Chapter one lecture slides for AP Chemistry Matter and Measurement

1. Chemistry

2. Chemistry
• is the study of properties of materials and
changes that they undergo.

3. Chemistry
• is the study of properties of materials and
changes that they undergo.
• can be applied to all aspects of life (e.g.,
development of pharmaceuticals, leaf color
change in fall, etc.).

4. The Atomic and Molecular
Perspective of Chemistry

5. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.

6. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:

7. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:
• is the physical material of the universe.

8. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:
• is the physical material of the universe.
• has mass.

9. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:
• is the physical material of the universe.
• has mass.
• occupies space.

10. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:
• is the physical material of the universe.
• has mass.
• occupies space.
• ~100 elements constitute all matter.

11. The Atomic and Molecular
Perspective of Chemistry
Chemistry involves the study of the properties and
the behavior of matter.
Matter:
• is the physical material of the universe.
• has mass.
• occupies space.
• ~100 elements constitute all matter.
• A property is any characteristic that allows us to
recognize a particular type of matter and to
distinguish it from other types of matter.

12. Elements

13. Elements

14. Elements
• are made up of unique atoms, the building
blocks of matter.

15. Elements
• are made up of unique atoms, the building
blocks of matter.
• Names of the elements are derived from a
wide variety of sources (e.g., Latin or
Greek, mythological characters, names of
people or places).

16. Elements
• are made up of unique atoms, the building
blocks of matter.
• Names of the elements are derived from a
wide variety of sources (e.g., Latin or
Greek, mythological characters, names of
people or places).
• Memorize element symbols

17. Molecules

18. Molecules

19. Molecules
• are combinations of atoms held together in
specific shapes.

20. Molecules
• are combinations of atoms held together in
specific shapes.
• Macroscopic (observable) properties of matter
relate to submicroscopic realms of atoms.

21. Molecules
• are combinations of atoms held together in
specific shapes.
• Macroscopic (observable) properties of matter
relate to submicroscopic realms of atoms.
• Properties relate to composition (types of atoms
present) and structure (arrangement of atoms)
present.

22. 1.2 Classifications of Matter
Matter is classified by state (solid, liquid or
gas) or by composition (element, compound
or mixture).

23. States of Matter

24. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.

25. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.

26. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.
Properties described on the macroscopic level:

27. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.
Properties described on the macroscopic level:
• gas (vapor): no fixed volume or shape, conforms
to shape of container, compressible.

28. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.
Properties described on the macroscopic level:
• gas (vapor): no fixed volume or shape, conforms
to shape of container, compressible.
• liquid: volume independent of container, no fixed
shape, incompressible.

29. States of Matter
Solids, liquids and gases are the three forms of
matter called the states of matter.
Properties described on the macroscopic level:
• gas (vapor): no fixed volume or shape, conforms
to shape of container, compressible.
• liquid: volume independent of container, no fixed
shape, incompressible.
• solid: volume and shape independent of
container, rigid, incompressible.

30. States of Matter

31. States of Matter
Properties described on the molecular level:

32. States of Matter
Properties described on the molecular level:
• gas: molecules far apart, move at high
speeds, collide often.

33. States of Matter
Properties described on the molecular level:
• gas: molecules far apart, move at high
speeds, collide often.
• liquid: molecules closer than gas, move
rapidly but can slide over each other.

34. States of Matter
Properties described on the molecular level:
• gas: molecules far apart, move at high
speeds, collide often.
• liquid: molecules closer than gas, move
rapidly but can slide over each other.
• solid: molecules packed closely in definite
arrangements.

35. Pure Substances

36. Pure Substances
Pure substances:

37. Pure Substances
Pure substances:
• are matter with fixed compositions and
distinct proportions.

38. Pure Substances
Pure substances:
• are matter with fixed compositions and
distinct proportions.
• are elements (cannot be decomposed into
simpler substances, i.e. only one kind of
atom) or compounds (consist of two or
more elements).

39. Mixtures

40. Mixtures
• are a combination of two or more pure
substances.

41. Mixtures
• are a combination of two or more pure
substances.
• Each substance retains its own identity.

42. Elements

43. Elements
• There are 116 known elements.

44. Elements
• There are 116 known elements.
• They vary in abundance.

45. Elements
• There are 116 known elements.
• They vary in abundance.
• Each is given a unique name and is
abbreviated by a chemical symbol.

