Frederick A. Bettelheim
William H. Brown
Mary K. Campbell
Shawn O. Farrell
www.cengage.com/chemistry/bettelheim
Chapter 1
Matter, Energy, and Measurement
William H. Brown • Beloit College
Chemistry
Chemistry is the study of matter.
• Matter is anything that has mass and takes up space.
Matter can change from one form to another.
• In a chemical change (chemical reaction), substances
are used up and others formed in their place.
Example: When propane (bottled or LP gas) burns in
air, propane and oxygen are converted to carbon
dioxide and water.
• In a physical change, matter does not lose its identity.
• A common physical change is a change of state.
Example: Ice (solid water) melts to become liquid
water; liquid water boils to become steam (gaseous
water).
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Scientific Method
Fact: A statement based on direct experience
Hypothesis: A statement that is proposed, without actual
proof, to explain a set of facts or their relationship.
Theory: The formulation of an apparent relationship
among certain observed phenomena, which has been
verified to some extent.
• In a sense, a theory is the same as a hypothesis except
that we have a stronger belief in it because more
evidence supports it.
• If, however, we find new evidence that conflicts with
the theory, it must be altered or rejected.
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Scientific Method
Experimentation
Theory
If exp eriments con trad ict the theory,
the th eory may have to b e modified
or d iscarded
A u nifying principle that explains a
body of observation s and th e law s bas ed
on them; suggests n ew experiments
Law
Su mmarizes and exp lains a
w ide range of observation s
Observation
D irect ob servations (facts )
about the b ehavior of matter
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Exponential Notation
Used to represent very large or very small numbers as
powers of 10.
• Examples:
0.00002 is written as 2 x 10-5
2,000,000 is written as 2 x 106
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Metric System
Table 1.1 Base units in the metric system
Length
Volume
Mass
Time
Temperature
Energy
Amount of matter
meter (m)
liter (L)
gram (g)
second (s )
Cels ius (°C)
calorie (cal)
mole (mol)
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Metric System
Table 1.2 The Most Common Metric Prefixes
Prefix
Symbol
Valu e
giga
G
109
mega
M
106
kilo
deci
k
d
103
10-1
cen ti
c
10-2
milli
m
10-3
micro

10-6
nan o
n
10-9
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Significant Figures
1. Nonzero digits are always significant.
• For example 233.1 has four significant figures; 2.3g has
two.
2. Zeros at the beginning of a number are never significant.
For example 0.0055 L has two significant figures.
3. Zeros at the end of a number that contains a decimal point
are always significant. For example 3.00L has three
significant figures. 0.0450 mm also has three.
4. Zeros at the end of of number that contains no decimal
point may or may not be significant. For example $36,000
contains two significant figures and $36,000.00 contains
seven significant figures. 5,000 mL contains one
significant Figure and 5,000. mL contains four.
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Metric & English Systems
Table 1.3 Some Conversion Factors
Length
1 in . = 2.54 cm
1 m = 39.37 in .
1 mile = 1.609 k m
Mass
Volu me
1 oz = 28.35 g
1 qt = 0.946 L
1 lb = 453.6 g
1 gal = 3.785 L
1 kg = 2.205 lb
1 L = 33.81 fl oz
1 g = 15.43 grains 1 fl oz = 29.57 mL
1 L = 1.057 qt
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Mass and Weight
Mass: The quantity of matter in an object.
• Mass is independent of location.
Weight: The result of mass acted upon by gravity.
• Weight depends on location; depends on the force of
gravity at the particular location.
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Temperature
Fahrenheit (F): Defined by setting the normal freezing point
of water at 32°F and the normal boiling point of water at
212°F.
Celsius (C): Defined by setting the normal freezing point of
water at 0°C and the normal boiling point of water at
100°C.
9
°F = _ °C + 32
5
5
°C = _ (°F - 32)
9
Kelvin (K): Zero is the lowest possible temperature; also
called the absolute scale.
