.13 Conversions Involving Density
What is the mass in grams of 1.00 gal of water? The density of water is 1.00 g/mL.
SOLUTION
Before we begin solving this exercise, we note the following:
(1) We are given 1.00 gal of water (the known, or given, quantity) and asked to calculate its mass
in grams (the unknown).
(2) We have the following conversion factors either given, commonly known, or available on the
back inside cover of the text:
1.00 g water 1L 1L 1 gal
1 mL water 1000 mL 1.057 qt 4 qt
The first of these conversion factors must be used as written (with grams in the numerator) to give
the desired result, whereas the last conversion factor must be inverted in order to cancel gallons:
Mass in grams = 11.00 gal2a ba ba ba b

4 qt 1L 1000 mL 1.00 g
1 gal 1.057 qt 1L 1 mL
= 3.78 * 103 g water
The unit of our final answer is appropriate, and we have taken care of our significant figures. We
can further check our calculation by estimating. We can round 1.057 off to 1. Then focusing on the
numbers that do not equal 1 gives 4 * 1000 = 4000 g, in agreement with the detailed calculation.
You should also use common sense to assess the reasonableness of your answer. In this case
we know that most people can lift a gallon of milk with one hand, although it would be tiring to
carry it around all day. Milk is mostly water and will have a density not too different from that of
water. Therefore, we might estimate that a gallon of water has mass that is more than 5 lb but less
than 50 lb. The mass we have calculated, 3.78 kg * 2.2 lb>kg = 8.3 lb, is thus reasonable as an
order-of-magnitude estimate.
Practice Exercise 1
Trex is a manufactured substitute for wood compounded from post-consumer plastic and wood.
It is frequently used in outdoor decks. Its density is reported as 60 lb>ft3. What is the density of
Trex in kg/L? (a) 138 kg/L, (b) 0.960 kg/L, (c) 259 kg/L, (d) 15.8 kg/L, (e) 11.5 kg/L.
Practice Exercise 2
The density of the organic compound benzene is 0.879 g/mL. Calculate the mass in grams of
1.00 qt of benzene.
a Trex deck.
The Features of This Book “Visualizing Concepts” are meant to test how well you understand a
concept without plugging a lot of numbers into a formula. The other
If, like most students, you haven’t yet read the part of the Preface to exercises are grouped in pairs, with the answers given at the back of
this text entitled TO THE STUDENT, you should do it now. In less the book to the odd-numbered exercises (those with red exercise num-
than two pages of reading you will encounter valuable advice on how bers). An exercise with a [bracket] around its number is designed to
to navigate your way through this book and through the course. We’re be more challenging. Additional Exercises appear after the regular
serious! This is advice you can use. exercises; the chapter sections that they cover are not identified, and
The TO THE STUDENT section describes how text features such they are not paired. Integrative Exercises, which start appearing from
as “What’s Ahead,” Key Terms, Learning Outcomes, and Key Equations Chapter 3, are problems that require skills learned in previous chap-
will help you remember what you have learned. We describe there also ters. Also first appearing in Chapter 3, are Design an Experiment ex-
how to take advantage of the text’s Web site, where many types of online ercises consisting of problem scenarios that challenge you to design
study tools are available. If you have registered for MasteringChemistry®, experiments to test hypotheses.
you will have access to many helpful animations, tutorials, and additional Many chemical databases are available, usually on the Web. The
problems correlated to specific topics and sections of each chapter. An in- CRC Handbook of Chemistry and Physics is the standard reference for
teractive eBook is also available online. many types of data and is available in libraries. The Merck Index is a stan-
As previously mentioned, working exercises is very important— dard reference for the properties of many organic compounds, especially
in fact, essential. You will find a large variety of exercises at the end ones of biological interest. WebElements (http://www.webelements
of each chapter that are designed to test your problem-solving skills .com/) is a good Web site for looking up the properties of the elements.
in chemistry. Your instructor will very likely assign some of these Wolfram Alpha (http://www.wolframalpha.com/) can also be a source
end-of-chapter exercises as homework. The first few exercises called of useful information on substances, numerical values, and other data.
Chapter Summary and Key Terms

THE STUDY OF CHEMISTRY (SECTION 1.1) Chemistry is the study the same consistent results, we speak of a scientific law, a general rule
of the composition, structure, properties, and changes of matter. The that summarizes how nature behaves.
composition of matter relates to the kinds of elements it contains.
The structure of matter relates to the ways the atoms of these elements UNITS OF MEASUREMENT (SECTION 1.4) Measurements in chem-
are arranged. A property is any characteristic that gives a sample of mat- istry are made using the metric system. Special emphasis is placed on
ter its unique identity. A molecule is an entity composed of two or more SI units, which are based on the meter, the kilogram, and the second as
atoms with the atoms attached to one another in a specific way. the basic units of length, mass, and time, respectively. SI units use pre-
fixes to indicate fractions or multiples of base units. The SI temperature
CLASSIFICATIONS OF MATTER (SECTION 1.2) Matter exists in scale is the Kelvin scale, although the Celsius scale is frequently used as
three physical states, gas, liquid, and solid, which are known as the well. Absolute zero is the lowest temperature attainable. It has the value
states of matter. There are two kinds of pure substances: elements and 0 K. A derived unit is obtained by multiplication or division of SI base
compounds. Each element has a single kind of atom and is represented units. Derived units are needed for defined quantities such as speed
by a chemical symbol consisting of one or two letters, with the first or volume. Density is an important defined quantity that equals mass
letter capitalized. Compounds are composed of two or more elements divided by volume.
joined chemically. The law of constant composition, also called the law
of definite proportions, states that the elemental composition of a pure UNCERTAINTY IN MEASUREMENT (SECTION 1.5) All measured
compound is always the same. Most matter consists of a mixture of quantities are inexact to some extent. The precision of a measurement
substances. Mixtures have variable compositions and can be either indicates how closely different measurements of a quantity agree with
homogeneous or heterogeneous; homogeneous mixtures are called one another. The accuracy of a measurement indicates how well a
solutions. measurement agrees with the accepted or “true” value. The significant
figures in a measured quantity include one estimated digit, the last
PROPERTIES OF MATTER (SECTION 1.3) Each substance has a digit of the measurement. The significant figures indicate the extent of
unique set of physical properties and chemical properties that can be used the uncertainty of the measurement. Certain rules must be followed so
to identify it. During a physical change, matter does not change its com- that a calculation involving measured quantities is reported with the
position. Changes of state are physical changes. In a chemical change appropriate number of significant figures.
(chemical reaction) a substance is transformed into a chemically different
substance. Intensive properties are independent of the amount of matter DIMENSIONAL ANALYSIS (SECTION 1.6) In the dimensional
examined and are used to identify substances. Extensive properties relate analysis approach to problem solving, we keep track of units as we
to the amount of substance present. Differences in physical and chemi- carry measurements through calculations. The units are multiplied
cal properties are used to separate substances. together, divided into each other, or canceled like algebraic quantities.
The scientific method is a dynamic process used to answer ques- Obtaining the proper units for the final result is an important means
tions about the physical world. Observations and experiments lead to of checking the method of calculation. When converting units and
tentative explanations or hypotheses. As a hypothesis is tested and re- when carrying out several other types of problems, conversion factors
fined, a theory may be developed that can predict the results of future can be used. These factors are ratios constructed from valid relations
observations and experiments. When observations repeatedly lead to between equivalent quantities.
exercises 33
Learning Outcomes After studying this chapter, you should be able to:
• Distinguish among elements, compounds, and mixtures. (Section 1.2) • Demonstrate the use of significant figures, scientific notation, and
• Identify symbols of common elements. (Section 1.2) SI units in calculations. (Section 1.5)
• Identify common metric prefixes. (Section 1.4) • Attach appropriate SI units to defined quantities, and employ
dimensional analysis in calculations. (Sections 1.4 and 1.6)
Key Equations
• K = °C + 273.15 [1.1] Converting between Celsius 1°C2 and Kelvin (K) temperature scales
• °C = 1°F - 322 or °F = 1°C2 + 32 Converting between Celsius 1°C2 and Fahrenheit 1°F2 tempera-
5 9
[1.2]
9 5
ture scales
mass
• Density = [1.3] Definition of density
volume
Exercises
Visualizing Concepts 1.3 Describe the separation method(s) involved in brewing a cup
of coffee. [Section 1.3]
1.1 Which of the following figures represents (a) a pure element,
(b) a mixture of two elements, (c) a pure compound,
(d) a mixture of an element and a compound? (More than
one picture might fit each description.) [Section 1.2]
(i) (ii) (iii)
1.4 Identify each of the following as measurements of length,

area, volume, mass, density, time, or temperature: (a) 25 ps,
(iv) (v) (vi)
(b) 374.2 mg, (c) 77 K, (d) 100,000 km2, (e) 1.06 mm,
1.2 Does the following diagram represent a chemical or physical (f) 16 nm2, (g) -78 °C, (h) 2.56 g>cm3, (i) 28 cm3. [Section 1.4]
1density = 2.70 g>cm32, s i l v e r 1density = 10.49 g>cm32,

change? How do you know? [Section 1.3] 1.5 (a) Three spheres of equal size are composed of aluminum
and nickel 1density = 8.90 g>cm32. List the spheres from
composed of gold 1density = 19.32 g>cm32, platinum

lightest to heaviest. (b) Three cubes of equal mass are
1density = 21.45 g>cm32, and lead 1density = 11.35 g>cm32.

List the cubes from smallest to largest. [Section 1.4]
1.6 The three targets from a rifle range shown on the next page
were produced by: (A) the instructor firing a newly acquired
target rifle; (B) the instructor firing his personal target rifle;
and (C) a student who has fired his target rifle only a few
times. (a) Comment on the accuracy and precision for each
of these three sets of results. (b) For the A and C results in
the future to look like those in B, what needs to happen?
[Section 1.5]
A B C
1.7 (a) What is the length of the pencil in the following figure if
the ruler reads in centimeters? How many significant figures
are there in this measurement? (b) An automobile speed-
ometer with circular scales reading both miles per hour and 1.12 The photo below shows a picture of an agate stone. Jack, who
kilometers per hour is shown. What speed is indicated, in picked up the stone on the Lake Superior shoreline and pol-
both units? How many significant figures are in the measure- ished it, insists that agate is a chemical compound. Ellen ar-
ments? [Section 1.5] gues that it cannot be a compound. Discuss the relative merits
of their positions. [Section 1.2]
1 2 3 4 5 6 7 8 9
Classification and Properties of Matter (Sections

1.2 and 1.3)
1.13 Classify each of the following as a pure substance or a mixture.

If a mixture, indicate whether it is homogeneous or hetero-
1.8 (a) How many significant figures should be reported for the geneous: (a) rice pudding, (b) seawater, (c) magnesium,
volume of the metal bar shown here? (b) If the mass of the bar (d) crushed ice.
is 104.72 g, how many significant figures should be reported 1.14 Classify each of the following as a pure substance or a
when its density is determined using the calculated volume? mixture. If a mixture, indicate whether it is homogeneous or
[Section 1.5] heterogeneous: (a) air, (b) tomato juice, (c) iodine crystals,
(d) sand.
1.15 Give the chemical symbol or name for the following elements,
2.5 cm
as appropriate: (a) sulfur, (b) gold, (c) potassium, (d) chlorine,
(e) copper, (f) U, (g) Ni, (h) Na, (i) Al, (j) Si.
1.25 cm 1.16 Give the chemical symbol or name for each of the follow-
ing elements, as appropriate: (a) carbon, (b) nitrogen,
5.30 cm
(c) titanium, (d) zinc, (e) iron, (f) P, (g) Ca, (h) He,
(i) Pb, (j) Ag.
1.9 When you convert units, how do you decide which part of the 1.17 A solid white substance A is heated strongly in the absence
conversion factor is in the numerator and which is in the de- of air. It decomposes to form a new white substance B and
nominator? [Section 1.6] a gas C. The gas has exactly the same properties as the prod-
1.10 Show the steps to convert the speed of sound, 344 meters per uct obtained when carbon is burned in an excess of oxygen.
second, into miles per hour. [Section 1.6] Based on these observations, can we determine whether solids
1.11 Consider the jar of jelly beans in the photo. To get an estimate A and B and gas C are elements or compounds? Explain your
of the number of beans in the jar you weigh six beans and conclusions for each substance.
obtain masses of 3.15, 3.12, 2.98, 3.14, 3.02, and 3.09 g. Then 1.18 You are hiking in the mountains and find a shiny gold nug-
you weigh the jar with all the beans in it, and obtain a mass of get. It might be the element gold, or it might be “fool’s gold,”
2082 g. The empty jar has a mass of 653 g. Based on these data which is a nickname for iron pyrite, FeS2. What kinds of ex-
estimate the number of beans in the jar. Justify the number of periments could be done to determine if the shiny nugget is
significant figures you use in your estimate. [Section 1.5] really gold?
exercises 35
1.19 In the process of attempting to characterize a substance, 1.28 (a) The temperature on a warm summer day is 87 °F.
a chemist makes the following observations: The sub- What is the temperature in °C? (b) Many scientific data are
stance is a silvery white, lustrous metal. It melts at 649 °C reported at 25 °C. What is this temperature in kelvins and
and boils at 1105 °C. Its density at 20 °C is 1.738 g>cm3. in degrees Fahrenheit? (c) Suppose that a recipe calls for
The substance burns in air, producing an intense white an oven temperature of 400 °F. Convert this temperature to
light. It reacts with chlorine to give a brittle white solid. degrees Celsius and to kelvins. (d) Liquid nitrogen boils at
The substance can be pounded into thin sheets or drawn 77 K. Convert this temperature to degrees Fahrenheit and to
into wires. It is a good conductor of electricity. Which of degrees Celsius.
these characteristics are physical properties, and which are 1.29 (a) A sample of tetrachloroethylene, a liquid used in dry
chemical properties? cleaning that is being phased out because of its potential to
1.20 (a) Read the following description of the element zinc and in- cause cancer, has a mass of 40.55 g and a volume of 25.0 mL
dicate which are physical properties and which are chemical at 25 °C. What is its density at this temperature? Will tetra-
than water will float.) (b) Carbon dioxide 1CO22 is a gas at

properties. chloroethylene float on water? (Materials that are less dense
room temperature and pressure. However, carbon dioxide

can be put under pressure to become a “supercritical fluid”
that is a much safer dry-cleaning agent than tetrachloroethyl-
ene. At a certain pressure, the density of supercritical CO2 is
0.469 g>cm3. What is the mass of a 25.0-mL sample of super-
critical CO2 at this pressure?
1.30 (a) A cube of osmium metal 1.500 cm on a side has a mass of
76.31 g at 25 °C. What is its density in g>cm3 at this tempera-
Zinc melts at 420 °C. When zinc granules are added to dilute ture? (b) The density of titanium metal is 4.51g>cm3 at 25 °C.
sulfuric acid, hydrogen is given off and the metal dissolves. What mass of titanium displaces 125.0 mL of water at 25 °C?
Zinc has a hardness on the Mohs scale of 2.5 and a density (c) The density of benzene at 15 °C is 0.8787g>mL. Calculate
of 7.13g>cm3 at 25 °C. It reacts slowly with oxygen gas at el- the mass of 0.1500 L of benzene at this temperature.
evated temperatures to form zinc oxide, ZnO. 1.31 (a) To identify a liquid substance, a student determined
(b) Which properties of zinc can you describe from the its density. Using a graduated cylinder, she measured out
photo? Are these physical or chemical properties? a 45-mL sample of the substance. She then measured the
1.21 Label each of the following as either a physical process or a mass of the sample, finding that it weighed 38.5 g. She
1density 0.785 g>mL2 or toluene 1density 0.866>mL2. What

chemical process: (a) rusting of a metal can, (b) boiling a cup knew that the substance had to be either isopropyl alcohol
of water, (c) pulverizing an aspirin, (d) digesting a candy bar,
(e) exploding of nitroglyerin. are the calculated density and the probable identity of the
substance? (b) An experiment requires 45.0 g of ethylene gly-
1.22 A match is lit and held under a cold piece of metal. The
col, a liquid whose density is 1.114 g>mL. Rather than weigh
following observations are made: (a) The match burns.
the sample on a balance, a chemist chooses to dispense the
(b) The metal gets warmer. (c) Water condenses on the metal.
liquid using a graduated cylinder. What volume of the liquid
(d) Soot (carbon) is deposited on the metal. Which of these
should he use? (c) Is a graduated cylinder such as that shown
occurrences are due to physical changes, and which are due to
in Figure 1.19 likely to afford the accuracy of measurement
chemical changes?
needed? (d) A cubic piece of metal measures 5.00 cm on each
1.23 Suggest a method of separating each of the following mixtures edge. If the metal is nickel, whose density is 8.90 g>cm3, what
into two components: (a) sugar and sand, (b) oil and vinegar. is the mass of the cube?
1.24 Three beakers contain clear, colorless liquids. One beaker 1.32 (a) After the label fell off a bottle containing a clear liquid be-
contains pure water, another contains salt water, and an- lieved to be benzene, a chemist measured the density of the
other contains sugar water. How can you tell which beaker is liquid to verify its identity. A 25.0-mL portion of the liquid
which? (No tasting allowed!) had a mass of 21.95 g. A chemistry handbook lists the den-
sity of benzene at 15 °C as 0.8787 g>mL. Is the calculated
density in agreement with the tabulated value? (b) An experi-
Units and Measurement (Section 1.4)
ment requires 15.0 g of cyclohexane, whose density at 25 °C is
1.25 What exponential notation do the following abbreviations 0.7781 g>mL. What volume of cyclohexane should be used?
represent? (a) d, (b) c, (c) f, (d) m, (e) M, (f) k, (g) n, (h) m, (i) p. (c) A spherical ball of lead has a diameter of 5.0 cm. What is
(The volume of a sphere is 14>32pr3, where r is the radius.)

the mass of the sphere if lead has a density of 11.34 g>cm3?
1.26 Use appropriate metric prefixes to write the following mea-
surements without use of exponents: (a) 2.3 * 10-10 L,
carbon dioxide 1CO22 was emitted worldwide due to fossil

