Chapter 1

Essential Ideas
Chapter 1: Essential Ideas
1.1 Chemistry in Context
1.2 Phases and Classification of Matter
1.3 Physical and Chemical Properties
1.4 Measurements
1.5 Measurement Uncertainty, Accuracy, and
Precision
1.6 Mathematical Treatment of
Measurements
1.1 — Chemistry in
Context
What is Chemistry?
 the study of the composition, properties, and interactions of
matter
The Central Science
 Chemistry is related to many disciplines
The Central Science
 Chemistry is a fundamental part of everyday life.
 metabolism of food
 synthesis of polymers for clothing and plastics
 refining crude oil for gasoline
 forensics … and much more!
 As you study chemistry, you will discover:
 that chemistry is involved in changing the composition of
matter.
 that these changes can be classified and understood by using
basic chemical principles.
 that changes in energy accompany changes in matter.
The Scientific Method
 What is this?
 a logical, problem-solving path to
discovery used by scientists

1. Ask a question or define a

problem.
2. Formulate a hypothesis, a
tentative explanation of
observations.
3. Use experiments, calculations,
and observations to answer the
original question or solve the
original problem.
The Domains of Chemistry
 Macroscopic Domain: Moisture
in the air, icebergs, and the ocean
represent water in the
macroscopic domain.
 Microscopic Domain: At the
molecular level, gas molecules
are far apart and disorganized,
solid water molecules are close
together and organized, and
liquid molecules are close
together and disorganized.
 Symbolic Domain: H2O is the
chemical formula for water, and
(g), (s), and (l) symbolize its
phases.
1.2 — Phases and the
Classification of
Matter
Why does matter matter?
What is Matter?
 Matter is anything that has a mass and
occupies volume.
 What are the phases/states of Matter?
 Solid (s) - rigid
and posses a
definite shape.
 Gas (g) - flows
and takes the
shape of the
container.
 Liquid (l) - takes
both the shape
and volume of the
container.
States of Matter
Mass vs. Weight
 Mass
 is a measure of the amount of
matter in an object.
 Weight
 refers to the force that gravity
exerts on an object

 An object will have the

same mass on both the
earth and the moon, but it
will have a different weight
on each.
The Law of Conservation of Matter
During a physical or chemical change, there is no
detectable change in the total quantity of matter present.
Matter cannot be created or destroyed.

Physical Change
 the phase of a substance changes
 between solid, liquid, gaseous states
Chemical Change
• matter is converted from one type to another
• beer brewing: water, yeast, grains, malt, hops, and sugar are
converted into beer (water, alcohol, carbonation, and flavoring)
Elements
 An element is a
pure substance
that cannot be
broken down into
simpler
substances by
chemical changes.
 There are
currently 118
elements
discovered
 98 of those are
naturally
occurring
Atoms and Molecules
Atoms
 Atoms are the smallest particles of an element that
have the properties of that element.
carbon nitrogen oxygen hydrogen

Molecules
• Molecules are two or more atoms joined together by
forces known as chemical bonds.
• The atoms in a molecule are bound together and the
move around as a unit.

water oxygen
molecule molecule glucose
molecule
Elements and Matter
Elements – Atoms
 Elements consist of a single type of atom.
 Only six elements exist as single atoms in nature.
 He, Ne, Ar, Xe, Kr, and Rn
 What are these?

Elements - Molecules
 Most elements exist as molecules with two or more atoms of that
element are bound together.
Molecules
 Some molecules are elements; they consist of one type
of atom.
 Most molecules consist of atoms of different elements
bound together. These molecules are known as
compounds.

 A water molecule is a compound that consists of two

hydrogen atoms and one oxygen atom held together by water
molecule
chemical bonds.