• Kelvin degree is the same size as a Celsius degree
• K = °C + 273
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Factor-Label Method
Conversion factor: A ratio of two different units, used as a
multiplier to change from one system or unit to another.
• For example, 1 lb = 453.6 g
• Example: Convert 381 grams to pounds.
381 g x
1 lb
453.6 g
= 0.840 lb
• Example: Convert 1.844 gallons to milliliters.
1.844 gal x
3.785 L 1000. mL = 6980. mL
x
1 gal
1L
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The Three States of Matter
Gas
• Has no definite shape or volume.
• Expands to fill whatever container it is put into.
• Is highly compressible.
Liquid
• Has no definite shape but a definite volume.
• Is only slightly compressible.
Solid
• Has a definite shape and volume.
• Is essentially incompressible.
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Density
Density: The ratio of mass to volume.
m
d=
V
d = den sity
m = mass
V = volume
• The most commonly used units are g/mL for liquids
and solids, and g/L for gases.
• Example: If 73.2 mL of a liquid has a mass of 61.5 g,
what is its density in g/mL?
d=
m
61.5 g
=
V
73.2 mL
= 0.840 g/mL
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Specific Gravity
Specific gravity: The density of a substance compared to
water as a standard.
• Because specific gravity is the ratio of two densities, it
has no units (it is dimensionless).
• Example: The density of copper at 20°C is 8.92 g/mL.
The density of water at this temperature is 1.00 g/mL.
What is the specific gravity of copper?
8.92 g/mL
Specific gravity =
= 8.92
1.00 g/mL
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Energy
Energy: The capacity to do work.
• May be either kinetic energy or potential energy
• The calorie (cal) is the base metric unit of energy.
Kinetic energy (KE): the energy of motion
1
KE = _ mv 2
2
• KE increases as the object’s velocity increases.
• At the same velocity, a heavier object has greater KE.
Potential energy: The energy an object has because of its
position; stored energy
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Energy
• Examples of kinetic energy are mechanical energy, light,
heat, and electrical energy.
• In chemistry, the most important forms of potential
energy are chemical energy and nuclear energy.
• Chemical energy is stored in chemical substances as,
for example, in foods such as carbohydrates and fats.
• It is given off when substances take part in chemical
reactions.
• The law of conservation of energy
• Energy can neither be created nor destroyed.
• Energy can be converted from one form to another.
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Heat and Temperature
Heat is a form of energy.
• Heating refers to the energy transfer process when two
objects of different temperature are brought into
contact.
• Heat energy always flows from the hotter object to the
cooler one until the two have the same temperature.
• Heat is commonly measured in calories (cal), which is
the heat necessary to raise the temperature of 1 g of
liquid water by 1°C.
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Specific Heat
Specific heat (SH): The amount of heat necessary to raise
the temperature of 1.00 g of a substance by 1.00°C.
Table 1.4 Specific Heats of Common Substances
Substance
Water
Ice
Steam
Iron
Aluminum
Copper
Lead
Specific Heat
(cal/g •°C)
1.00
0.48
0.48
0.11
0.22
0.092
0.038
Specific Heat
Substance
(cal/g •°C)
Wood
0.42
Glass
0.22
Rock
0.20
Ethanol
0.59
Methanol
0.61
Ether
0.56
Acetone
0.52
Carbon tetrachloride 0.21
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Specific Heat
The following equation gives the relationship between
specific heat, amount of heat, the mass of an object, and
the change in temperature.
Amount of h eat = sp ecific heat x mass x change in temperature
= S H x m x (T2 - T1 )
• Example: How many calories are required to heat 352 g
of water from 23°C to 95°C?
Amoun t of heat =
1.00 cal
g • °C
x 352 g x (95 - 23)°C
4
= 2.5 x 10 cal = 25 kcal
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Chapter 1 Matter, Energy, and Measurement
End
Chapter 1
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