( b ) 4.7 * 10-6 g, ( c ) 1.85 * 10-12 m, ( d ) 16.7 * 106 s, 1.33 In the year 2011, an estimated amount of 35 billion tons of
(e) 15.7 * 103 g, (f) 1.34 * 10-3 m, (g) 1.84 * 102 cm.
fuel combustion and cement production. Express this mass of
1.27 Make the following conversions: (a) 72 °F to °C, (b) 216.7 °C
CO2 in grams without exponential notation, using an appro-
to °F, (c) 233 °C to K, (d) 315 K to °F, (e) 2500 °F to K,
priate metric prefix.
(f) 0 K to °F.
1.34 Silicon for computer chips is grown in large cylinders called (a) 320.5 - 16104.5>2.32
“boules” that are 300 mm in diameter and 2 m in length, as (b) 31285.3 * 1052 - 11.200 * 10324 * 2.8954
(c) 10.0045 * 20,000.02 + 12813 * 122
shown. The density of silicon is 2.33 g>cm3. Silicon wafers for
(d) 863 * 31255 - 13.45 * 10824

making integrated circuits are sliced from a 2.0 m boule and
are typically 0.75 mm thick and 300 mm in diameter. (a) How
many wafers can be cut from a single boule? (b) What is the 1.43 You weigh an object on a balance and read the mass in grams
mass of a silicon wafer? (The volume of a cylinder is given by according to the picture. How many significant figures are in
pr2h, where r is the radius and h is its height.) this measurement?
Diamond blade
0.75 mm
thickness
Si boule
300 mm
diameter
Cut wafers
2m
1.44 You have a graduated cylinder that contains a liquid (see pho-
Uncertainty in Measurement (Section 1.5) tograph). Write the volume of the liquid, in milliliters, using
the proper number of significant figures.
1.35 Indicate which of the following are exact numbers: (a) the
mass of a 3 by 5–inch index card, (b) the number of ounces
in a pound, (c) the volume of a cup of Seattle’s Best coffee,
(d) the number of inches in a mile, (e) the number of micro-
seconds in a week, (f) the number of pages in this book.
1.36 Indicate which of the following are exact numbers: (a) the
mass of a 32-oz can of coffee, (b) the number of students in
your chemistry class, (c) the temperature of the surface of the
Sun, (d) the mass of a postage stamp, (e) the number of mil-
liliters in a cubic meter of water, (f) the average height of NBA
basketball players.
1.37 What is the number of significant figures in each of the fol-
lowing measured quantities? (a) 601 kg, (b) 0.054 s, (c) 6.3050
cm, (d) 0.0105 L, (e) 7.0500 * 10-3 m3, (f) 400 g. Dimensional Analysis (Section 1.6)
1.38 Indicate the number of significant figures in each of the
1.45 Using your knowledge of metric units, English units, and the
following measured quantities: (a) 3.774 km, (b) 205 m2,
information on the back inside cover, write down the conver-
(c) 1.700 cm, (d) 350.00 K, (e) 307.080 g, (f) 1.3 * 103 m>s.
sion factors needed to convert (a) mm to nm, (b) mg to kg,
1.39 Round each of the following numbers to four significant fig- (c) km to ft, (d) in.3 to cm3.
ures and express the result in standard exponential notation:
1.46 Using your knowledge of metric units, English units, and the
(a) 102.53070, (b) 656.980, (c) 0.008543210, (d) 0.000257870,
information on the back inside cover, write down the conver-
(e) - 0.0357202.
sion factors needed to convert (a) mm to mm, (b) ms to ns,
1.40 (a) The diameter of Earth at the equator is 7926.381 mi. Round (c) mi to km, (d) ft3 to L.
this number to three significant figures and express it in stan-
1.47 (a) A bumblebee flies with a ground speed of 15.2 m/s. Cal-
dard exponential notation. (b) The circumference of Earth
culate its speed in km/hr. (b) The lung capacity of the blue
through the poles is 40,008 km. Round this number to four sig-
whale is 5.0 * 103 L. Convert this volume into gallons.
nificant figures and express it in standard exponential notation.
(c) The Statue of Liberty is 151 ft tall. Calculate its height in
1.41 Carry out the following operations and express the answers meters. (d) Bamboo can grow up to 60.0 cm/day. Convert this
with the appropriate number of significant figures. growth rate into inches per hour.
(a) 14.3505 + 2.65 1.48 (a) The speed of light in a vacuum is 2.998 * 108 m>s.
(b) 952.7 - 140.7389 Calculate its speed in miles per hour. (b) The Sears Tower in
(c) 13.29 * 104210.25012 Chicago is 1454 ft tall. Calculate its height in meters. (c) The
Vehicle Assembly Building at the Kennedy Space Center in
(d) 0.0588/0.677
Florida has a volume of 3,666,500 m3. Convert this volume to
1.42 Carry out the following operations and express the answer liters and express the result in standard exponential notation.
with the appropriate number of significant figures. (d) An individual suffering from a high cholesterol level in her
additional exercises 37
blood has 242 mg of cholesterol per 100 mL of blood. If the 1.54 (a) If an electric car is capable of going 225 km on a single charge,
total blood volume of the individual is 5.2 L, how many grams how many charges will it need to travel from Seattle, Washing-
of total blood cholesterol does the individual’s body contain? ton, to San Diego, California, a distance of 1257 mi, assuming
1.49 The inside dimension of a box that is cubic is 24.8 cm on each that the trip begins with a full charge? (b) If a migrating loon flies
edge with an uncertainty of 0.2 cm. What is the volume of the at an average speed of 14 m/s, what is its average speed in mi/hr?
box? What do you estimate to be the uncertainty in the calcu- (c) What is the engine piston displacement in liters of an engine
lated volume? whose displacement is listed as 450 in.3? (d) In March 1989 the
Exxon Valdez ran aground and spilled 240,000 barrels of crude
1.50 The distance from Grand Rapids, Michigan, to Detroit is listed petroleum off the coast of Alaska. One barrel of petroleum is
in a road atlas as 153 miles. Describe some of the factors that equal to 42 gal. How many liters of petroleum were spilled?
contribute to the uncertainty in this number. To make the num-
1.55 The density of air at ordinary atmospheric pressure and 25 °C
ber more precise, what would you need to specify and measure?
is 1.19 g>L. What is the mass, in kilograms, of the air in a
1.51 Perform the following conversions: (a) 5.00 days to s, room that measures 14.5 ft * 16.5 ft * 8.0 ft?
(b) 0.0550 mi to m, (c) $1.89/gal to dollars per liter,
1.56 The concentration of carbon monoxide in an urban apart-
(d) 0.510 in./ms to km/hr, (e) 22.50 gal/min to L/s,
ment is 48 mg>m3. What mass of carbon monoxide in grams
(f) 0.02500 ft3 to cm3.
is present in a room measuring 10.6 ft * 14.8 ft * 20.5 ft?
1.52 Carry out the following conversions: (a) 0.105 in. to mm, 1.57 Gold can be hammered into extremely thin sheets called gold
(b) 0.650 qt to mL, (c) 8.75 mm>s to km>hr, (d) 1.955 m3 to yd3, leaf. An architect wants to cover a 100 ft * 82 ft ceiling with
(e) $3.99/lb to dollars per kg, (f) 8.75 lb>ft3 to g>mL. gold leaf that is five–millionths of an inch thick. The density
11 troy ounce = 31.1034768 g2. How much will it cost the

1.53 (a) How many liters of wine can be held in a wine barrel of gold is 19.32 g>cm3, and gold costs $1654 per troy ounce
whose capacity is 31 gal? (b) The recommended adult dose of
Elixophyllin®, a drug used to treat asthma, is 6 mg/kg of body architect to buy the necessary gold?
mass. Calculate the dose in milligrams for a 185-lb person. 1.58 A copper refinery produces a copper ingot weighing 150 lb.
(c) If an automobile is able to travel 400 km on 47.3 L of gaso- If the copper is drawn into wire whose diameter is 7.50 mm,
line, what is the gas mileage in miles per gallon? (d) When how many feet of copper can be obtained from the ingot? The
beans yields 50 cups of coffee 14 cups = 1 qt2. How many kg

the coffee is brewed according to directions, a pound of coffee density of copper is 8.94 g>cm3. (Assume that the wire is a
cylinder whose volume V = pr2h, where r is its radius and h
of coffee are required to produce 200 cups of coffee? is its height or length.)
Additional Exercises
1.59 (a) Classify each of the following as a pure substance, a solu-
tion, or a heterogeneous mixture: a gold coin, a cup of coffee,
a wood plank. (b) What ambiguities are there in answering
part (a) from the descriptions given?
1.60 (a) What is the difference between a hypothesis and a theory?
(b) Explain the difference between a theory and a scientific
law. Which addresses how matter behaves, and which ad-
dresses why it behaves that way?
1.61 A sample of ascorbic acid (vitamin C) is synthesized in the
laboratory. It contains 1.50 g of carbon and 2.00 g of oxy-
gen. Another sample of ascorbic acid isolated from citrus
fruits contains 6.35 g of carbon. How many grams of oxygen
does it contain? Which law are you assuming in answering
this question?
1.62 Ethyl chloride is sold as a liquid (see photo) under pres-
sure for use as a local skin anesthetic. Ethyl chloride boils at
12 °C at atmospheric pressure. When the liquid is sprayed
onto the skin, it boils off, cooling and numbing the skin as it
vaporizes. (a) What changes of state are involved in this use
1.63 Two students determine the percentage of lead in a sample
of ethyl chloride? (b) What is the boiling point of ethyl chlo-
as a laboratory exercise. The true percentage is 22.52%. The
ride in degrees Fahrenheit? (c) The bottle shown contains
students’ results for three determinations are as follows:
103.5 mL of ethyl chloride. The density of ethyl chloride at
25 °C is 0.765 g>cm3. What is the mass of ethyl chloride in (1) 22.52, 22.48, 22.54
the bottle? (2) 22.64, 22.58, 22.62
units of O, using the freezing point 113 °C2 and boiling point
(a) Calculate the average percentage for each set of data and 1.72 Suppose you decide to define your own temperature scale with
1360 °C2 of oleic acid, the main component of olive oil. If you
state which set is the more accurate based on the average.
(b) Precision can be judged by examining the average of the
deviations from the average value for that data set. (Calculate set the freezing point of oleic acid as 0 °O and the boiling point as
the average value for each data set; then calculate the average 100 °O, what is the freezing point of water on this new scale?
value of the absolute deviations of each measurement from 1.73 The liquid substances mercury 1density = 13.6 g>mL2,
the average.) Which set is more precise? water 11.00 g>mL2, and cyclohexane 10.778 g>mL2 do not
1.64 Is the use of significant figures in each of the following form a solution when mixed but separate in distinct layers.
statements appropriate? Why or why not? (a) Apple sold Sketch how the liquids would position themselves in a test tube.
22,727,000 iPods during the last three months of 2008. 1.74 Two spheres of equal volume are placed on the scales as
(b) New York City receives 49.7 inches of rain, on average, shown. Which one is more dense?
per year. (c) In the United States, 0.621% of the population
has the surname Brown. (d) You calculate your grade point
average to be 3.87562.
1.65 What type of quantity (for example, length, volume, density)
do the following units indicate? (a) mL, (b) cm2, (c) mm3,
(d) mg/L, (e) ps, (f) nm, (g) K.
1.66 Give the derived SI units for each of the following quantities
in base SI units:
(a) acceleration = distance>time2
(b) force = mass * acceleration
(c) work = force * distance
(d) pressure = force>area
(e) power = work>time 1.75 Water has a density of 0.997 g>cm3 at 25 °C; ice has a density
(f) velocity = distance>time of 0.917 g>cm3 at - 10 °C. (a) If a soft-drink bottle whose vol-
(g) energy = mass * 1velocity22 ume is 1.50 L is completely filled with water and then frozen
to -10 °C, what volume does the ice occupy? (b) Can the ice
1.67 The distance from Earth to the Moon is approximately
be contained within the bottle?
240,000 mi. (a) What is this distance in meters? (b) The per-
egrine falcon has been measured as traveling up to 350 km/ 1.76 A 32.65-g sample of a solid is placed in a flask. Toluene, in
hr in a dive. If this falcon could fly to the Moon at this speed, which the solid is insoluble, is added to the flask so that the
how many seconds would it take? (c) The speed of light is total volume of solid and liquid together is 50.00 mL. The
3.00 * 108 m>s. How long does it take for light to travel from solid and toluene together weigh 58.58 g. The density of
Earth to the Moon and back again? (d) Earth travels around toluene at the temperature of the experiment is 0.864 g>mL.
the Sun at an average speed of 29.783 km>s. Convert this What is the density of the solid?
speed to miles per hour. 1.77 A thief plans to steal a gold sphere with a radius of 28.9 cm
1.68 Which of the following would you characterize as a pure or from a museum. If the gold has a density of 19.3 g>cm3, what
is V = 14>32pr3.4 Is the thief likely to be able to walk off

nearly pure substance? (a) baking powder; (b) lemon juice; is the mass of the sphere in pounds? [The volume of a sphere
(c) propane gas, used in outdoor gas grills; (d) aluminum foil;
(e) ibuprofen; (f) bourbon whiskey; (g) helium gas; (h) clear with the gold sphere unassisted?
water pumped from a deep aquifer. 1.78 Automobile batteries contain sulfuric acid, which is com-
1.69 The U.S. quarter has a mass of 5.67 g and is approximately monly referred to as “battery acid.” Calculate the number of
1.55 mm thick. (a) How many quarters would have to grams of sulfuric acid in 1.00 gal of battery acid if the solution
be stacked to reach 575 ft, the height of the Washing- has a density of 1.28 g/mL and is 38.1% sulfuric acid by mass.
ton Monument? (b) How much would this stack weigh? 1.79 A 40-lb container of peat moss measures 14 * 20 * 30 in. A
(c) How much money would this stack contain? (d) The U.S. 40-lb container of topsoil has a volume of 1.9 gal. (a) Calculate
National Debt Clock showed the outstanding public debt the average densities of peat moss and topsoil in units of g>cm3.
to be $16,213,166,914,811 on October 28, 2012. How many Would it be correct to say that peat moss is “lighter” than topsoil?
stacks like the one described would be necessary to pay off Explain. (b) How many bags of peat moss are needed to cover an
this debt? area measuring 15.0 ft * 20.0 ft to a depth of 3.0 in.?
1.70 In the United States, water used for irrigation is measured in 1.80 A package of aluminum foil contains 50 ft2 of foil, which
acre-feet. An acre-foot of water covers an acre to a depth of weighs approximately 8.0 oz. Aluminum has a density of
exactly 1 ft. An acre is 4840 yd2. An acre-foot is enough water 2.70 g>cm3. What is the approximate thickness of the foil
to supply two typical households for 1.00 yr. (a) If desalinated in millimeters?
water costs $1950 per acre-foot, how much does desalinated 1.81 The total rate at which power used by humans worldwide
water cost per liter? (b) How much would it cost one house- is approximately 15 TW (terawatts). The solar flux aver-
hold per day if it were the only source of water? aged over the sunlit half of Earth is 680 W>m2. (assuming
1.71 By using estimation techniques, determine which of the follow- no clouds). The area of Earth’s disc as seen from the sun
a 5-kg bag of sugar, or 1 gal of water 1density = 1.0 g>mL2.