 A glucose molecule is made from six carbon atoms, glucose

twelve hydrogen atoms, and six oxygen atoms boundmolecule
Classifying Matter
Pure Substances
 Pure substances have constant
composition.
 Pure substances that are elements:
 cannot be broken down into simpler
substances chemically.
 gold (Au), oxygen (O2), iron (Fe) for example.
 Pure substances that are compounds:
 can be broken down into simpler substances
or into its elements by chemical changes.
 H2O can be broken down into H2 and O2.
 have physical properties that are different
from the physical properties of its elements.
Classifying Matter
Mixtures
 A mixture is a combination of
substances.
 two or more compounds, two or more
elements, or a combination
 can be separated by physical methods.
Classifying Matter
Determine the type of matter
 NaCl  Sucrose (Sugar)
 P.S.  compound  P.S.  compound
 Honey  Saltwater solution
 P.S.  compound  P.S.  compound
 Shampoo
 Mix  homogenous
 Peroxide
 Carbon  P.S.  compound
 P.S.  element  Apple Juice
 Rocky Road Ice Cream  Mix  homogenous
 Mix  heterogenous
 Sodium
 Coffee
 P.S.  element
 Mix  homogenous
 Water
 P.S.  compound
 Copper
1.3 — Physical and
Chemical Properties
Properties of Matter:

Physical Chemical
 are characteristics of matter  The ability to change
that are not associated with from one type of matter
Density
a change in chemical Reactivity
into another type of
composition matter arises from
 Density
Conductivity
 Flammability chemical
Smell properties
Color

 Flammability
Hardness Color
 Conductivity
 Reactivity
 Viscosity Hardness Acidity  Acidity Shape
 Smell Combustibility Phases of
 Combustibility
 Shape  Ability to oxidize

Matter
Phases of matter
 SolubilitySolubility Ability to Oxidize
Intensive vs. Extensive Properties
Intensive Properties
 do not depend on the amount of substance present.
 examples: density, temperature
Extensive
 depend on the amount of substance present.
 examples: mass, volume, heat

A drop of cooking oil and a pot of cooking oil

may have the same temperature (intensive)
but the drop will not burn you as badly as an
entire pot would (extensive) because the pot
contains more heat.
Quick Review
 What is Chemistry?
 The study of matter
 What is Matter?
 Something with mass and volume
 How do you talk about Matter?
 Describe mass and volume
1.4 — Measurements
Why do we measure things?
Measurements
 provide the information that is the basis of most
hypotheses, theories, and laws in science

 Number—size or magnitude
 Units—a standard for measurement
 A measurement without a unit is meaningless!

The diameter of
The diameter of
the earth is
DNA 23
12,760,000
nanometers.
meters.
SI Units
 International System of Units (SI units)
 have been in use by US National Institute of Standards and
Technology (NIST) since 1964.
 Is not used in everyday life here in America, but it is the standard
almost everywhere else in the world.

Property Unit Symbol

Length meter m
Mass kilogram kg
Time second s
Temperature Kelvin or K or °C
Celsius
Amount mole mol
Current ampere amp
SI Units
 Fractional or multiple SI units are named by using a
prefix and the name of the base unit.
Prefix Symbol Factor 1 cm = 10-2 m
femto- f 10-15 or
pico- p 10-12 100 cm = 1 m

nano- n 10-9
micro- μ 10-6 1 mL = 10-3 L
milli- m 10-3 or
1000 mL = 1 L
centi- c 10-2
deci- d 10-1
kilo- k 10+3 1 μg = 10-6 g
or
giga- G 10+9 1,000,000 μg = 1 g
tera- T 10+12
What are basic SI units used in
Chemistry?
 For Length  Volume
 Meters (m)  Liters (L)
 nm, km, cm  mL, µL
 1 mL = 1 cm3
 For Mass
 Grams (g)  Time
 kg, mg, µg  Seconds (s or sec)
 ms, ns, ps, µs
 For Temperature  Speed = distance/time (m/s)
 Kelvin (K) or Celsius)
 K = 273 + ℃
History of the Kilogram
 kilogram was originally defined in 1795 as the
mass of a liter of water
 From 1879 to this May (2019) a platinum-iridium
cylinder, known as the
International Prototype of the Kilogram (IPK) was the
standard of the unit of mass for the metric system.
 It is/was kept in a triple-locked vault in the outskirts of
Paris.