ing is the heaviest and which is the lightest: a 5-lb bag of potatoes, is 1.28 * 1014 m2. The surface area of Earth is approxi-
mately 197,000,000 square miles. How much of Earth’s
additional exercises 39
surface would we need to cover with solar energy collectors 1.86 Judge the following statements as true or false. If you believe a
to power the planet for use by all humans? Assume that the statement to be false, provide a corrected version.
solar energy collectors can convert only 10% of the available (a) Air and water are both elements.
sunlight into useful power.
(b) All mixtures contain at least one element and one
1.82 In 2005, J. Robin Warren and Barry J. Marshall shared the compound.
Nobel Prize in Medicine for discovery of the bacterium
(c) Compounds can be decomposed into two or more other
Helicobacter pylori, and for establishing experimental proof
substances; elements cannot.
that it plays a major role in gastritis and peptic ulcer disease.
The story began when Warren, a pathologist, noticed that (d) Elements can exist in any of the three states of matter.
bacilli were associated with the tissues taken from patients (e) When yellow stains in a kitchen sink are treated with
suffering from ulcers. Look up the history of this case and bleach water, the disappearance of the stains is due to a
describe Warren’s first hypothesis. What sorts of evidence did physical change.
it take to create a credible theory based on it? (f) A hypothesis is more weakly supported by experimental
1.83 A 25.0-cm long cylindrical glass tube, sealed at one end, is evidence than a theory.
filled with ethanol. The mass of ethanol needed to fill the tube (g) The number 0.0033 has more significant figures than
is found to be 45.23 g. The density of ethanol is 0.789 g/mL. 0.033.
Calculate the inner diameter of the tube in centimeters.
(h) Conversion factors used in converting units always have
1.84 Gold is alloyed (mixed) with other metals to increase its hard- a numerical value of one.
ness in making jewelry. (a) Consider a piece of gold jewelry
that weighs 9.85 g and has a volume of 0.675 cm3. The jew- (i) Compounds always contain at least two different
elry contains only gold and silver, which have densities of elements.
19.3 and 10.5 g>cm3, respectively. If the total volume of the 1.87 You are assigned the task of separating a desired granular ma-
jewelry is the sum of the volumes of the gold and silver that terial with a density of 3.62 g>cm3 from an undesired granular
it contains, calculate the percentage of gold (by mass) in the material that has a density of 2.04 g>cm3. You want to do this
jewelry. (b) The relative amount of gold in an alloy is com- by shaking the mixture in a liquid in which the heavier mate-
monly expressed in units of carats. Pure gold is 24 carat, and rial will fall to the bottom and the lighter material will float.
the percentage of gold in an alloy is given as a percentage of A solid will float on any liquid that is more dense. Using an
this value. For example, an alloy that is 50% gold is 12 carat. Internet-based source or a handbook of chemistry, find the
State the purity of the gold jewelry in carats. densities of the following substances: carbon tetrachloride,
1.85 Paper chromatography is a simple but reliable method for sep- hexane, benzene, and diiodomethane. Which of these liquids
arating a mixture into its constituent substances. You have a will serve your purpose, assuming no chemical interaction be-
mixture of two vegetable dyes, one red and one blue, that you tween the liquid and the solids?
are trying to separate. You try two different chromatography 1.88 In 2009, a team from Northwestern University and Western
procedures and achieve the separations shown in the figure. Washington University reported the preparation of a new
Which procedure worked better? Can you suggest a method “spongy” material composed of nickel, molybdenum, and
to quantify how good or poor the separation was? sulfur that excels at removing mercury from water. The den-
sity of this new material is 0.20 g>cm3, and its surface area is
1242 m2 per gram of material. (a) Calculate the volume of a
10.0-mg sample of this material. (b) Calculate the surface area
for a 10.0-mg sample of this material. (c) A 10.0-mL sample
of contaminated water had 7.748 mg of mercury in it. After
treatment with 10.0 mg of the new spongy material, 0.001 mg
of mercury remained in the contaminated water. What
percentage of the mercury was removed from the water?
(d) What is the final mass of the spongy material after the
exposure to mercury?
APPENDIX
Mathematical Operations
A
A.1 | Exponential Notation
The numbers used in chemistry are often either extremely large or extremely small.
Such numbers are conveniently expressed in the form
N * 10n
where N is a number between 1 and 10, and n is the exponent. Some examples of this
exponential notation, which is also called scientific notation, follow.
1,200,000 is 1.2 * 106 (read “one point two multi ten to the sixth power”)
0.000604 is 6.04 * 10-4 (read “six point zero four times ten to the negative fourth
power”)
A positive exponent, as in the first example, tells us how many times a number
must be multiplied by 10 to give the long form of the number:
1.2 * 106 = 1.2 * 10 * 10 * 10 * 10 * 10 * 10 1six tens2
= 1,200,000
It is also convenient to think of the positive exponent as the number of places the deci-
mal point must be moved to the left to obtain a number greater than 1 and less than 10.
For example, if we begin with 3450 and move the decimal point three places to the left,
we end up with 3.45 * 103.
In a related fashion, a negative exponent tells us how many times we must divide a
number by 10 to give the long form of the number.
6.04
6.04 * 10-4 = = 0.000604
10 * 10 * 10 * 10
It is convenient to think of the negative exponent as the number of places the decimal
point must be moved to the right to obtain a number greater than 1 but less than 10.
For example, if we begin with 0.0048 and move the decimal point three places to the
right, we end up with 4.8 * 10-3.
In the system of exponential notation, with each shift of the decimal point one
place to the right, the exponent decreases by 1:
4.8 * 10-3 = 48 * 10-4
Similarly, with each shift of the decimal point one place to the left, the exponent
increases by 1:
4.8 * 10-3 = 0.48 * 10-2
Many scientific calculators have a key labeled EXP or EE, which is used to enter
numbers in exponential notation. To enter the number 5.8 * 103 on such a calculator,
the key sequence is
5 # 8 EXP (or EE ) 3
On some calculators the display will show 5.8, then a space, followed by 03, the
exponent. On other calculators, a small 10 is shown with an exponent 3.
1092
A.1 Exponential Notation 1093
To enter a negative exponent, use the key labeled + > - . For example, to enter the
number 8.6 * 10-5, the key sequence is
8 # 6 EXP + > - 5
When entering a number in exponential notation, do not key in the 10 if you use the EXP
or EE button.
In working with exponents, it is important to recall that 100 = 1. The following
rules are useful for carrying exponents through calculations.
1. Addition and Subtraction In order to add or subtract numbers expressed in expo-
nential notation, the powers of 10 must be the same.
15.22 * 1042 + 13.21 * 1022 = 1522 * 1022 + 13.21 * 1022

= 525 * 102 13 significant figures2
= 5.25 * 104
16.25 * 10-22 - 15.77 * 10-32 = 16.25 * 10-22 - 10.577 * 10-22
= 5.67 * 10-2 13 significant figures2
When you use a calculator to add or subtract, you need not be concerned with hav-
ing numbers with the same exponents because the calculator automatically takes
care of this matter.
2. Multiplication and Division When numbers expressed in exponential notation
are multiplied, the exponents are added; when numbers expressed in exponential
notation are divided, the exponent of the denominator is subtracted from the
exponent of the numerator.
15.4 * 102212.1 * 1032 = 15.4212.12 * 102+3

= 11 * 105
= 1.1 * 106
11.2 * 105213.22 * 10-32 = 11.2213.222 * 105+1-32 = 3.9 * 102
3.2 * 105 3.2

2 = * 105-2 = 0.49 * 103 = 4.9 * 102
6.5 * 10 6.5
5.7 * 107 5.7
= * 107-1-22 = 0.67 * 109 = 6.7 * 108
8.5 * 10-2 8.5
3. Powers and Roots When numbers expressed in exponential notation are raised
to a power, the exponents are multiplied by the power. When the roots of num-
bers expressed in exponential notation are taken, the exponents are divided by
the root.
11.2 * 10523 = 11.223 * 105*3
= 1.7 * 1015
3 6 3
22.5 * 10 = 2 2.5 * 106>3
= 1.3 * 102
Scientiic calculators usually have keys labeled x2 and 2x for squaring and taking
the square root of a number, respectively. To take higher powers or roots, many
x
calculators have yx and 2y (or INV yx) keys. For example, to perform the opera-
3
tion 27.5 * 10-4 on such a calculator, you would key in 7.5 * 10-4, press the
x
2y key (or the INV and then the yx keys), enter the root, 3, and inally press =.
he result is 9.1 * 10-2.
1094 APPENDIX A Mathematical Operations
SAMPLE
EXERCISE 1 Using Exponential Notation
Perform each of the following operations, using your calculator where possible:
(a) Write the number 0.0054 in standard exponential notation. (c) Performing this operation longhand, we have
(b) 15.0 * 10-22 + 14.7 * 10-32 15.98 * 2.772 * 1012-5 = 16.6 * 107 = 1.66 * 108
(c) 15.98 * 1012212.77 * 10-52
4 (d)To perform this operation on a calculator, we enter the number,
x
(d) 21.75 * 10-12 press the 1y key (or the INV and yx keys), enter 4, and press the
= key. The result is 1.15 * 10-3.
SOLUTION
(a) Because we move the decimal point three places to the right to Practice Exercise
convert 0.0054 to 5.4, the exponent is - 3: Perform the following operations:
5.4 * 10-3 (a) Write 67,000 in exponential notation, showing two significant
Scientific calculators are generally able to convert numbers to ex- figures.
ponential notation using one or two keystrokes; frequently “SCI” (b) 13.378 * 10-32 - 14.97 * 10-52
for “scientific notation” will convert a number into exponential no- (c) 11.84 * 1015217.45 * 10-22
tation. Consult your instruction manual to see how this operation (d) 16.67 * 10-823
is accomplished on your calculator.
(b) To add these numbers longhand, we must convert them to the
same exponent.
15.0 * 10-22 + 10.47 * 10-22 = 15.0 + 0.472 * 10-2 = 5.5 * 10-2
A.2 | Logarithms
Common Logarithms
The common, or base-10, logarithm (abbreviated log) of any number is the power to
which 10 must be raised to equal the number. For example, the common logarithm of
1000 (written log 1000) is 3 because raising 10 to the third power gives 1000.
103 = 1000, therefore, log 1000 = 3
Further examples are
log 105 = 5
log 1 = 0 Remember that 100 = 1
log 10-2 = -2
In these examples the common logarithm can be obtained by inspection. However, it is
not possible to obtain the logarithm of a number such as 31.25 by inspection. The loga-
rithm of 31.25 is the number x that satisfies the following relationship:
10x = 31.25
Most electronic calculators have a key labeled LOG that can be used to obtain loga-
rithms. For example, on many calculators we obtain the value of log 31.25 by entering
31.25 and pressing the LOG key. We obtain the following result:
log 31.25 = 1.4949
Notice that 31.25 is greater than 10 11012 and less than 100 11022. The value for log
31.25 is accordingly between log 10 and log 100, that is, between 1 and 2.
Significant Figures and Common Logarithms

For the common logarithm of a measured quantity, the number of digits after the deci-
mal point equals the number of significant figures in the original number. For example,
if 23.5 is a measured quantity (three significant figures), then log 23.5 = 1.371 (three
significant figures after the decimal point).
ANSWERS TO SELECTED EXERCISES
Chapter 1 is not correct to say that peat is “lighter” than topsoil. Volumes must
be specified in order to compare masses. (b) Buy 16 bags of peat (more
1.1 (a) Pure element: i (b) mixture of elements: v, vi (c) pure compound: than 15 are needed). (Results to 1 significant figure are not meaning-
iv (d) mixture of an element and a compound: ii, iii 1.3 This kind of ful.) 1.83 The inner diameter of the tube is 1.71 cm. 1.85 The separa-
separation based on solubility differences is called extraction. The in- tion is successful if two distinct spots are seen on the paper. To quantify
soluble grounds are then separated from the coffee solution by filtra- the characteristics of the separation, calculate a reference value for
tion. 1.5 (a) The aluminum sphere is lightest, then nickel, then silver. each spot: distance traveled by spot/distance traveled by solvent. If the
(b) The platinum sphere is smallest, then gold, then lead. 1.7 (a) 7.5 cm; values for the two spots are fairly different, the separation is successful.
two significant figures (sig figs) (b) 72 mi>hr (inner scale, two sig- 1.88 (a) Volume = 0.050 mL (b) surface area = 12.4 m2 (c) 99.99%
nificant figures) or 115 km>hr (outer scale, three significant figures) of the mercury was removed. (c) The spongy material weighs 17.7 mg
1.9 Arrange the conversion factor so that the given unit cancels and after exposure to mercury.
the desired unit is in the correct position. 1.11 464 jelly beans. The
mass of an average bean has 2 decimal places and 3 sig figs. The num-
ber of beans then has 3 sig figs, by the rules for multiplication and
Chapter 2
division. 1.13 (a) Heterogeneous mixture (b) homogeneous mixture 2.1 (a) The path of the charged particle bends because the particle is
(heterogeneous if there are undissolved particles) (c) pure substance repelled by the negatively charged plate and attracted to the positively
(d) pure substance. 1.15 (a) S (b) Au (c) K (d) Cl (e) Cu (f) uranium charged plate. (b) 1- 2 (c) increase (d) decrease 2.4 The particle is
(g) nickel (h) sodium (i) aluminum (j) silicon 1.17 C is a compound; an ion. 32 2-
16S 2.6 Formula: IF5; name: iodine pentafluoride; the com-
it contains both carbon and oxygen. A is a compound; it contains pound is molecular. 2.8 Only Ca1NO322, calcium nitrate, is consistent
at least carbon and oxygen. B is not defined by the data given; it is with the diagram. 2.10 (a) In the presence of an electric field, there
probably also a compound because few elements exist as white solids. is electrostatic attraction between the negatively charged oil drops
1.19 Physical properties: silvery white; lustrous; melting point and the positively charged plate as well as electrostatic repulsion be-
= 649 °C, boiling point = 1105 °C; density at 20 °C = 1.738 g>cm3; tween the negatively charged oil drops and the negatively charged
pounded into sheets; drawn into wires; good conductor. Chemical plate. These electrostatic forces oppose the force of gravity and change
properties: burns in air; reacts with Cl2. 1.21 (a) Chemical (b) physical the rate of fall of the drops. (b) Each individual drop has a differ-
(c) physical (d) chemical (e) chemical 1.23 (a) Add water to dissolve ent number of electrons associated with it. If the combined electro-
the sugar; filter this mixture, collecting the sand on the filter paper and static forces are greater than the force of gravity, the drop moves up.
the sugar water in the flask. Evaporate water from the flask to recover 2.11 Postulate 4 of the atomic theory states that the relative number
solid sugar. (b) Allow the mixture to settle so that there are two dis- and kinds of atoms in a compound are constant, regardless of the
tinct layers. Carefully pour off most of the top oil layer. After the lay- source. Therefore, 1.0 g of pure water should always contain the same
ers reform, use a dropper to remove any remaining oil. Vinegar is in relative amounts of hydrogen and oxygen, no matter where or how
the original vessel and oil is in a second container. 1.25 (a) 1 * 10-1 the sample is obtained. 2.13 (a) 0.5711 g O>1 g N; 1.142 g O>1 g N;
(b) 1 * 10-2 (c) 1 * 10-15 (d) 1 * 10-6 (e) 1 * 106 (f) 1 * 103 2.284 g O>1 g N; 2.855 g O>1 g N (b) The numbers in part (a) obey
(g) 1 * 10-9 (h) 1 * 10-3 (i) 1 * 10-12 1.27 (a) 22 °C (b) 422.1 °F the law of multiple proportions. Multiple proportions arise because
(c) 506 K (d) 107 °F (e) 1644 K (f) - 459.67 °F 1.29 (a) 1.62 g>mL. atoms are the indivisible entities combining, as stated in Dalton’s
Tetrachloroethylene, 1.62 g>mL, is more dense than water, 1.00 g>mL; atomic theory. 2.15 (i) Electric and magnetic fields deflected the
tetrachloroethylene will sink rather than float on water. (b) 11.7 g rays in the same way they would deflect negatively charged particles.
1.31 (a) Calculated density = 0.86 g>mL. The substance is prob- (ii) A metal plate exposed to cathode rays acquired a negative charge.
ably toluene, density = 0.866 g>mL. (b) 40.4 mL ethylene glycol 2.17 (a) Most of the volume of an atom is empty space in which elec-
(c) 1.11 * 103 g nickel 1.33 32 Pg 1.35 Exact: (b), (d), and (f) trons move. Most alpha particles passed through this space. (b) The
1.37 (a) 3 (b) 2 (c) 5 (d) 3 (e) 5 (f) 1 1.39 (a) 1.025 * 102 (b) 6.570 * 102 few alpha particles that hit the massive, positively charged gold nuclei
(c) 8.543 * 10-3 (d) 2.579 * 10-4 (e) - 3.572 * 10-2 1.41 (a) 17.00 were strongly repelled and deflected back in the direction they came
(b) 812.0 (c) 8.23 * 103 (d) 8.69 * 10-2 1.43 5 significant figures from. (c) Because the Be nuclei have a smaller volume and a smaller
1 * 10-3 m 1 nm 1 * 10-3 g 1 kg positive charge than the Au nuclei, fewer alpha particles will be scat-
1.45 (a) * (b) * tered and fewer will be strongly back scattered. 2.19 (a) 0.135 nm;
1 mm 1 * 10 m -9
1 mg 1000 g
1.35 * 102 or 135 pm (b) 3.70 * 106 Au atoms (c) 1.03 * 10-23 cm3
1000 m 1 cm 1 in. 1 ft 12.5423 cm3 2.21 (a) Proton, neutron, electron (b) proton = 1 +, neutron = 0,
(c) * * * (d)
1 km 1 * 10-2 m 2.54 cm 12 in. 13 in.3 electron = 1 - (c) The neutron is most massive. (The neutron and
3
1.47 (a) 54.7 km>hr (b) 1.3 * 10 gal (c) 46.0 m (d) 0.984 in/hr proton have very similar masses.) (d) The electron is least massive.
1.49 The volume of the box is 1.52 * 104 cm3.The uncertainty in 2.23 (a) Not isotopes (b) isotopes (c) isotopes 2.25 (a) Atomic number
the calculated volume is {0.4 * 104 cm3 cm. 1.51 (a) 4.32 * 105 s is the number of protons in the nucleus of an atom. Mass number is the
(b) 88.5 m (c) +0.499>L (d) 46.6 km>hr (e) 1.420 L>s (f) 707.9 cm3 total number of nuclear particles, protons plus neutrons, in an atom.
1.53 (a) 1.2 * 102 L (b) 5 * 102 mg (c) 19.9 mi>gal (2 * 101 mi>gal (b) mass number 2.27 (a) 40Ar: 18 p, 22 n, 18 e (b) 65Zn: 30 p, 35 n,
for 1 significant figure) (d) 1.81 kg 1.55 64 kg air 1.57 $6 * 104 30 e (c) 70Ga: 31 p, 39 n, 31 e (d) 80Br: 35 p, 45 n, 35 e (e) 184W: 74 p,
1.61 8.47 g O; the law of constant composition 1.63 (a) Set I, 22.51; 110 n, 74 e (f) 243Am: 95 p, 148 n, 95e
set II, 22.61. Based on the average, set I is more accurate. (b) The 2.29
average deviation for both set I and set II is 0.02. The two sets dis- Symbol 79
Br 55
Mn 112
Cd 222
Rn 207
Pb
play the same precision. 1.65 (a) Volume (b) area (c) volume
(d) density (e) time (f) length (g) temperature 1.68 Substances (c), (d), Protons 35 25 48 86 82
(e), (g) and (h) are pure or nearly pure. 1.69 (a) 1.13 * 105 quar- Neutrons 44 30 64 136 125
ters (b) 6.41 * 105 g (c) $2.83 * 104 (d) 5.74 * 108 stacks 1.73 The
most dense liquid, Hg, will sink; the least dense, cyclohexane, will Electrons 35 25 48 86 82
float; H2O will be in the middle. 1.76 density of solid = 1.63 g>mL Mass no. 79 55 112 222 207
1.79 (a) Density of peat = 0.13 g>cm3, density of soil = 2.5 g>cm3. It A-1
ANSWERS TO SELECTED PRACTICE EXERCISES
Chapter 1 Sample Exercise 2.7
Practice Exercise 2: 34 protons, 45 neutrons, and 36 electrons
Sample Exercise 1.1
Practice Exercise 2: It is a compound because it has constant composi- Sample Exercise 2.8
tion and can be separated into several elements. Practice Exercise 2: (a) 3 + , (b) 1 -
Sample Exercise 1.2 Sample Exercise 2.9
Practice Exercise 2: (a) 1012 pm, (b) 6.0 km, (c) 4.22 * 10-3 g, Practice Exercise 2: (a) Rb is from group 1, and readily loses one electron
(d) 0.00422 g. to attain the electron configuration of the nearest noble gas element, Kr.
(b) Nitrogen and the halogens are all nonmetallic elements, which
Sample Exercise 1.3 form molecular compounds with one another. (c) Krypton, Kr, is
Practice Exercise 2: (a) 261.7 K, (b) 11.3 °F a noble gas element and is chemically inactive except under special
Sample Exercise 1.4 conditions. (d) Na and K are both from group 1 and adjacent to
Practice Exercise 2: (a) 8.96 g>cm3, (b) 19.0 mL, (c) 340 g. one another in the periodic table. They would be expected to behave
very similarly. (e) Calcium is an active metal and readily loses two
Sample Exercise 1.5 electrons to attain the noble gas configuration of Ar.
Practice Exercise 2: Five, as in the measurement 24.995 g, the uncer-
tainty being in the third decimal place. Sample Exercise 2.10
Practice Exercise 2: (a) Na3PO4, (b) ZnSO4, (c) Fe21CO323
Sample Exercise 1.6
Practice Exercise 2: No. The number of feet in a mile is a defined quan- Sample Exercise 2.11
tity and is therefore exact, but the distance represented by one foot is Practice Exercise 2: BrO- and BrO2-
not exact, although it is known to high accuracy. Sample Exercise 2.12
Sample Exercise 1.7 Practice Exercise 2: (a) ammonium bromide, (b) chromium(III)
Practice Exercise 2: (a) four, (b) two, (c) three. oxide, (c) cobalt(II) nitrate
Practice Exercise 2: 9.52 m>s (three significant figures). Practice Exercise 2: (a) HBr, (b) H2CO3
Practice Exercise 2: No. Even though the mass of the gas would then be Practice Exercise 2: (a) SiBr4, (b) S2Cl2, (c) P2O6.
known to four significant figures, the volume of the container would Sample Exercise 2.15
still be known to only three. Practice Exercise 2: No, they are not isomers because they have different
Sample Exercise 1.10 molecular formulas. Butane is C4H10, whereas cyclobutane is C4H8.
Practice Exercise 2: 804.7 km
Sample Exercise 1.11 Chapter 3
Practice Exercise 2: 12 km>L. Sample Exercise 3.1
Sample Exercise 1.12 Practice Exercise 2: (a) C2H4 + 3 O2 ¡ 2 CO2 + 2 H2O. (b) Nine
Practice Exercise 2: 1.2 * 104 ft. O2 molecules

Practice Exercise 2: 832 g Practice Exercise 2: (a) 4, 3, 2; (b) 2, 6, 2, 3; (c) 1, 2, 1, 1, 1
Sample Exercise 3.3
Chapter 2 Practice Exercise 2: (a) HgS1s2 ¡ Hg1l2 + S1s2,
(b) 4 Al1s2 + 3 O21g2 ¡ 2 Al2O31s2
Sample Exercise 2.1
Practice Exercise 2: (a) 154 pm, (b) 1.3 * 106 C atoms Sample Exercise 3.4
Practice Exercise 2: C2H5OH1l2 + 3 O21g2 ¡ 2 CO21g2 + 3 H2O1g2
Sample Exercise 2.2
Practice Exercise 2: (a) 56 protons, 56 electrons, and 82 neutrons, Sample Exercise 3.5
(b) 15 protons, 15 electrons, and 16 neutrons. Practice Exercise 2: (a) 78.0 amu, (b) 32.0 amu, (c) 211.0 amu
Practice Exercise 2: 208
82Pb Practice Exercise 2: 16.1%
Practice Exercise 2: 28.09 amu Practice Exercise 2: 1 mol H2O 16 * 1023 O atoms2 6 3 * 1023
molecules O3 19 * 1023 O atoms2 6 1 mol CO2112 * 1023 O atoms2
Sample Exercise 2.5
Practice Exercise 2: Na, atomic number 11, is a metal; Br, atomic num- Sample Exercise 3.8
ber 35, is a nonmetal. Practice Exercise 2: (a) 9.0 * 1023, (b) 2.71 * 1024
Practice Exercise 2: B5H7 Practice Exercise 2: 164.1 g>mol
A-44
GLOSSARY
absolute zero The lowest attainable temperature; aldehyde An organic compound that contains anion A negatively charged ion. (Section 2.7)
0 K on the Kelvin scale and - 273.15 °C on the a carbonyl group 1C “ O2 to which at least one anode An electrode at which oxidation occurs.
Celsius scale. (Section 1.4) hydrogen atom is attached. (Section 24.4) (Section 20.3)
absorption spectrum A pattern of variation in the alkali metals Members of group 1A in the antibonding molecular orbital A molecular
amount of light absorbed by a sample as a function periodic table. (Section 7.7) orbital in which electron density is concentrated
of wavelength. (Section 23.5) outside the region between the two nuclei of
alkaline earth metals Members of group 2A in
accuracy A measure of how closely individual the periodic table. (Section 7.7) bonded atoms. Such orbitals, designated as s* or
measurements agree with the correct value. p*, are less stable (of higher energy) than bonding
alkanes Compounds of carbon and hydrogen
(Section 1.5) molecular orbitals. (Section 9.7)
containing only carbon–carbon single bonds.
acid A substance that is able to donate a H + ion (Sections 2.9 and 24.2) antiferromagnetism A form of magnetism in
(a proton) and, hence, increases the concentration which unpaired electron spins on adjacent sites
alkenes Hydrocarbons containing one or more
of H + 1aq2 when it dissolves in water. (Section 4.3) carbon–carbon double bonds. (Section 24.2)
point in opposite directions and cancel each
other’s effects. (Section 23.1)
acid-dissociation constant (Ka) An equilibrium
alkyl group A group that is formed by removing a
constant that expresses the extent to which an acid aqueous solution A solution in which water is the
hydrogen atom from an alkane. (Section 25.3)
transfers a proton to solvent water. (Section 16.6) solvent. (Chapter 4: Introduction)
alkynes Hydrocarbons containing one or more
acidic anhydride (acidic oxide) An oxide that aromatic hydrocarbons Hydrocarbon compounds
carbon–carbon triple bonds. (Section 24.2)
forms an acid when added to water; soluble that contain a planar, cyclic arrangement of carbon
nonmetal oxides are acidic anhydrides. (Section 22.5) alloy A substance that has the characteristic atoms linked by both s and delocalized p bonds.
properties of a metal and contains more than one (Section 24.2)
acidic oxide (acidic anhydride) An oxide that
element. Often there is one principal metallic Arrhenius equation An equation that relates the
either reacts with a base to form a salt or with
component, with other elements present in rate constant for a reaction to the frequency factor,
water to form an acid. (Section 22.5)
smaller amounts. Alloys may be homogeneous or A, the activation energy, Ea, and the temperature,
acid rain Rainwater that has become excessively heterogeneous. (Section 12.3) T: k = Ae -E a>RT. In its logarithmic form it is
acidic because of absorption of pollutant oxides,
alpha decay A type of radioactive decay in which written ln k = - Ea >RT + ln A. (Section 14.5)
notably SO 3, produced by human activities.
an atomic nucleus emits an alpha particle and atmosphere (atm) A unit of pressure equal to 760
(Section 18.2) thereby transforms (or “decays”) into an atom with torr; 1 atm = 101.325 kPa. (Section 10.2)
actinide element Element in which the 5f orbitals a mass number 4 less and atomic number 2 less.
are only partially occupied. (Section 6.8) (Section 21.1) atom The smallest representative particle of an
element. (Sections 1.1 and 2.1)
activated complex (transition state) The alpha 1A 2 helix A protein structure in which
particular arrangement of atoms found at the top of the protein is coiled in the form of a helix with atomic mass unit (amu) A unit based on the
the potential-energy barrier as a reaction proceeds hydrogen bonds between C “ O and N ¬ H value of exactly 12 amu for the mass of the isotope
from reactants to products. (Section 14.5) groups on adjacent turns. (Section 24.7) of carbon that has six protons and six neutrons in
the nucleus. (Sections 2.3 and 3.3)
activation energy (Ea) The minimum energy alpha particles Particles that are identical to
needed for reaction; the height of the energy helium-4 nuclei, consisting of two protons and two atomic number The number of protons in the
barrier to formation of products. (Section 14.5) neutrons, symbol 42He or 42a. (Section 21.1) nucleus of an atom of an element. (Section 2.3)
active site Specific site on a heterogeneous amide An organic compound that has an NR 2 atomic radius An estimate of the size of an atom.
catalyst or an enzyme where catalysis occurs. group attached to a carbonyl. (Section 24.4) See bonding atomic radius. (Section 7.3)
(Section 14.7) atomic weight The average mass of the atoms
amine A compound that has the general formula
of an element in atomic mass units (amu); it is
activity The decay rate of a radioactive R 3N, where R may be H or a hydrocarbon group.
numerically equal to the mass in grams of one
material, generally expressed as the number of (Section 16.7)
mole of the element. (Section 2.4)
disintegrations per unit time. (Section 21.4) amino acid A carboxylic acid that contains an
amino 1 ¬ NH 22 group attached to the carbon
autoionization The process whereby water
activity series A list of metals in order of
atom adjacent to the carboxylic acid 1 ¬ COOH2
spontaneously forms low concentrations of
H + 1aq2 and OH - 1aq2 ions by proton transfer
decreasing ease of oxidation. (Section 4.4)
functional group. (Section 24.7)
addition polymerization Polymerization that from one water molecule to another.
occurs through coupling of monomers with one amorphous solid A solid whose molecular (Section 16.3)
another, with no other products formed in the arrangement lacks the regularly repeating long-
Avogadro’s hypothesis A statement that equal
reaction. (Section 12.8) range pattern of a crystal. (Section 12.2)
volumes of gases at the same temperature and
addition reaction A reaction in which a reagent amphiprotic Refers to the capacity of a substance pressure contain equal numbers of molecules.
adds to the two carbon atoms of a carbon–carbon to either add or lose a proton 1H + 2. (Section 16.2) (Section 10.3)
multiple bond. (Section 24.3) amphoteric oxides and hydroxides Oxides and Avogadro’s law A statement that the volume of
adsorption The binding of molecules to a surface. hydroxides that are only slightly soluble in water a gas maintained at constant temperature and
(Section 14.7) but that dissolve in either acidic or basic solutions. pressure is directly proportional to the number of
(Section 17.5) moles of the gas. (Section 10.3)
alcohol An organic compound obtained by
substituting a hydroxyl group 1 ¬ OH2 for a angstrom A common non-SI unit of length, Avogadro’s number (NA) The number of
12
hydrogen on a hydrocarbon. (Sections 2.9 denoted Å, that is used to measure atomic C atoms in exactly 12 g of 12C; it equals
and 24.4) dimensions: 1Å = 10 -10 m. (Section 2.3) 6.022 * 1023 mol-1. (Section 3.4)
G-1
G-2 GLOSSARY
band An array of closely spaced molecular bond dipole The dipole moment that is due to capillary action The process by which a liquid
orbitals occupying a discrete range of energy. unequal electron sharing between two atoms in a rises in a tube because of a combination of
(Section 12.4) covalent bond. (Section 9.3) adhesion to the walls of the tube and cohesion
bond enthalpy The enthalpy change, ΔH, between liquid particles. (Section 11.3)
band gap The energy gap between a fully
occupied band called a valence band and an empty required to break a particular bond when the carbide A binary compound of carbon with a
band called the conduction band. (Section 12.7) substance is in the gas phase. (Section 8.8) metal or metalloid. (Section 22.9)
band structure The electronic structure of a solid, bonding atomic radius The radius of an atom carbohydrates A class of substances formed from
defining the allowed ranges of energy for electrons as defined by the distances separating it from polyhydroxy aldehydes or ketones. (Section 24.8)
in a solid. (Section 12.7) other atoms to which it is chemically bonded.
carbon black A microcrystalline form of carbon.
(Section 7.3)
bar A unit of pressure equal to 105 Pa. (Section 22.9)
(Section 10.2) bonding molecular orbital A molecular orbital
carbonyl group The C “ O double bond, a
in which the electron density is concentrated in
base A substance that is an H + acceptor; a base characteristic feature of several organic functional
produces an excess of OH - 1aq2 ions when it
the internuclear region. The energy of a bonding
groups, such as ketones and aldehydes.
molecular orbital is lower than the energy of the
dissolves in water. (Section 4.3) (Section 24.4)
separate atomic orbitals from which it forms.
base-dissociation constant (Kb) An equilibrium (Section 9.7) carboxylic acid A compound that contains the
constant that expresses the extent to which a base ¬ COOH functional group. (Sections 16.10 and
bonding pair In a Lewis structure a pair of
reacts with solvent water, accepting a proton and 24.4)
forming OH - 1aq2. (Section 16.7)
electrons that is shared by two atoms. (Section 9.2)
bond length The distance between the centers of catalyst A substance that changes the speed of
basic anhydride (basic oxide) An oxide that a chemical reaction without itself undergoing
two bonded atoms. (Section 8.3)
forms a base when added to water; soluble metal a permanent chemical change in the process.
oxides are basic anhydrides. (Section 22.5) bond order The number of bonding electron pairs (Section 14.7)
shared between two atoms, minus the number
basic oxide (basic anhydride) An oxide that cathode An electrode at which reduction occurs.
of antibonding electron pairs: bond order =
either reacts with water to form a base or (Section 20.3)
(number of bonding electrons - number of
reacts with an acid to form a salt and water.
antibonding electrons)/2. (Section 9.7) cathode rays Streams of electrons that are
(Section 22.5)
bond polarity A measure of the degree to which produced when a high voltage is applied to
battery A self-contained electrochemical power electrodes in an evacuated tube. (Section 2.2)
the electrons are shared unequally between two
source that contains one or more voltaic cells.
atoms in a chemical bond. (Section 8.4) cathodic protection A means of protecting
(Section 20.7)
boranes Covalent hydrides of boron. a metal against corrosion by making it the
becquerel The SI unit of radioactivity. It cathode in a voltaic cell. This can be achieved by
(Section 22.11)
corresponds to one nuclear disintegration per attaching a more easily oxidized metal, which
second. (Section 21.4) Born–Haber cycle A thermodynamic cycle based serves as an anode, to the metal to be protected.
on Hess’s law that relates the lattice energy of an (Section 20.8)
Beer’s law The light absorbed by a substance
ionic substance to its enthalpy of formation and to
(A) equals the product of its extinction coefficient cation A positively charged ion. (Section 2.7)
1e2, the path length through which the light
other measurable quantities. (Section 8.2)
passes (b), and the molar concentration of the Boyle’s law A law stating that at constant cell potential The potential difference between
substance (c): A = ebc. (Section 14.2) temperature, the product of the volume and the cathode and anode in an electrochemical cell;
pressure of a given amount of gas is a constant. it is measured in volts: 1 V = 1 J>C. Also called
beta emission A nuclear decay process where electromotive force. (Section 20.4)
(Section 10.3)
a beta particle is emitted from the nucleus; also
called beta decay. (Section 21.1) Brønsted–Lowry acid A substance (molecule or cellulose A polysaccharide of glucose; it is the
ion) that acts as a proton donor. (Section 16.2) major structural element in plant matter.
beta particles Energetic electrons emitted from
(Section 24.8)
the nucleus, symbol -10 e or b- . (Section 21.1) Brønsted–Lowry base A substance (molecule
or ion) that acts as a proton acceptor. Celsius scale A temperature scale on which
beta sheet A structural form of protein in which
(Section 16.2) water freezes at 0° and boils at 100° at sea level.
two strands of amino acids are hydrogen-bonded
(Section 1.4)
together in a zipperlike configuration. buffer capacity The amount of acid or base a
(Section 24.7) buffer can neutralize before the pH begins to chain reaction A series of reactions in which one
change appreciably. (Section 17.2) reaction initiates the next. (Section 21.7)
bidentate ligand A ligand in which two linked
coordinating atoms are bound to a metal. buffered solution (buffer) A solution that changes of state Transformations of matter from
(Section 23.3) undergoes a limited change in pH upon addition one state to a different one, for example, from a
of a small amount of acid or base. (Section 17.2) gas to a liquid. (Section 1.3)
bimolecular reaction An elementary reaction
that involves two molecules. (Section 14.6) calcination The heating of an ore to bring about charcoal A form of carbon produced when
its decomposition and the elimination of a volatile wood is heated strongly in a deficiency of air.
biochemistry The study of the chemistry of living
product. For example, a carbonate ore might be (Section 22.9)
systems. (Chapter 24: Introduction)
calcined to drive off CO 2. (Section 23.2) Charles’s law A law stating that at constant
biodegradable Organic material that bacteria are
calorie A unit of energy; it is the amount of pressure, the volume of a given quantity of gas is
able to oxidize. (Section 18.4)
energy needed to raise the temperature of 1 g of proportional to absolute temperature.
body-centered lattice A crystal lattice in which water by 1 °C from 14.5 °C to 15.5 °C. A related (Section 10.3)
the lattice points are located at the center and unit is the joule: 1 cal = 4.184 J. (Section 5.1)
corners of each unit cell. (Section 12.2) chelate effect The generally larger formation
calorimeter An apparatus that measures the heat constants for polydentate ligands as compared
bomb calorimeter A device for measuring the released or absorbed in a chemical or physical with the corresponding monodentate ligands.
heat evolved in the combustion of a substance process. (Section 5.5) (Section 23.3)
under constant-volume conditions. (Section 5.5)
calorimetry The experimental measurement of chelating agent A polydentate ligand that is
bond angles The angles made by the lines joining heat produced in chemical and physical processes. capable of occupying two or more sites in the
the nuclei of the atoms in a molecule. (Section 9.1) (Section 5.5) coordination sphere. (Section 23.3)
GLOSSARY G-3
chemical bond A strong attractive force that terms of the frequency of collisions, the number of conversion factor A ratio relating the same
exists between atoms in a molecule. (Section 8.1) collisions with energies exceeding the activation quantity in two systems of units that is used to
energy, and the probability that the collisions convert the units of measurement. (Section 1.6)
chemical changes Processes in which one
occur with suitable orientations. (Section 14.5)
or more substances are converted into other coordination compound A compound containing
substances; also called chemical reactions. colloids (colloidal dispersions) Mixtures a metal ion bonded to a group of surrounding
(Section 1.3) containing particles larger than normal solutes molecules or ions that act as ligands. (Section 23.2)
but small enough to remain suspended in the
chemical equation A representation of a chemical coordination number The number of adjacent
dispersing medium. (Section 13.6)
reaction using the chemical formulas of the atoms to which an atom is directly bonded. In a
reactants and products; a balanced chemical combination reaction A chemical reaction in complex the coordination number of the metal
equation contains equal numbers of atoms of each which two or more substances combine to form a ion is the number of donor atoms to which it is
element on both sides of the equation. (Section 3.1) single product. (Section 3.2) bonded. (Sections 12.37 and 24.2)
chemical equilibrium A state of dynamic balance combustion reaction A chemical reaction that coordination sphere The metal ion and its
in which the rate of formation of the products of proceeds with evolution of heat and usually also surrounding ligands. (Section 23.2)
a reaction from the reactants equals the rate of a flame; most combustion involves reaction with
oxygen, as in the burning of a match. (Section 3.2) coordination-sphere isomers Structural isomers
formation of the reactants from the products; at
of coordination compounds in which the ligands
equilibrium the concentrations of the reactants common-ion effect A shift of an equilibrium
within the coordination sphere differ. (Section 23.4)
and products remain constant. (Section 4.1; induced by an ion common to the equilibrium.
Chapter 15: Introduction) For example, added Na 2SO 4 decreases the copolymer A complex polymer resulting from
solubility of the slightly soluble salt BaSO 4, or the polymerization of two or more chemically
chemical formula A notation that uses chemical
added NaF decreases the percent ionization of HF. different monomers. (Section 12.8)
symbols with numerical subscripts to convey
the relative proportions of atoms of the different (Section 17.1) core electrons The electrons that are not in the
elements in a substance. (Section 2.6) complementary colors Colors that, when mixed outermost shell of an atom. (Section 6.8)
in proper proportions, appear white or colorless.
chemical kinetics The area of chemistry corrosion The process by which a metal is
(Section 23.5)
concerned with the speeds, or rates, at which oxidized by substances in its environment.
chemical reactions occur. (Chapter 14: complete ionic equation A chemical equation (Section 20.8)
Introduction) in which dissolved strong electrolytes (such
as dissolved ionic compounds) are written as covalent bond A bond formed between two or
chemical nomenclature The rules used in more atoms by a sharing of electrons. (Section 8.1)
separate ions. (Section 4.2)
naming substances. (Section 2.8)
complex ion (complex) An assembly of a metal covalent-network solids Solids in which the units
chemical properties Properties that describe a ion and the Lewis bases (ligands) bonded to it. that make up the three-dimensional network are
substance’s composition and its reactivity; how the (Section 17.5) joined by covalent bonds. (Section 12.1)
substance reacts or changes into other substances.
compound A substance composed of two or critical mass The amount of fissionable material
(Section 1.3)
more elements united chemically in definite necessary to maintain a nuclear chain reaction.
chemical reactions Processes in which one proportions. (Section 1.2) (Section 21.7)
or more substances are converted into other
compound semiconductor A semiconducting critical pressure The pressure at which a gas at
substances; also called chemical changes.
material formed from two or more elements. its critical temperature is converted to a liquid
(Section 1.3)
(Section 12.7) state. (Section 11.4)
chemistry The scientific discipline that studies the
concentration The quantity of solute present in a critical temperature The highest temperature at
composition, properties, and transformations of
given quantity of solvent or solution. (Section 4.5) which it is possible to convert the gaseous form of
matter. (Chapter 1: Introduction)
concentration cell A voltaic cell containing the a substance to a liquid. The critical temperature
chiral A term describing a molecule or an ion same electrolyte and the same electrode materials increases with an increase in the magnitude of
that cannot be superimposed on its mirror image. in both the anode and cathode compartments. The intermolecular forces. (Section 11.4)
(Sections 23.4 and 24.5) emf of the cell is derived from a difference in the crystal-field theory A theory that accounts for the
chlorofluorocarbons Compounds composed concentrations of the same electrolyte solutions in colors and the magnetic and other properties of
entirely of chlorine, fluorine, and carbon. the compartments. (Section 20.6) transition-metal complexes in terms of the splitting of
(Section 18.3) condensation polymerization Polymerization the energies of metal ion d orbitals by the electrostatic
chlorophyll A plant pigment that plays a major in which molecules are joined together through interaction with the ligands. (Section 23.6)
role in conversion of solar energy to chemical condensation reactions. (Section 12.8)
crystal lattice An imaginary network of points
energy in photosynthesis. (Section 23.3) condensation reaction A chemical reaction in on which the repeating motif of a solid may be
cholesteric liquid crystalline phase A liquid which a small molecule (such as a molecule of imagined to be laid down so that the structure of the
crystal formed from flat, disc-shaped molecules water) is split out from between two reacting crystal is obtained. The motif may be a single atom
that align through a stacking of the molecular molecules. (Sections 12.6 and 22.8) or a group of atoms. Each lattice point represents an
discs. (Section 11.7) conduction band A band of molecular orbitals identical environment in the crystal. (Section 12.2)
coal A naturally occurring solid containing lying higher in energy than the occupied valence crystalline solid (crystal) A solid whose internal
hydrocarbons of high molecular weight, as well band and distinctly separated from it. (Section 12.7) arrangement of atoms, molecules, or ions
as compounds containing sulfur, oxygen, and conjugate acid A substance formed by addition of possesses a regularly repeating pattern in any
nitrogen. (Section 5.8) a proton to a Brønsted–Lowry base. (Section 16.2) direction through the solid. (Section 12.2)
colligative property A property of a solvent conjugate acid–base pair An acid and a base, crystallization The process in which molecules,
(vapor-pressure lowering, freezing-point lowering, such as H 2O and OH - , that differ only in the ions, or atoms come together to form a crystalline
boiling-point elevation, osmotic pressure) that presence or absence of a proton. (Section 16.2) solid. (Section 13.2)
depends on the total concentration of solute conjugate base A substance formed by the loss of a cubic close packing A crystal structure where
particles present. (Section 13.5) proton from a Brønsted–Lowry acid. (Section 16.2) the atoms are packed together as close as possible,