 In November 2018, the international scientific

community voted to redefine the kilogram,
freeing it from its embodiment in one golf-ball-
sized artifact, and basing it instead on a
constant of nature.
 Planck’s constant
 defining the kilogram in terms of the second and the meter
Density
 the ratio of the mass of sample to its volume
 mass per unit volume
 SI units: Density in chemistry usually: , ,
How density differs?
1. Mercury has a density of 13.535 g/cm3 and a volume if
32.0 mL. What is the mass?
2. Ethylene glycol had a density of 1.11 g/mL and a mass of
1850 grams. What is the volume?
3. The density of seawater at 15°C is 1.025 g/cm3. What is
this in kilograms per cubic meter?
4. The density of air is 1.18 x 10 -3 g/mL. What is the volume
of 15.5 g of air?
5. What is the mass of a compound if the density is 0.970
g/mL and has a volume of 4.06 cm3?
6. A piece of plastic (shown below) has a mass of 50.2 g.
Determine the density.

7. A graduated cylinder with water has a volume at 200 mL.

A piece of copper is added and the water level increases
to 415 mL. The density of copper is 8.96 g/mL. What is
1.5—Uncertainty,
Accuracy, and
Precision
Significant figures
Exact Numbers
 Counting is the only type of measurement that is free
from uncertainty
 A count is an exact number
 Examples?
 12 eggs, 3 cars, 12 atoms

 Defined quantities are exact

 We know exactly how many inches are in a foot
 Centimeters in an inch
 Grams in a kilogram
Measurements and Uncertainty
 Measurements are not exact
 What are the practical limitations of the process?
 When making a measurement, estimate one uncertain
digit.

The meniscus is between 21 mL and 22 mL.

The next digit must be estimated.
What should the final measurement be?
Uncertainty, Accuracy, and Precision
 Every measurement has some uncertainty
Significant Figures
 All the digits in a measurement, including the
uncertain last digit, are called significant figures or
significant digits
 All nonzero digits1267
four significant figures

captive zero
 Captive zeros
3085
four significant figures

 Trailing zeros after a decimal

leading point
zeros (not significant)

0.008020
four significant figures
Significant Figures
 trailing zeros not after a decimal?
ambiguous

3200
significant
 avoid ambiguity by using scientific notation
3.2 x 103 3.20 x 103
two significant figures three significant figures

 How many significant figures?

0.082057 8.3145x103 96485
6.022x1023
Significant Figures in Calculations

Multiplication/
Addition/Subtraction Division
 Result depends upon the  Result depends upon the
value with the least value with the least
number of decimal number of significant
places.
13.4637 + 1.62 = 15.08
digits.
13.4637 x 1.62 = 21.8
four
significant three significant
figures figures
13.4637
+ 1.62
15.0837 0.082057 x 298.15 = 24.465

limited to five significant

two decimal figures
places
Rounding
• After the number of required significant figures is determined,
the calculated value must be rounded properly.
• dropped digit > 5  round up
0.028673 to three sig figs  0.0287

• dropped digit < 5  round down

18.3384 to three sig figs  18.3

• dropped digit = 5, keep retained digit even  must round up

6.8752 to three sig figs  6.88

• dropped digit = 5, keep retained digit even  must round down

92.05 to three sig figs  92.0
Zero is an even number.
Accuracy and Precision
 Accurate measurements yield results close to the true
or accepted value.
 Precise measurements yield similar results when
repeated
1.6 — Mathematical
Treatment of
Measurement Results
Dimensional Analysis
 The units associated with measurements must be
subjected to the same mathematical treatment as the
numerical values of those measurements.
 A conversion factor is a ratio of two equivalent quantities
expressed with different measurement units.