collision model A model of reaction rates based continuous spectrum A spectrum that contains and the close-packed layers of atoms adopt a
on the idea that molecules must collide to react; it radiation distributed over all wavelengths. three-layer repeating pattern that leads to a face-
explains the factors influencing reaction rates in (Section 6.3) centered cubic unit cell. (Section 12.3)
G-4 GLOSSARY
curie A measure of radioactivity: 1 curie = dipole moment A measure of the separation and electromotive force (emf) A measure of the
3.7 * 1010 nuclear disintegrations per second. magnitude of the positive and negative charges in driving force, or electrical pressure, for the
(Section 21.4) polar molecules. (Section 8.4) completion of an electrochemical reaction.
dispersion forces Intermolecular forces resulting Electromotive force is measured in volts:
cycloalkanes Saturated hydrocarbons of general
from attractions between induced dipoles. Also 1 V = 1 J>C. Also called the cell potential.
formula CnH 2n in which the carbon atoms form a
called London dispersion forces. (Section 11.2) (Section 20.4)
closed ring. (Section 24.2)
displacement reaction A reaction in which an electron A negatively charged subatomic particle
Dalton’s law of partial pressures A law stating
element reacts with a compound, displacing an found outside the atomic nucleus; it is a part of all
that the total pressure of a mixture of gases is the
element from it. (Section 4.4) atoms. An electron has a mass 1>1836 times that
sum of the pressures that each gas would exert if it
of a proton. (Section 2.3)
were present alone. (Section 10.6) donor atom The atom of a ligand that bonds to
the metal. (Section 23.2) electron affinity The energy change that occurs
d–d transition The transition of an electron in a
when an electron is added to a gaseous atom or
transition-metal compound from a lower-energy d doping Incorporation of a hetero atom into a solid ion. (Section 7.5)
orbital to a higher-energy d orbital. (Section 23.6) to change its electrical properties. For example,
incorporation of P into Si. (Section 12.7) electron capture A mode of radioactive decay in
decomposition reaction A chemical reaction in
which an inner-shell orbital electron is captured
which a single compound reacts to give two or double bond A covalent bond involving two by the nucleus. (Section 21.1)
more products. (Section 3.2) electron pairs. (Section 8.3)
electron configuration The arrangement of
degenerate A situation in which two or more double helix The structure for DNA that involves electrons in the orbitals of an atom or molecule
orbitals have the same energy. (Section 6.7) the winding of two DNA polynucleotide chains (Section 6.8)
delocalized electrons Electrons that are spread together in a helical arrangement. The two
strands of the double helix are complementary electron density The probability of finding an
over a number of atoms in a molecule or a crystal
in that the organic bases on the two strands are electron at any particular point in an atom; this
rather than localized on a single atom or a pair of
paired for optimal hydrogen bond interaction. probability is equal to c2, the square of the wave
atoms. (Section 9.6)
(Section 24.10) function. Also called the probability density.
density The ratio of an object’s mass to its (Section 6.5)
volume. (Section 1.4) dynamic equilibrium A state of balance in
which opposing processes occur at the same rate. electron domain In the VSEPR model, a region
deoxyribonucleic acid (DNA) A polynucleotide about a central atom in which an electron pair is
(Section 11.5)
in which the sugar component is deoxyribose. concentrated. (Section 9.2)
(Section 24.10) effective nuclear charge The net positive charge
experienced by an electron in a many-electron electron-domain geometry The three-
desalination The removal of salts from seawater, dimensional arrangement of the electron domains
atom; this charge is not the full nuclear charge
brine, or brackish water to make it fit for human around an atom according to the VSEPR model.
because there is some shielding of the nucleus by
consumption. (Section 18.4) (Section 9.2)
the other electrons in the atom. (Section 7.2)
deuterium The isotope of hydrogen whose electronegativity A measure of the ability of an
effusion The escape of a gas through an orifice or
nucleus contains a proton and a neutron: 21H. atom that is bonded to another atom to attract
hole. (Section 10.8)
(Section 22.2) electrons to itself. (Section 8.4)
elastomer A material that can undergo a
dextrorotatory, or merely dextro or d A term electronic charge The negative charge
substantial change in shape via stretching,
used to label a chiral molecule that rotates the carried by an electron; it has a magnitude of
bending, or compression and return to its original
plane of polarization of plane-polarized light to 1.602 * 10 -19 C. (Section 2.3)
shape upon release of the distorting force.
the right (clockwise). (Section 23.4)
(Section 12.6) electronic structure The arrangement of
diamagnetism A type of magnetism that causes a electrons in an atom or molecule. (Chapter 6:
electrochemistry The branch of chemistry
substance with no unpaired electrons to be weakly Introduction)
that deals with the relationships between
repelled from a magnetic field. (Section 9.8)
electricity and chemical reactions. (Chapter 20: electron-sea model A model for the behavior of
diatomic molecule A molecule composed of only Introduction) electrons in metals. (Section 12.4)
two atoms. (Section 2.6)
electrolysis reaction A reaction in which a electron shell A collection of orbitals that have
diffusion The spreading of one substance through nonspontaneous redox reaction is brought about the same value of n. For example, the orbitals with
a space occupied by one or more other substances. by the passage of current under a sufficient n = 3 (the 3s, 3p, and 3d orbitals) comprise the
(Section 10.8) external electrical potential. The devices in which third shell. (Section 6.5)
dilution The process of preparing a less electrolysis reactions occur are called electrolytic
electron spin A property of the electron that
concentrated solution from a more concentrated cells. (Section 20.9)
makes it behave as though it were a tiny magnet.
one by adding solvent. (Section 4.5) electrolyte A solute that produces ions in The electron behaves as if it were spinning on its
dimensional analysis A method of problem solution; an electrolytic solution conducts an axis; electron spin is quantized. (Section 6.7)
solving in which units are carried through all electric current. (Section 4.1)
element A substance consisting of atoms of the
calculations. Dimensional analysis ensures that the electrolytic cell A device in which a same atomic number. Historically defined as a
final answer of a calculation has the desired units. nonspontaneous oxidation–reduction reaction substance that cannot be separated into simpler
(Section 1.6) is caused to occur by passage of current under a substances by chemical means. (Sections 1.1
dipole A molecule with one end having a partial sufficient external electrical potential. and 1.2)
negative charge and the other end having a partial (Section 20.9)
elemental semiconductor A semiconducting
positive charge; a polar molecule. (Section 8.4) electromagnetic radiation (radiant energy) A material composed of just one element.
dipole–dipole force A force that becomes form of energy that has wave characteristics (Section 12.7)
significant when polar molecules come in close and that propagates through a vacuum at the
elementary reaction A process in a
contact with one another. The force is attractive characteristic speed of 3.00 * 108 m >s.
chemical reaction that occurs in a single step.
when the positive end of one polar molecule (Section 6.1)
An overall chemical reaction consists of one or
approaches the negative end of another. electrometallurgy The use of electrolysis to more elementary reactions or steps.
(Section 11.2) reduce or refine metals. (Section 20.9) (Section 14.6)
GLOSSARY G-5
empirical formula A chemical formula that shows molecular equation appears to involve the free energy (Gibbs free energy, G) A
the kinds of atoms and their relative numbers in a exchange of ions between the two reactants. thermodynamic state function that gives a criterion
substance in the smallest possible whole-number (Section 4.2) for spontaneous change in terms of enthalpy and
ratios. (Section 2.6) entropy: G = H - TS. (Section 19.5)
excited state A higher energy state than the
enantiomers Two mirror-image molecules ground state. (Section 6.3) free radical A substance with one or more
of a chiral substance. The enantiomers are unpaired electrons. (Section 21.9)
exothermic process A process in which a system
nonsuperimposable. (Section 23.4)
releases heat to its surroundings. (Section 5.2) frequency The number of times per second that
endothermic process A process in which a system one complete wavelength passes a given point.
extensive property A property that depends on
absorbs heat from its surroundings. (Section 5.2) (Section 6.1)
the amount of material considered; for example,
energy The capacity to do work or to transfer mass or volume. (Section 1.3) frequency factor (A) A term in the Arrhenius
heat. (Section 5.1) equation that is related to the frequency of
face-centered lattice A crystal lattice in which
energy-level diagram A diagram that shows collision and the probability that the collisions are
the lattice points are located at the faces and
the energies of molecular orbitals relative to the favorably oriented for reaction. (Section 14.5)
corners of each unit cell. (Section 12.2)
atomic orbitals from which they are derived. Also fuel cell A voltaic cell that utilizes the oxidation
Faraday constant (F ) The magnitude of charge
of one mole of electrons: 96,500 C >mol. (Section
called a molecular-orbital diagram. (Section 9.7) of a conventional fuel, such as H 2 or CH 4, in the
enthalpy A quantity defined by the relationship cell reaction. (Section 20.7)
20.5)
H = E + PV; the enthalpy change, ΔH, for fuel value The energy released when 1 g of a
f-block metals Lanthanide and actinide elements
a reaction that occurs at constant pressure is substance is combusted. (Section 5.8)
in which the 4f or 5f orbitals are partially
the heat evolved or absorbed in the reaction:
occupied. (Section 6.9) functional group An atom or group of atoms that
ΔH = qp. (Section 5.3)
imparts characteristic chemical properties to an
ferrimagnetism A form of magnetism in which
enthalpy of formation The enthalpy change that organic compound. (Section 24.1)
unpaired electron spins on different-type ions
accompanies the formation of a substance from
point in opposite directions but do not fully cancel fusion The joining of two light nuclei to form a
the most stable forms of its component elements.
out. (Section 23.1) more massive one. (Section 21.6)
(Section 5.7)
ferromagnetism A form of magnetism in which galvanic cell See voltaic cell. (Section 20.3)
enthalpy of reaction The enthalpy change
unpaired electron spins align parallel to one
associated with a chemical reaction. (Section 5.4) gamma radiation Energetic electromagnetic
another. (Section 23.1) radiation emanating from the nucleus of a
entropy A thermodynamic function associated
first law of thermodynamics A statement that radioactive atom. (Section 21.1)
with the number of different equivalent energy
energy is conserved in any process. One way
states or spatial arrangements in which a system gas Matter that has no fixed volume or shape; it
to express the law is that the change in internal
may be found. It is a thermodynamic state conforms to the volume and shape of its container.
energy, ΔE, of a system in any process is equal
function, which means that once we specify the (Section 1.2)
to the heat, q, added to the system, plus the
conditions for a system—that is, the temperature, gas constant (R) The constant of proportionality
work, w, done on the system by its surroundings:
pressure, and so on—the entropy is defined. in the ideal-gas equation. (Section 10.4)
ΔE = q + w. (Section 5.2)
(Section 19.2)
first-order reaction A reaction in which the geometric isomerism A form of isomerism in
enzyme A protein molecule that acts to catalyze which compounds with the same type and number
reaction rate is proportional to the concentration
specific biochemical reactions. (Section 14.7) of atoms and the same chemical bonds have
of a single reactant, raised to the first power.
equilibrium constant The numerical value of the (Section 14.4) different spatial arrangements of these atoms and
equilibrium-constant expression for a system at bonds. (Sections 23.4 and 24.4)
fission The splitting of a large nucleus into two
equilibrium. The equilibrium constant is most Gibbs free energy A thermodynamic state
smaller ones. (Section 21.6)
usually denoted by Kp for gas-phase systems or Kc function that combines enthalpy and entropy, in
for solution-phase systems. (Section 15.2) folding The process by which a protein adopts its the form G = H - TS. For a change occurring
biologically active shape. (Section 24.7)
equilibrium-constant expression The expression at constant temperature and pressure, the
that describes the relationship among the force A push or a pull. (Section 5.1) change in free energy is ΔG = ΔH - TΔS.
concentrations (or partial pressures) of the (Section 19.5)
formal charge The number of valence electrons
substances present in a system at equilibrium. in an isolated atom minus the number of electrons glass An amorphous solid formed by fusion of
The numerator is obtained by multiplying the assigned to the atom in the Lewis structure. SiO 2, CaO, and Na 2O. Other oxides may also be
concentrations of the substances on the product (Section 8.5) used to form glasses with differing characteristics.
side of the equation, each raised to a power equal (Section 22.10)
to its coefficient in the chemical equation. The formation constant For a metal ion complex, the
denominator similarly contains the concentrations equilibrium constant for formation of the complex glucose A polyhydroxy aldehyde whose
of the substances on the reactant side of the from the metal ion and base species present in formula is CH 2OH1CHOH24CHO; it is the
equation. (Section 15.2) solution. It is a measure of the tendency of the most important of the monosaccharides.
complex to form. (Section 17.5) (Section 24.8)
equivalence point The point in a titration at
which the added solute reacts completely with the formula weight The mass of the collection of glycogen The general name given to a group of
solute present in the solution. (Section 4.6) atoms represented by a chemical formula. For polysaccharides of glucose that are synthesized
example, the formula weight of NO 2 (46.0 amu) in mammals and used to store energy from
ester An organic compound that has an OR is the sum of the masses of one nitrogen atom and carbohydrates. (Section 24.7)
group attached to a carbonyl; it is the product of a two oxygen atoms. (Section 3.3)
reaction between a carboxylic acid and an alcohol. Graham’s law A law stating that the rate of
(Section 24.4) fossil fuels Coal, oil, and natural gas, which are effusion of a gas is inversely proportional to the
presently our major sources of energy. (Section 5.8) square root of its molecular weight. (Section 10.8)
ether A compound in which two hydrocarbon
groups are bonded to one oxygen. (Section 24.4) fracking The practice in which water laden gray (Gy) The SI unit for radiation dose
with sand and other materials is pumped at high corresponding to the absorption of 1 J of
exchange (metathesis) reaction A reaction pressure into rock formations to release natural energy per kilogram of biological material;
between compounds that when written as a gas and other petroleum materials. (Section 18.4) 1 Gy = 100 rads. (Section 21.9)
G-6 GLOSSARY
green chemistry Chemistry that promotes the two or more distinct phases with characteristic ideal gas A hypothetical gas whose pressure,
design and application of chemical products compositions are present. (Section 12.3) volume, and temperature behavior is completely
and processes that are compatible with human described by the ideal-gas equation. (Section 10.4)
heterogeneous catalyst A catalyst that is
health and that preserve the environment. ideal-gas equation An equation of state for
in a different phase from that of the reactant
(Section 18.5) gases that embodies Boyle’s law, Charles’s law, and
substances. (Section 14.7)
greenhouse gases Gases in an atmosphere that Avogadro’s hypothesis in the form PV = nRT.
heterogeneous equilibrium The equilibrium
absorb and emit infrared radiation (radiant heat), (Section 10.4)
established between substances in two or more
“trapping” heat in the atmosphere. (Section 18.2) ideal solution A solution that obeys Raoult’s law.
different phases, for example, between a gas and a
ground state The lowest-energy, or most stable, solid or between a solid and a liquid. (Section 13.5)
state. (Section 6.3) (Section 15.4) immiscible liquids Liquids that do not dissolve
group Elements that are in the same column of hexagonal close packing A crystal structure in one another to a significant extent.
the periodic table; elements within the same group where the atoms are packed together as closely (Section 13.3)
or family exhibit similarities in their chemical as possible. The close-packed layers adopt a two- indicator A substance added to a solution that
behavior. (Section 2.5) layer repeating pattern, which leads to a primitive changes color when the added solute has reacted
Haber process The catalyst system and hexagonal unit cell. (Section 12.3) with all the solute present in solution. The
conditions of temperature and pressure developed high-spin complex A complex whose electrons most common type of indicator is an acid–base
by Fritz Haber and coworkers for the formation of populate the d orbitals to give the maximum indicator whose color changes as a function of pH.
NH 3 from H 2 and N 2. (Section 15.2) number of unpaired electrons. (Section 23.6) (Section 4.6)
half-life The time required for the concentration hole A vacancy in the valence band of a instantaneous rate The reaction rate at a
of a reactant substance to decrease to half its initial semiconductor, created by doping. (Section 12.7) particular time as opposed to the average rate over
value; the time required for half of a sample of a an interval of time. (Section 14.2)
particular radioisotope to decay. (Sections 14.4 homogeneous catalyst A catalyst that is in the same
phase as the reactant substances. (Section 14.7) insulators Materials that do not conduct
and 21.4) electricity. (Section 12.7)
half-reaction An equation for either an homogeneous equilibrium The equilibrium
established between reactant and product substances intensive property A property that is independent
oxidation or a reduction that explicitly of the amount of material considered, for example,
shows the electrons involved, for example, that are all in the same phase. (Section 15.4)
density. (Section 1.3)
Zn2 + 1aq2 + 2 e - ¡ Zn1s2. (Section 20.2) Hund’s rule A rule stating that electrons occupy
degenerate orbitals in such a way as to maximize interhalogens Compounds formed between two
halogens Members of group 7A in the periodic different halogen elements. Examples include IBr
table. (Section 7.8) the number of electrons with the same spin. In
other words, each orbital has one electron placed and BrF3. (Section 22.4)
hard water Water that contains appreciable in it before pairing of electrons in orbitals occurs. intermediate A substance formed in one elementary
concentrations of Ca2 + and Mg 2 + ; these ions (Section 6.8) step of a multistep mechanism and consumed in
react with soaps to form an insoluble material. another; it is neither a reactant nor an ultimate
(Section 18.4) hybridization The mixing of different types of
product of the overall reaction. (Section 14.6)
atomic orbitals to produce a set of equivalent
heat The flow of energy from a body at intermetallic compound A homogeneous alloy
hybrid orbitals. (Section 9.5)
higher temperature to one at lower temperature with definite properties and a fixed composition.
when they are placed in thermal contact. hybrid orbital An orbital that results from the
Intermetallic compounds are stoichiometric
(Section 5.1) mixing of different kinds of atomic orbitals on
compounds that form between metallic elements.
the same atom. For example, an sp3 hybrid results
heat capacity The quantity of heat required to (Section 12.3)
from the mixing, or hybridizing, of one s orbital
raise the temperature of a sample of matter by 1 °C intermolecular forces The short-range attractive
and three p orbitals. (Section 9.5)
(or 1 K). (Section 5.5) forces operating between the particles that make
hydration Solvation when the solvent is water.
heat of fusion The enthalpy change, ΔH, for up the units of a liquid or solid substance. These
(Section 13.1)
melting a solid. (Section 11.4) same forces also cause gases to liquefy or solidify
hydride ion An ion formed by the addition of an at low temperatures and high pressures.
heat of sublimation The enthalpy change, ΔH,
electron to a hydrogen atom: H - . (Section 7.7) (Chapter 11: Introduction)
for vaporization of a solid. (Section 11.4)
hydrocarbons Compounds composed of only internal energy The total energy possessed by
heat of vaporization The enthalpy change, ΔH,
carbon and hydrogen. (Section 2.9) a system. When a system undergoes a change,
for vaporization of a liquid. (Section 11.4)
hydrogen bonding Bonding that results from the change in internal energy, ΔE, is defined as
Henderson–Hasselbalch equation The intermolecular attractions between molecules the heat, q, added to the system, plus the work,
relationship among the pH, pKa, and the containing hydrogen bonded to an electronegative w, done on the system by its surroundings:
concentrations of acid and conjugate base in an ΔE = q + w. (Section 5.2)
3base4
element. The most important examples involve
aqueous solution: pH = pKa + log . OH, NH, and HF. (Section 11.2)
3acid4
interstitial alloy An alloy in which smaller
(Section 17.2) hydrolysis A reaction with water. When a cation atoms fit into spaces between larger atoms.
or anion reacts with water, it changes the pH. The larger atoms are metallic elements and
Henry’s law A law stating that the concentration (Sections 16.9 and 24.4) the smaller atoms are typically nonmetallic
of a gas in a solution, Sg, is proportional to the hydronium ion 1H3O + 2 The predominant form of elements. (Section 12.3)
pressure of gas over the solution: S g = kPg . the proton in aqueous solution. (Section 16.2) ion Electrically charged atom or group of
(Section 13.3) atoms (polyatomic ion); ions can be positively
hydrophilic Water attracting. The term is often
Hess’s law The heat evolved in a given process used to describe a type of colloid. (Section 13.6) or negatively charged, depending on whether
can be expressed as the sum of the heats of several electrons are lost (positive) or gained (negative) by
processes that, when added, yield the process of hydrophobic Water repelling. The term is often the atoms. (Section 2.7)
interest. (Section 5.6) used to describe a type of colloid. (Section 13.6)
ion–dipole force The force that exists between an
heterogeneous alloy An alloy in which the hypothesis A tentative explanation of a series of ion and a neutral polar molecule that possesses a
components are not distributed uniformly; instead, observations or of a natural law. (Section 1.3) permanent dipole moment. (Section 11.2)
GLOSSARY G-7
ionic bond A bond between oppositely charged lanthanide (rare earth) element Element in liquid Matter that has a distinct volume but no
ions. The ions are formed from atoms by transfer which the 4f subshell is only partially occupied. specific shape. (Section 1.2)
of one or more electrons. (Section 8.1) (Sections 6.8 and 6.9)
liquid crystal A substance that exhibits one or
ionic compound A compound composed of lattice energy The energy required to separate more partially ordered liquid phases above the
cations and anions. (Section 2.7) completely the ions in an ionic solid. (Section 8.2) melting point of the solid form. By contrast, in
ionic hydrides Compounds formed when nonliquid crystalline substances the liquid phase
lattice points Points in a crystal all of which have
hydrogen reacts with alkali metals and also the that forms upon melting is completely unordered.
identical environments. (Section 12.2)
heavier alkaline earths (Ca, Sr, and Ba); these (Section 11.7)
lattice vectors The vectors a, b, and c that define
compounds contain the hydride ion, H - . lock-and-key model A model of enzyme action
a crystal lattice. The position of any lattice point in
(Section 22.2) in which the substrate molecule is pictured as
a crystal can be represented by summing integer
ionic solids Solids that are composed of ions. fitting rather specifically into the active site on the
multiples of the lattice vectors. (Section 12.2)
(Section 12.1) enzyme. It is assumed that in being bound to the
law of constant composition A law that active site, the substrate is somehow activated for
ionization energy The energy required to remove states that the elemental composition of a reaction. (Section 14.7)
an electron from a gaseous atom when the atom is pure compound is always the same, regardless
in its ground state. (Section 7.4) low-spin complex A metal complex in which
of its source; also called the law of definite
the electrons are paired in lower-energy orbitals.
ionizing radiation Radiation that has sufficient proportions. (Section 1.2)
(Section 23.6)
energy to remove an electron from a molecule, law of definite proportions A law that states that
thereby ionizing it. (Section 21.9) magic numbers Numbers of protons and
the elemental composition of a pure substance
neutrons that result in very stable nuclei.
ion-product constant For water, Kw is is always the same, regardless of its source; also
(Section 21.2)
the product of the aquated hydrogen called the law of constant composition.
ion and hydroxide ion concentrations: (Section 1.2) main-group elements Elements in the s and
3H + 43OH - 4 = Kw = 1.0 * 10 -14 at 25 °C. law of mass action The rules by which the
p blocks of the periodic table. (Section 6.9)
(Section 16.3) equilibrium constant is expressed in terms of mass A measure of the amount of material in
irreversible process A process that cannot the concentrations of reactants and products, in an object. It measures the resistance of an object
be reversed to restore both the system and accordance with the balanced chemical equation to being moved. In SI units, mass is measured in
its surroundings to their original states. Any for the reaction. (Section 15.2) kilograms. (Section 1.4)
spontaneous process is irreversible. Le Châtelier’s principle A principle stating that mass defect The difference between the mass of
(Section 19.1) when we disturb a system at chemical equilibrium, a nucleus and the total masses of the individual
isoelectronic series A series of atoms, ions, or the relative concentrations of reactants and nucleons that it contains. (Section 21.6)
molecules having the same number of electrons. products shift so as to partially undo the effects of mass number The sum of the number of protons
(Section 7.3) the disturbance. (Section 15.7) and neutrons in the nucleus of a particular atom.
isomers Compounds whose molecules have the levorotatory, or merely levo or l A term used to (Section 2.3)
same overall composition but different structures. label a chiral molecule that rotates the plane of mass percentage The number of grams of solute
(Sections 2.9 and 23.4) polarization of plane-polarized light to the left in each 100 g of solution. (Section 13.4)
(counterclockwise). (Section 24.4)
isothermal process One that occurs at constant mass spectrometer An instrument used to
temperature. (Section 19.1) Lewis base An electron-pair donor. (Section 16.11) measure the precise masses and relative amounts
Lewis structure A representation of covalent of atomic and molecular ions. (Section 2.4)
isotopes Atoms of the same element containing
different numbers of neutrons and therefore bonding in a molecule that is drawn using Lewis matter Anything that occupies space and has
having different masses. (Section 2.3) symbols. Shared electron pairs are shown as lines, mass; the physical material of the universe.
and unshared electron pairs are shown as pairs of
joule (J) The SI unit of energy, 1 kg@m2 >s2.
(Section 1.1)
dots. Only the valence-shell electrons are shown.
A related unit is the calorie: 4.184 J = 1 cal. matter waves The term used to describe the wave
(Section 8.3)
(Section 5.1) characteristics of a moving particle. (Section 6.4)
Lewis symbol (electron-dot symbol) The
Kelvin scale The absolute temperature scale; the SI mean free path The average distance traveled by a
chemical symbol for an element, with a dot for
unit for temperature is the kelvin. Zero on the Kelvin gas molecule between collisions. (Section 10.8)
each valence electron. (Section 8.1)
scale corresponds to - 273.15 °C. (Section 1.4) metal complex An assembly of a metal ion and
ligand An ion or molecule that coordinates to a the Lewis bases bonded to it. (Section 23.2)
ketone A compound in which the carbonyl group metal atom or to a metal ion to form a complex.
1C “ O2 occurs at the interior of a carbon chain (Section 23.2) metallic bond Bonding, usually in solid metals, in
and is therefore flanked by carbon atoms. which the bonding electrons are relatively free to
(Section 24.4) limiting reactant (limiting reagent) The reactant move throughout the three-dimensional structure.
present in the smallest stoichiometric quantity in (Section 8.1)
kinetic energy The energy that an object a mixture of reactants; the amount of product that
possesses by virtue of its motion. (Section 5.1) can form is limited by the complete consumption metallic character The extent to which an
kinetic-molecular theory A set of assumptions of the limiting reactant. (Section 3.7) element exhibits the physical and chemical
about the nature of gases. These assumptions, properties characteristic of metals, for example,
line spectrum A spectrum that contains radiation luster, malleability, ductility, and good thermal and
when translated into mathematical form, yield the
at only certain specific wavelengths. (Section 6.3) electrical conductivity. (Section 7.6)
ideal-gas equation. (Section 10.7)
linkage isomers Structural isomers of metallic elements (metals) Elements that are
lanthanide contraction The gradual decrease
coordination compounds in which a ligand differs usually solids at room temperature, exhibit high
in atomic and ionic radii with increasing atomic
in its mode of attachment to a metal ion. electrical and heat conductivity, and appear
number among the lanthanide elements, atomic
(Section 23.4) lustrous. Most of the elements in the periodic table
numbers 57 through 70. The decrease arises
because of a gradual increase in effective lipid A nonpolar molecule derived from glycerol are metals. (Sections 2.5 and 12.1)
nuclear charge through the lanthanide series. and fatty acids that is used by organisms for long- metallic hydrides Compounds formed when
(Section 23.1) term energy storage. (Section 24.9) hydrogen reacts with transition metals; these
G-8 GLOSSARY
compounds contain the hydride ion, H - . molecular equation A chemical equation in nematic liquid crystalline phase A liquid crystal
(Section 22.2) which the formula for each substance is written in which the molecules are aligned in the same
without regard for whether it is an electrolyte or a general direction, along their long axes, but in
metallic solids Solids that are composed of metal
nonelectrolyte. (Section 4.2) which the ends of the molecules are not aligned.
atoms. (Section 12.1)
molecular formula A chemical formula that (Section 11.7)
metalloids Elements that lie along the diagonal
indicates the actual number of atoms of each Nernst equation An equation that relates the cell emf,
line separating the metals from the nonmetals in
element in one molecule of a substance. E, to the standard emf, E°, and the reaction quotient,
the periodic table; the properties of metalloids
are intermediate between those of metals and
(Section 2.6) Q: E = E° - 1RT>nF2 ln Q. (Section 20.6)
nonmetals. (Section 2.5) molecular geometry The arrangement in space of net ionic equation A chemical equation for
the atoms of a molecule. (Section 9.2) a solution reaction in which soluble strong
metallurgy The science of extracting metals from
their natural sources by a combination of chemical molecular hydrides Compounds formed when electrolytes are written as ions and spectator ions
and physical processes. It is also concerned with hydrogen reacts with nonmetals and metalloids. are omitted. (Section 4.2)
the properties and structures of metals and alloys. (Section 22.2) neutralization reaction A reaction in which
(Section 23.1) molecularity The number of molecules that an acid and a base react in stoichiometrically
metathesis (exchange) reaction A reaction in participate as reactants in an elementary reaction. equivalent amounts; the neutralization reaction
which two substances react through an exchange of (Section 14.6) between an acid and a metal hydroxide produces
their component ions: AX + BY ¡ AY + BX. water and a salt. (Section 4.3)
molecular orbital (MO) An allowed state for an
Precipitation and acid–base neutralization reactions electron in a molecule. According to molecular- neutron An electrically neutral particle found in
are examples of metathesis reactions. orbital theory, a molecular orbital is entirely the nucleus of an atom; it has approximately the
(Section 4.2) analogous to an atomic orbital, which is an allowed same mass as a proton. (Section 2.3)
metric system A system of measurement used in state for an electron in an atom. Most bonding noble gases Members of group 8A in the
science and in most countries. The meter and the molecular orbitals can be classified as s or p, periodic table. (Section 7.8)
gram are examples of metric units. (Section 1.4) depending on the disposition of electron density
node Points in an atom at which the electron
with respect to the internuclear axis. (Section 9.7)
microstate The state of a system at a particular density is zero. For example, the node in a 2s
instant; one of many possible energetically molecular-orbital diagram A diagram that shows orbital is a spherical surface. (Section 6.6)
equivalent ways to arrange the components of a the energies of molecular orbitals relative to the
nonbonding pair In a Lewis structure a pair of
system to achieve a particular state. (Section 19.3) atomic orbitals from which they are derived; also
electrons assigned completely to one atom; also
called an energy-level diagram. (Section 9.7)
mineral A solid, inorganic substance occurring in called a lone pair. (Section 9.2)
nature, such as calcium carbonate, which occurs as molecular-orbital theory A theory that accounts
nonelectrolyte A substance that does not ionize
calcite. (Section 23.1) for the allowed states for electrons in molecules.
in water and consequently gives a nonconducting
(Section 9.7)
miscible liquids Liquids that mix in all solution. (Section 4.1)
proportions. (Section 13.3) molecular solids Solids that are composed of
nonionizing radiation Radiation that does not
molecules. (Sections 12.1 and 12.6)
mixture A combination of two or more substances have sufficient energy to remove an electron from
in which each substance retains its own chemical molecular weight The mass of the collection of a molecule. (Section 21.9)
identity. (Section 1.2) atoms represented by the chemical formula for a
nonmetallic elements (nonmetals) Elements
molecule. (Section 3.3)
molal boiling-point-elevation constant (Kb) A in the upper right corner of the periodic table;
constant characteristic of a particular solvent molecule A chemical combination of two or more nonmetals differ from metals in their physical and
that gives the increase in boiling point as a atoms. (Sections 1.1 and 2.6) chemical properties. (Section 2.5)
function of solution molality: ΔTb = Kbm. mole fraction The ratio of the number of moles of nonpolar covalent bond A covalent bond in which
(Section 13.5) one component of a mixture to the total moles of the electrons are shared equally. (Section 8.4)
molal freezing-point-depression constant (Kf ) A all components; abbreviated X, with a subscript to
identify the component. (Section 10.6) normal boiling point The boiling point at 1 atm
constant characteristic of a particular solvent that
pressure. (Section 11.5)
gives the decrease in freezing point as a function momentum The product of the mass, m, and
of solution molality: ΔTf = - Kf m. (Section 13.5) velocity, v, of an object. (Section 6.4) normal melting point The melting point at 1 atm
pressure. (Section 11.6)
molality The concentration of a solution monodentate ligand A ligand that binds to the
expressed as moles of solute per kilogram of metal ion via a single donor atom. It occupies one nuclear binding energy The energy required to
solvent; abbreviated m. (Section 13.4) position in the coordination sphere. (Section 23.3) decompose an atomic nucleus into its component
protons and neutrons. (Section 21.6)
molar heat capacity The heat required to raise monomers Molecules with low molecular weights,
the temperature of one mole of a substance by which can be joined together (polymerized) to nuclear disintegration series A series of nuclear
1 °C. (Section 5.5) form a polymer. (Section 12.8) reactions that begins with an unstable nucleus
and terminates with a stable one; also called a
molarity The concentration of a solution monosaccharide A simple sugar, most commonly radioactive series. (Section 21.2)
expressed as moles of solute per liter of solution; containing six carbon atoms. The joining together
abbreviated M. (Section 4.5) of monosaccharide units by condensation nuclear model Model of the atom with a nucleus
reactions results in formation of polysaccharides. containing protons and neutrons and with electrons
molar mass The mass of one mole of a substance
(Section 24.8) in the space outside the nucleus. (Section 2.2)
in grams; it is numerically equal to the formula
weight in atomic mass units. (Section 3.4) nanomaterial A solid whose dimensions range nuclear transmutation A conversion of one kind
from 1 to 100 nm and whose properties differ of nucleus to another. (Section 21.3)
mole A collection of Avogadro’s number
16.022 * 10232 of objects; for example, a mole of from those of a bulk material with the same nucleic acids Polymers of high molecular weight
H 2O is 6.022 * 1023 H 2O molecules. composition. (Section 12.1) that carry genetic information and control protein
(Section 3.4) natural gas A naturally occurring mixture of synthesis. (Section 24.10)
molecular compound A compound that consists gaseous hydrocarbon compounds composed of nucleon A particle found in the nucleus of an
of molecules. (Section 2.6) hydrogen and carbon. (Section 5.8) atom. (Section 21.1)
GLOSSARY G-9
nucleotide Compounds formed from a oxyacid A compound in which one or more OH phase diagram A graphic representation of the
molecule of phosphoric acid, a sugar molecule, groups, and possibly additional oxygen atoms, are equilibria among the solid, liquid, and gaseous
and an organic nitrogen base. Nucleotides bonded to a central atom. (Section 16.10) phases of a substance as a function of temperature
form linear polymers called DNA and RNA, and pressure. (Section 11.6)
oxyanion A polyatomic anion that contains one or
which are involved in protein synthesis and cell phospholipid A form of lipid molecule that
more oxygen atoms. (Section 2.8)
reproduction. (Section 24.10) contains charged phosphate groups. (Section 24.9)
ozone The name given to O 3, an allotrope of
nucleus The very small, very dense, positively photochemical smog A complex mixture of
oxygen. (Section 7.8)
charged portion of an atom; it is composed of undesirable substances produced by the action of
protons and neutrons. (Section 2.2) paramagnetism A property that a substance
sunlight on an urban atmosphere polluted with
possesses if it contains one or more unpaired
octet rule A rule stating that bonded atoms tend automobile emissions. The major starting ingredients
electrons. A paramagnetic substance is drawn into
to possess or share a total of eight valence-shell are nitrogen oxides and organic substances, notably
a magnetic field. (Section 9.8)
electrons. (Section 8.1) olefins and aldehydes. (Section 18.2)
partial pressure The pressure exerted by a
optical isomerism A form of isomerism in which photodissociation The breaking of a molecule
particular gas in a mixture. (Section 10.6)
the two forms of a compound (stereoisomers) are into two or more neutral fragments as a result of
nonsuperimposable mirror images. (Section 23.4) particle accelerator A device that uses strong absorption of light. (Section 18.2)
magnetic and electrostatic fields to accelerate
optically active Possessing the ability to rotate photoelectric effect The emission of electrons
charged particles. (Section 21.3)
the plane of polarized light. (Section 23.4) from a metal surface induced by light.
parts per billion (ppb) The concentration of a (Section 6.2)
orbital An allowed energy state of an electron in
solution in grams of solute per 109 (billion) grams
the quantum mechanical model of the atom; the photoionization The removal of an electron
of solution; equals micrograms of solute per liter
term orbital is also used to describe the spatial from an atom or molecule by absorption of light.
of solution for aqueous solutions. (Section 13.4)
distribution of the electron. An orbital is defined (Section 18.2)
by the values of three quantum numbers: n, l, and parts per million (ppm) The concentration of a
photon The smallest increment (a quantum) of
ml (Section 6.5) solution in grams of solute per 106 (million) grams
radiant energy; a photon of light with frequency n
of solution; equals milligrams of solute per liter of
organic chemistry The study of carbon- has an energy equal to hn. (Section 6.2)
solution for aqueous solutions. (Section 13.4)
containing compounds, typically containing photosynthesis The process that occurs in plant
carbon–carbon bonds. (Section 2.9; Chapter 24: pascal (Pa) The SI unit of pressure:
1 Pa = 1 N >m2. (Section 10.2)
leaves by which light energy is used to convert
Introduction) carbon dioxide and water to carbohydrates and
osmosis The net movement of solvent through Pauli exclusion principle A rule stating that oxygen. (Section 23.3)
a semipermeable membrane toward the solution no two electrons in an atom may have the same
physical changes Changes (such as a phase
with greater solute concentration. (Section 13.5) four quantum numbers (n, l, ml, and ms ). As a
change) that occur with no change in chemical
reflection of this principle, there can be no more
osmotic pressure The pressure that must be composition. (Section 1.3)
than two electrons in any one atomic orbital.
applied to a solution to stop osmosis from pure physical properties Properties that can be
(Section 6.7)
solvent into the solution. (Section 13.5) measured without changing the composition of a
peptide bond A bond formed between two amino
Ostwald process An industrial process used substance, for example, color and freezing point.
acids. (Section 24.7)
to make nitric acid from ammonia. The NH 3 (Section 1.3)
is catalytically oxidized by O 2 to form NO; NO percent ionization The percent of a substance pi 1P 2 bond A covalent bond in which electron
in air is oxidized to NO 2; HNO 3 is formed in a that undergoes ionization on dissolution in water. density is concentrated above and below the
disproportionation reaction when NO 2 dissolves The term applies to solutions of weak acids and internuclear axis. (Section 9.6)
in water. (Section 22.7) bases. (Section 16.6)
pi 1P 2 molecular orbital A molecular orbital
overall reaction order The sum of the reaction percent yield The ratio of the actual that concentrates the electron density on opposite
orders of all the reactants appearing in the rate (experimental) yield of a product to its theoretical sides of an imaginary line that passes through the
expression when the rate can be expressed as (calculated) yield, multiplied by 100. (Section 3.7) nuclei. (Section 9.8)
rate = k3A4a3B4b… . (Section 14.3) period The row of elements that lie in a horizontal Planck constant (h) The constant that relates the
overlap The extent to which atomic orbitals on row in the periodic table. (Section 2.5) energy and frequency of a photon, E = hn. Its
different atoms share the same region of space. periodic table The arrangement of elements in value is 6.626 * 10 -34 J@s. (Section 6.2)
When the overlap between two orbitals is large, a order of increasing atomic number, with elements plastic A material that can be formed into
strong bond may be formed. (Section 9.4) having similar properties placed in vertical particular shapes by application of heat and
oxidation A process in which a substance loses columns. (Section 2.5) pressure. (Section 12.8)
one or more electrons. (Section 4.4) petroleum A naturally occurring combustible polar covalent bond A covalent bond in which
oxidation number (oxidation state) A positive or liquid composed of hundreds of hydrocarbons and the electrons are not shared equally. (Section 8.4)
negative whole number assigned to an element in other organic compounds. (Section 5.8)
polarizability The ease with which the electron
a molecule or ion on the basis of a set of formal pH The negative log in base 10 of the aquated cloud of an atom or a molecule is distorted by
rules; to some degree it reflects the positive or hydrogen ion concentration: pH = - log3H + 4. an outside influence, thereby inducing a dipole
negative character of that atom. (Section 4.4) (Section 16.4) moment. (Section 11.2)
oxidation–reduction (redox) reaction A pH titration curve A graph of pH as a function of polar molecule A molecule that possesses a
chemical reaction in which the oxidation states added titrant. (Section 17.3) nonzero dipole moment. (Section 8.4)
of certain atoms change. (Section 4.4; Chapter 20:
phase change The conversion of a substance polyatomic ion An electrically charged group of
Introduction)
from one state of matter to another. The phase two or more atoms. (Section 2.7)
oxidizing agent, or oxidant The substance that changes we consider are melting and freezing
is reduced and thereby causes the oxidation of 1solid Δ liquid2, sublimation and polydentate ligand A ligand in which two or
some other substance in an oxidation–reduction deposition, and vaporization and condensation more donor atoms can coordinate to the same
reaction. (Section 20.1) 1liquid Δ gas2. (Section 11.4) metal ion. (Section 23.3)
G-10 GLOSSARY
polymer A large molecule of high molecular mass, protium The most common isotope of hydrogen. rate law An equation that relates the reaction rate
formed by the joining together, or polymerization, (Section 22.2) to the concentrations of reactants (and sometimes
of a large number of molecules of low molecular of products also). (Section 14.3)
proton A positively charged subatomic particle