or
1 𝑖𝑛 2.54 𝑐𝑚
 One inch is equal to 2.54 centimeters.2.54 𝑐𝑚 1 𝑖𝑛

1000 𝑚𝐿 or
1𝐿 1000 𝑚𝐿
 One liter is equal to 1000 milliliters. 1𝐿

or
1𝑙𝑏 453.59 𝑔
 One pound is equal to 453.59 grams. 453.59 𝑔 1𝑙𝑏
Dimensional Analysis
 In a calculation, arrange the conversion factor so that
the original units cancel out and the desired units
remain.
 Example:
 A basketball player’s vertical jump is 34 inches. What is the
player’s vertical jump in centimeters?
 Given: 34in Want: cm
1 in = 2.54 cm
2.54 𝑐𝑚
2.54 𝑐𝑚 34 𝑖𝑛 × =86 𝑐𝑚 The inches cancel
1 𝑖𝑛
1 𝑖𝑛 and units of
two two centimeters are left.
significant significant
figures figures
exact
value
Convert the Following
• 25 inches to meters

• 10 m to mm

• 2.5 kg to lbs

• 2 mg to kg

• 515 km to cm
 Note: Exact
Example numbers/conversions do not
limit your significant figures.
A world record for the men’s marathon was set by
Dennis Kimetto of Kenya on September 28, 2014. He ran
the race in 2:02:57. The official marathon length is
26.219 miles. What speed, in meters per second, did
Dennis Kimetto run for that marathon? What was his
speed
Know:in2:02:57
miles per hour?Want: speed in meters/second and miles/hour
26.219 miles
2 : 02 : 57=2 h𝑟 + 2𝑚𝑖𝑛+ 57 𝑠=7377 𝑠
1 in = 2.54 cm
26.219𝑚𝑖 5280 𝑓𝑡 𝑓𝑡
1 mi = 5280 ft 𝑆𝑝𝑒𝑒𝑑 = × =18.7659
7377 𝑠 1 𝑚𝑖 𝑠
1 hr = 60 min
1 ft = 12 in 𝑓𝑡 12𝑖𝑛 2.54 𝑐𝑚 1𝑚 𝑚
18.7659 × × × =5.720
𝑠 1 𝑓𝑡 1𝑖𝑛 100 𝑐𝑚 𝑠

You convert to miles/hour. 54

Example
The surface area of one hexagonal gold mirror
on NASA’s James Webb Space Telescope is
12.01 ft2. What is the area in meters squared?

and and
12𝑖𝑛 2.54 𝑐𝑚 100 𝑐𝑚
1 𝑓𝑡 1 𝑖𝑛 1𝑚

2 12 𝑖𝑛 12 𝑖𝑛 2 Keep extra digits until

12.01 𝑓𝑡 =12.01 𝑓𝑡 ∙ 𝑓𝑡 × × =1729.44 𝑖𝑛 the end of the
1 𝑓𝑡 1 𝑓𝑡 calculation to avoid
rounding errors.

2 2.54 𝑐𝑚 2.54 𝑐𝑚 2
1729.44 𝑖𝑛 =1729.44 𝑖𝑛 ∙ 𝑖𝑛 × × =11157.6 5 𝑐𝑚
1𝑖𝑛 1𝑖𝑛

2 1𝑚 1𝑚 2
11 157 . 6 5 𝑐𝑚 =11 157 . 6 5 𝑐𝑚 ∙𝑐𝑚 × × =1.116 𝑚
100 𝑐𝑚 100 𝑐𝑚 12.01 has four
significant figures, so
Note: Exact the answer must have
numbers/conversions do not four significant figures. 55
limit your significant figures.
Practice
The volume of one of the DNA crystals shown in the
photograph is 0.144 nm3. What is the volume in
milliliters?
Practice
A factory has a solid copper sphere that needs to be drawn
into a wire. The mass of the copper sphere is 76.5 kg. The
copper needs to be drawn into a wire with a diameter of 9.50
mm. What length of wire, in meters, can be produced? Copper
has a density of 8.96 g/cm3.
Practice
Golden eagles in the alps hunt from high altitudes. While
hunting, eagles can reach diving speeds of up to 322 km/h.
Only peregrine falcons can dive faster – at speeds of up to
389 km/h. How fast do each of these bird's dive in miles per
hour? In meters per second?
Golden Eagle Peregrine Falcon

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