mass. The individual molecules forming the found in the nucleus of an atom. (Section 2.3) reactant A starting substance in a chemical
polymer are called monomers. (Sections 12.1
pure substance Matter that has a fixed reaction; it appears to the left of the arrow in a
and 12.8)
composition and distinct properties. (Section 1.2) chemical equation. (Section 3.1)
polypeptide A polymer of amino acids that has
pyrometallurgy A process in which heat converts reaction mechanism A detailed picture, or
a molecular weight of less than 10,000.
a mineral in an ore from one chemical form to model, of how the reaction occurs; that is, the
(Section 24.7)
another and eventually to the free metal. order in which bonds are broken and formed and
polyprotic acid A substance capable of (Section 23.2) the changes in relative positions of the atoms as
dissociating more than one proton in water; the reaction proceeds. (Section 14.6)
H 2SO 4 is an example. (Section 16.6) qualitative analysis The determination of the
presence or absence of a particular substance in a reaction order The power to which the
polysaccharide A substance made up of many concentration of a reactant is raised in a rate law.
mixture. (Section 17.7)
monosaccharide units joined together. (Section 14.3)
(Section 24.8) quantitative analysis The determination of the
amount of a given substance that is present in a reaction quotient (Q) The value that is obtained
porphyrin A complex derived from the porphine when concentrations of reactants and products
sample. (Section 17.7)
molecule. (Section 23.3) are inserted into the equilibrium expression. If the
quantum The smallest increment of radiant concentrations are equilibrium concentrations,
positron A particle with the same mass as an
energy that may be absorbed or emitted; the Q = K; otherwise, Q ≠ K. (Section 15.6)
electron but with a positive charge, +10e, or b+ .
magnitude of radiant energy is hn. (Section 6.2)
(Section 21.1) reaction rate A measure of the decrease in
quaternary structure The structure of a protein
positron emission A nuclear decay process where concentration of a reactant or the increase in
resulting from the clustering of several individual
a positron, a particle with the same mass as an concentration of a product with time. (Section 14.2)
protein chains into a final specific shape.
electron but with a positive charge, symbol +10e, or redox (oxidation–reduction) reaction A reaction
(Section 24.7)
b+ is emitted from the nucleus. (Section 21.1) in which certain atoms undergo changes in
racemic mixture A mixture of equal amounts of
potential energy The energy that an object oxidation states. The substance increasing
the dextrorotatory and levorotatory forms of a
possesses as a result of its composition or its in oxidation state is oxidized; the substance
chiral molecule. A racemic mixture will not rotate
position with respect to another object. (Section 5.1) decreasing in oxidation state is reduced. (Section
the plane of polarized light. (Section 23.4)
4.4; Chapter 20: Introduction)
precipitate An insoluble substance that forms in, rad A measure of the energy absorbed from
and separates from, a solution. (Section 4.2) reducing agent, or reductant The substance that
radiation by tissue or other biological material;
is oxidized and thereby causes the reduction of
precipitation reaction A reaction that occurs 1 rad = transfer of 1 * 10 -2 J of energy per
some other substance in an oxidation–reduction
between substances in solution in which one of the kilogram of material. (Section 21.9)
reaction. (Section 20.1)
products is insoluble. (Section 4.2) radial probability function The probability that
reduction A process in which a substance gains
precision The closeness of agreement among the electron will be found at a certain distance
one or more electrons. (Section 4.4)
several measurements of the same quantity; the from the nucleus. (Section 6.6)
reproducibility of a measurement. (Section 1.5) radioactive Possessing radioactivity, the rem A measure of the biological damage caused by
pressure A measure of the force exerted on a unit spontaneous disintegration of an unstable atomic radiation; rems = rads * RBE. (Section 21.9)
area. In chemistry, pressure is often expressed in nucleus with accompanying emission of radiation. renewable energy sources Energy such as solar
units of atmospheres (atm) or torr: 760 torr = (Section 2.2; Chapter 21: Introduction) energy, wind energy, and hydroelectric energy
1 atm; in SI units pressure is expressed in pascals radioactive decay chain A series of nuclear derived from essentially inexhaustible sources.
(Pa). (Section 10.2) reactions that begins with an unstable nucleus and (Section 5.8)
pressure–volume (PV) work Work performed terminates with a stable one. Also called nuclear representative (main-group) element An element
by expansion of a gas against a resisting pressure. disintegration series. (Section 21.2) from within the s and p blocks of the periodic
(Section 5.3) radioisotope An isotope that is radioactive; that table (Figure 6.29). (Section 6.9)
primary cell A voltaic cell that cannot be is, it is undergoing nuclear changes with emission resonance structures (resonance
recharged. (Section 20.7) of radiation. (Section 21.1) forms) Individual Lewis structures in cases where
primary structure The sequence of amino acids radionuclide A radioactive nuclide. (Section 21.1) two or more Lewis structures are equally good
along a protein chain. (Section 24.7) descriptions of a single molecule. The resonance
radiotracer A radioisotope that can be used to structures in such an instance are “averaged”
primitive lattice A crystal lattice in which the trace the path of an element in a chemical system. to give a more accurate description of the real
lattice points are located only at the corners of (Section 21.5) molecule. (Section 8.6)
each unit cell. (Section 12.2) Raoult’s law A law stating that the partial pressure
reverse osmosis The process by which water
probability density 1c22 A value that represents of a solvent over a solution, Psolution, is given by
molecules move under high pressure through
the probability that an electron will be found at a the vapor pressure of the pure solvent, P °solvent,
a semipermeable membrane from the more
given point in space. Also called electron density. times the mole fraction of a solvent in the solution,
concentrated to the less concentrated solution.
(Section 6.5) X solvent: Psolution = X solventP °solvent. (Section 13.5)
(Section 18.4)
product A substance produced in a chemical rare earth element See lanthanide element.
reversible process A process that can go back and
reaction; it appears to the right of the arrow in a (Sections 6.8 and 6.9)
forth between states along exactly the same path;
chemical equation. (Section 3.1)
rate constant A constant of proportionality a system at equilibrium is reversible if equilibrium
property A characteristic that gives a sample of between the reaction rate and the concentrations of can be shifted by an infinitesimal modification of a
matter its unique identity. (Section 1.1) reactants that appear in the rate law. (Section 14.3) variable such as temperature. (Section 19.1)
protein A biopolymer formed from amino acids. rate-determining step The slowest elementary ribonucleic acid (RNA) A polynucleotide in which
(Section 24.7) step in a reaction mechanism. (Section 14.6) ribose is the sugar component. (Section 24.10)
GLOSSARY G-11
root-mean-square (rms) speed 1M 2 The square SI units The preferred metric units for use in standard enthalpy change 1 �H ° 2 The change in
root of the average of the squared speeds of the gas science. (Section 1.4) enthalpy in a process when all reactants and products
molecules in a gas sample. (Section 10.7) smectic liquid crystalline phase A liquid crystal are in their stable forms at 1 atm pressure and a
rotational motion Movement of a molecule as in which the molecules are aligned along their specified temperature, commonly 25 °C. (Section 5.7)
though it is spinning like a top. (Section 19.3) long axes and arranged in sheets, with the ends of standard enthalpy of formation 1 �H f° 2 The
salinity A measure of the salt content of seawater, the molecules aligned. There are several different change in enthalpy that accompanies the
brine, or brackish water. It is equal to the mass kinds of smectic phases. (Section 12.8) formation of one mole of a substance from its
in grams of dissolved salts present in 1 kg of elements, with all substances in their standard
solid Matter that has both a definite shape and a
seawater. (Section 18.3) states. (Section 5.7)
definite volume. (Section 1.2)
salt An ionic compound formed by replacing one standard free energy of formation 1 �G f° 2 The
solubility The amount of a substance that
or more hydrogens of an acid by other cations. change in free energy associated with the
dissolves in a given quantity of solvent at a
(Section 4.3) formation of a substance from its elements under
given temperature to form a saturated solution.
standard conditions. (Section 19.5)
saponification Hydrolysis of an ester in the (Sections 4.2 and 13.2)
presence of a base. (Section 24.4) standard hydrogen electrode (SHE) An
solubility-product constant (solubility product)
1Ksp 2 An equilibrium constant related to the
electrode based on the half-reaction
saturated solution A solution in which 2 H + 11 M2 + 2 e - ¡ H 211 atm2. The
undissolved solute and dissolved solute are in equilibrium between a solid salt and its ions in
standard electrode potential of the standard
equilibrium. (Section 13.2) solution. It provides a quantitative measure of the
hydrogen electrode is defined as 0 V. (Section 20.4)
solubility of a slightly soluble salt. (Section 17.4)
scientific law A concise verbal statement or a standard molar entropy 1S ° 2 The entropy value
mathematical equation that summarizes a wide solute A substance dissolved in a solvent to form a
for a mole of a substance in its standard state.
range of observations and experiences. (Section 1.3) solution; it is normally the component of a solution
(Section 19.4)
present in the smaller amount. (Section 4.1)
scientific method The general process of standard reduction potential 1E°red 2 The
advancing scientific knowledge by making solution A mixture of substances that has a potential of a reduction half-reaction under
experimental observations and by formulating uniform composition; a homogeneous mixture. standard conditions, measured relative to
hypotheses, theories, and laws. (Section 1.3) (Section 1.2) the standard hydrogen electrode. A standard
secondary cell A voltaic cell that can be solution alloy A homogeneous alloy, where two reduction potential is also called a standard
recharged. (Section 20.7) or more elements are distributed randomly and electrode potential. (Section 20.4)
secondary structure The manner in which a uniformly throughout the solid. (Section 12.3) standard solution A solution of known
protein is coiled or stretched. (Section 24.7) solvation The clustering of solvent molecules concentration. (Section 4.6)
second law of thermodynamics A statement around a solute particle. (Section 13.1) standard temperature and pressure (STP)
of our experience that there is a direction to solvent The dissolving medium of a solution; it is Defined as 0 °C and 1 atm pressure; frequently used
the way events occur in nature. When a process normally the component of a solution present in as reference conditions for a gas. (Section 10.4)
occurs spontaneously in one direction, it is the greater amount. (Section 4.1) starch The general name given to a group of
specific heat 1Cs 2 The heat capacity of 1 g of a
nonspontaneous in the reverse direction. It is polysaccharides that acts as energy-storage
possible to state the second law in many different substances in plants. (Section 24.8)
substance; the heat required to raise the temperature
forms, but they all relate back to the same idea
of 1 g of a substance by 1 °C. (Section 5.5) state function A property of a system that is
about spontaneity. One of the most common
statements found in chemical contexts is that spectator ions Ions that go through a reaction determined by its state or condition and not by
in any spontaneous process the entropy of the unchanged and that appear on both sides of the how it got to that state; its value is fixed when
universe increases. (Section 19.2) complete ionic equation. (Section 4.2) temperature, pressure, composition, and physical
form are specified; P, V, T, E, and H are state
second-order reaction A reaction in which the spectrochemical series A list of ligands arranged functions. (Section 5.2)
overall reaction order (the sum of the concentration- in order of their abilities to split the d-orbital
term exponents) in the rate law is 2. (Section 14.4) energies (using the terminology of the crystal-field states of matter The three forms that matter can
model). (Section 23.6) assume: solid, liquid, and gas. (Section 1.2)
semiconductor A material that has electrical
conductivity between that of a metal and that of an spectrum The distribution among various stereoisomers Compounds possessing the same
insulator. (Section 12.7) wavelengths of the radiant energy emitted or formula and bonding arrangement but differing in
absorbed by an object. (Section 6.3) the spatial arrangements of the atoms. (Section 23.4)
sigma 1S 2 bond A covalent bond in which
spin magnetic quantum number 1ms 2 A quantum
electron density is concentrated along the stoichiometry The relationships among the
internuclear axis. (Section 9.6) quantities of reactants and products involved in
number associated with the electron spin; it may
chemical reactions. (Chapter 3: Introduction)
sigma 1S 2 molecular orbital A molecular orbital have values of + 12 or - 21 . (Section 6.7)
that centers the electron density about an imaginary stratosphere The region of the atmosphere
spin-pairing energy The energy required to pair
line passing through two nuclei. (Section 9.7) directly above the troposphere. (Section 18.1)
an electron with another electron occupying an
significant figures The digits that indicate the orbital. (Section 23.6) strong acid An acid that ionizes completely in
precision with which a measurement is made; all digits water. (Section 4.3)
spontaneous process A process that is capable
of a measured quantity are significant, including the of proceeding in a given direction, as written or strong base A base that ionizes completely in
last digit, which is uncertain. (Section 1.5) described, without needing to be driven by an outside water. (Section 4.3)
silica Common name for silicon dioxide. source of energy. A process may be spontaneous even strong electrolyte A substance (strong acids,
(Section 22.4) though it is very slow. (Section 19.1) strong bases, and most salts) that is completely
silicates Compounds containing silicon and standard atmospheric pressure Defined as 760 ionized in solution. (Section 4.1)
oxygen, structurally based on SiO 4 tetrahedra. torr or, in SI units, 101.325 kPa. (Section 10.2) structural formula A formula that shows not only
(Section 22.10) standard emf, also called the standard cell the number and kinds of atoms in the molecule
single bond A covalent bond involving one potential 1E ° 2 The emf of a cell when all reagents but also the arrangement (connections) of the
electron pair. (Section 8.3) are at standard conditions. (Section 20.4) atoms. (Section 2.6)
G-12 GLOSSARY
structural isomers Compounds possessing thermoplastic A polymeric material that can valence-bond theory A model of chemical
the same formula but differing in the bonding be readily reshaped by application of heat and bonding in which an electron-pair bond is formed
arrangements of the atoms. (Sections 23.4 pressure. (Section 12.8) between two atoms by the overlap of orbitals on
and 24.2) the two atoms. (Section 9.4)
thermosetting plastic A plastic that, once formed
subatomic particles Particles such as protons, in a particular mold, is not readily reshaped by valence electrons The outermost electrons of
neutrons, and electrons that are smaller than an application of heat and pressure. (Section 12.8) an atom; those that occupy orbitals not occupied
atom. (Section 2.2) in the nearest noble-gas element of lower atomic
third law of thermodynamics A law stating that
subshell One or more orbitals with the same set number. The valence electrons are the ones the
the entropy of a pure, crystalline solid at absolute
of quantum numbers n and l. For example, we atom uses in bonding. (Section 6.8)
zero temperature is zero: S10 K2 = 0.
speak of the 2p subshell 1n = 2, l = 12, which (Section 19.3) valence orbitals Orbitals that contain the
is composed of three orbitals (2px, 2py , and 2pz ). outer-shell electrons of an atom. (Chapter 7:
titration The process of reacting a solution of
(Section 6.5) Introduction)
unknown concentration with one of known
substitutional alloy A homogeneous (solution) concentration (a standard solution). (Section 4.6) valence-shell electron-pair repulsion (VSEPR)
torr A unit of pressure 11 torr = 1 mm Hg2.
alloy in which atoms of different elements model A model that accounts for the geometric
randomly occupy sites in the lattice. (Section 23.6) arrangements of shared and unshared electron
(Section 10.2)
substitution reactions Reactions in which pairs around a central atom in terms of the
transition elements (transition metals) Elements repulsions between electron pairs.
one atom (or group of atoms) replaces another
in which the d orbitals are partially occupied. (Section 9.2)
atom (or group) within a molecule; substitution
(Section 6.8)
reactions are typical for alkanes and aromatic van der Waals equation An equation of state for
hydrocarbons. (Section 24.3) transition state (activated complex) The nonideal gases that is based on adding corrections
particular arrangement of reactant and product to the ideal-gas equation. The correction terms
substrate A substance that undergoes a reaction
molecules at the point of maximum energy in the account for intermolecular forces of attraction and
at the active site in an enzyme. (Section 14.7)
rate-determining step of a reaction. (Section 14.5) for the volumes occupied by the gas molecules
supercritical mass An amount of fissionable
translational motion Movement in which an themselves. (Section 10.9)
material larger than the critical mass.
entire molecule moves in a definite direction.
(Section 21.7) vapor Gaseous state of any substance that
(Section 19.3)
normally exists as a liquid or solid. (Section 10.1)
supersaturated solution A solution containing
transuranium elements Elements that follow
more solute than an equivalent saturated solution. vapor pressure The pressure exerted by a vapor
uranium in the periodic table. (Section 21.3)
(Section 13.2) in equilibrium with its liquid or solid phase.
triple bond A covalent bond involving three (Section 11.5)
surface tension The intermolecular, cohesive
electron pairs. (Section 8.3)
attraction that causes a liquid to minimize its vibrational motion Movement of the atoms
surface area. (Section 11.3) triple point The temperature at which solid, within a molecule in which they move periodically
liquid, and gas phases coexist in equilibrium. toward and away from one another. (Section 19.3)
surroundings In thermodynamics, everything
(Section 11.6)
that lies outside the system that we study. viscosity A measure of the resistance of fluids to
(Section 5.1) tritium The isotope of hydrogen whose nucleus flow. (Section 11.3)
contains a proton and two neutrons.
system In thermodynamics, the portion of the volatile Tending to evaporate readily.
(Section 22.2)
universe that we single out for study. We must be (Section 11.5)
careful to state exactly what the system contains troposphere The region of Earth’s atmosphere
voltaic (galvanic) cell A device in which a
and what transfers of energy it may have with its extending from the surface to about 12 km
spontaneous oxidation–reduction reaction occurs
surroundings. (Section 5.1) altitude. (Section 18.1)
with the passage of electrons through an external
termolecular reaction An elementary reaction Tyndall effect The scattering of a beam of visible circuit. (Section 20.3)
that involves three molecules. Termolecular light by the particles in a colloidal dispersion.
reactions are rare. (Section 14.6) vulcanization The process of cross-linking
(Section 13.6)
polymer chains in rubber. (Section 12.6)
tertiary structure The overall shape of a large uncertainty principle A principle stating there
protein, specifically, the manner in which sections watt A unit of power; 1 W = 1 J>s.
is an inherent uncertainty in the precision with
of the protein fold back upon themselves or (Section 20.5)
which we can simultaneously specify the position
intertwine. (Section 24.7) and momentum of a particle. This uncertainty is wave function A mathematical description of an
theoretical yield The quantity of product that is significant only for particles of extremely small allowed energy state (an orbital) for an electron in
calculated to form when all of the limiting reagent mass, such as electrons. (Section 6.4) the quantum mechanical model of the atom; it is
reacts. (Section 3.7) usually symbolized by the Greek letter c.
unimolecular reaction An elementary reaction
(Section 6.5)
theory A tested model or explanation that that involves a single molecule. (Section 14.6)
satisfactorily accounts for a certain set of wavelength The distance between identical points
unit cell The smallest portion of a crystal that
phenomena. (Section 1.3) on successive waves. (Section 6.1)
reproduces the structure of the entire crystal when
thermochemistry The relationship between repeated in different directions in space. It is the weak acid An acid that only partly ionizes in
chemical reactions and energy changes. repeating unit or building block of the crystal water. (Section 4.3)
(Chapter 5: Introduction) lattice. (Section 12.2) weak base A base that only partly ionizes in
thermodynamics The study of energy and its unsaturated solution A solution containing less water. (Section 4.3)
transformation. (Chapter 5: Introduction) solute than a saturated solution. (Section 13.2) weak electrolyte A substance that only partly
thermonuclear reaction Another name for fusion valence band A band of closely spaced molecular ionizes in solution. (Section 4.1)
reactions; reactions in which two light nuclei are orbitals that is essentially fully occupied by work The movement of an object against some
joined to form a more massive one. (Section 21.8) electrons. (Section 12.7) force. (Section 5.1)
Useful Conversion Factors and Relationships
Length Energy (derived)
SI unit: meter (m) SI unit: Joule (J)
1 km = 0.62137 mi 1J = 1 kg-m2/s2
1 mi = 5280 ft = 0.2390 cal
= 1.6093 km = 1C-V
1 m = 1.0936 yd 1 cal = 4.184 J
1 in. = 2.54 cm (exactly) l eV = 1.602 * 10-19J
1 cm = 0.39370 in.
1 Å = 10-10 m Pressure (derived)
SI unit: Pascal (Pa)
Mass 1 Pa = 1 N/m2
SI unit: kilogram (kg) = 1 kg/m-s2
1 kg = 2.2046 lb 1 atm = 1.01325 * 105Pa
1 lb = 453.59 g = 760 torr
= 16 oz = 14.70 lb/in2
1 amu = 1.660538921 * 10-27 kg 1 bar = 105 Pa
1 torr = 1 mm Hg
Temperature
SI unit: Kelvin (K) Volume (derived)
0 K = -273.15 °C SI unit: cubic meter (m3)
= -459.67 °F 1 L = 10-3 m3
K = °C + 273.15 = 1 dm3
°C = -95 (°F - 32°) = 103 cm3
°F = -59 °C + 32° = 1.0567 qt
1 gal = 4 qt
= 3.7854 L
1 cm3 = 1 mL
1 in3 = 16.4 cm3
Color Chart for Common Elements
Generic metal
Ag Au Br C Ca Cl
Silver Gold Bromine Carbon Calcium Chlorine
Cu F H I K Mg
Copper Fluorine Hydrogen Iodine Potassium Magnesium
N Na O P S Si
Nitrogen Sodium Oxygen Phosphorus Sulfur Silicon

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CHE1401, Chapter 1

H. Eugene LeMay, Jr.

1.1 | The Study of Chemistry

The Atomic and Molecular Perspective of Chemistry

Carbon dioxide Ethylene glycol Aspirin

Give It Some Thought

Why Study Chemistry?

Chemistry Put to Work

Ethylene C2H4 50 Plastics, antifreeze

Lime CaO 45 Paper, cement, steel

Propylene C3H6 35 Plastics

Ammonia NH3 18 Fertilizers

Chlorine Cl2 21 Bleaches, plastics, water purification

Phosphoric acid H3PO4 20 Fertilizers

Sodium hydroxide NaOH 16 Aluminum production, soap

1.2 | Classifications of Matter

Only one kind of atom is in any element. Compounds must have at

▲ Figure 1.5 Molecular comparison of elements, compounds, and mixtures.

*U.S. Geological Survey Circular 285, U.S Department of the Interior.

Water, H2O Hydrogen gas, H2

Hydrogen atom Oxygen atom Water molecule

Oxygen molecule (written O2)

Hydrogen molecule (written H2)

Table 1.3 Comparison of Water, Hydrogen, and Oxygen

Give It Some Thought

NO Does it have a YES

NO Does it contain YES

1.3 | Properties of Matter

Give It Some Thought

Physical and Chemical Changes

Give It Some Thought

*Remsen, Ira, The Principles of Theoretical Chemistry, 1887.

▲ Figure 1.14 Separation of three substances using column chromatography.

1.4 | Units of Measurement

Table 1.4 SI Base Units

Give It Some Thought

Peta P 1015 1 petawatt (PW) = 1 * 1015 wattsa

Tera T 1012 1 terawatt (TW) = 1 * 1012 watts

Giga G 109 1 gigawatt (GW) = 1 * 109 watts

Mega M 106 1 megawatt (MW) = 1 * 106 watts

Kilo k 103 1 kilowatt (kW) = 1 * 103 watts

Deci d 10-1 1 deciwatt (dW) = 1 * 10-1 watt

Centi c 10-2 1 centiwatt (cW) = 1 * 10-2 watt

Milli m 10-3 1 milliwatt (mW) = 1 * 10-3 watt

Micro mb 10-6 1 microwatt 1mW2 = 1 * 10-6 watt

Nano n 10-9 1 nanowatt (nW) = 1 * 10-9 watt

Pico p 10-12 1 picowatt (pW) = 1 * 10-12 watt

Femto f 10-15 1 femtowatt (fW) = 1 * 10-15 watt

Atto a 10-18 1 attowatt (aW) = 1 * 10-18 watt

Zepto z 10-21 1 zeptowatt (zW) = 1 * 10-21 watt

Give It Some Thought

Length and Mass

373 K 100 °C 212 °F Water boils

273 K 0 °C 32 °F Water freezes

Kelvin scale Celsius scale Fahrenheit scale

degree sign 1°2 with temperatures on the Kelvin scale.

1°F - 322 or °F = 1°C2 + 32

These deliver variable volumes Pipette delivers a Volumetric flask contains

▲ Figure 1.19 Common volumetric glassware.

Volume = 12.00 cm23 = 12.0023 cm3 = 8.00 cm3