46. Elements
• There are 116 known elements.
• They vary in abundance.
• Each is given a unique name and is
abbreviated by a chemical symbol.
• they are organized in the periodic table.

47. Elements
• There are 116 known elements.
• They vary in abundance.
• Each is given a unique name and is
abbreviated by a chemical symbol.
• they are organized in the periodic table.
• Each is given a one- or two-letter symbol
derived from its name.

48. Compounds

49. Compounds
• Compounds are combinations of elements.

50. Compounds
• Compounds are combinations of elements.
Example: The compound H2O is a
combination of the elements H and O.

51. Compounds
• Compounds are combinations of elements.
Example: The compound H2O is a
combination of the elements H and O.
• The opposite of compound formation is
decomposition.

52. Compounds
• Compounds are combinations of elements.
Example: The compound H2O is a
combination of the elements H and O.
• The opposite of compound formation is
decomposition.
• Compounds have different properties than their
component elements (e.g., water is liquid, but
hydrogen and oxygen are both gases at the same
temperature and pressure).

53. Law of Constant (Definite)
Proportions

54. Law of Constant (Definite)
Proportions
(Proust): A compound always consists of the
same combination of elements (e.g., water
is always 11% H and 89% O).

56. Mixtures

57. Mixtures
• A mixture is a combination of two or more pure
substances.

58. Mixtures
• A mixture is a combination of two or more pure
substances.
• Each substance retains its own identity; each
substance is a component of the mixture.

59. Mixtures
• A mixture is a combination of two or more pure
substances.
• Each substance retains its own identity; each
substance is a component of the mixture.
• Mixtures have variable composition.

60. Mixtures
• A mixture is a combination of two or more pure
substances.
• Each substance retains its own identity; each
substance is a component of the mixture.
• Mixtures have variable composition.
• Heterogeneous mixtures do not have uniform
composition, properties, and appearance, e.g.,
sand.

61. Mixtures
• A mixture is a combination of two or more pure
substances.
• Each substance retains its own identity; each
substance is a component of the mixture.
• Mixtures have variable composition.
• Heterogeneous mixtures do not have uniform
composition, properties, and appearance, e.g.,
sand.
• Homogeneous mixtures are uniform throughout,
e.g., air; they are solutions.

63. 1.3 Properties of Matter

64. 1.3 Properties of Matter
Each substance has a unique set of physical and
chemical properties.

65. 1.3 Properties of Matter
Each substance has a unique set of physical and
chemical properties.
• Physical properties are measured without
changing the substance (e.g., color, density, odor,
melting point, etc.).

66. 1.3 Properties of Matter
Each substance has a unique set of physical and
chemical properties.
• Physical properties are measured without
changing the substance (e.g., color, density, odor,
melting point, etc.).
• Chemical properties describe how substances
react or change to form different substances (e.g.,
hydrogen burns in oxygen).

67. 1.3 Properties of Matter

68. 1.3 Properties of Matter
Properties may be categorized as intensive or
extensive.

69. 1.3 Properties of Matter
Properties may be categorized as intensive or
extensive.
• Intensive properties do not depend on the
amount of substance present (e.g., temperature,
melting point etc.).

70. 1.3 Properties of Matter
Properties may be categorized as intensive or
extensive.
• Intensive properties do not depend on the
amount of substance present (e.g., temperature,
melting point etc.).
• Extensive properties depend on the quantity
of substance present (e.g., mass, volume etc.).

71. 1.3 Properties of Matter
Properties may be categorized as intensive or
extensive.
• Intensive properties do not depend on the
amount of substance present (e.g., temperature,
melting point etc.).
• Extensive properties depend on the quantity
of substance present (e.g., mass, volume etc.).
• Intensive properties give an idea of the
composition of a substance whereas extensive
properties give an indication of the quantity of
substance present.

72. Physical and Chemical Changes

73. Physical and Chemical Changes
• Physical change: substance changes physical
appearance without altering its identity (e.g.,
changes of state).

74. Physical and Chemical Changes
• Physical change: substance changes physical
appearance without altering its identity (e.g.,
changes of state).
• Chemical change (or chemical reaction):
substance transforms into a chemically different
substance (i.e. identity changes, e.g.,
decomposition of water, explosion of nitrogen
triiodide).

75. Separation of Mixtures

76. Separation of Mixtures
Key: separation techniques exploit differences
in properties of the components.

77. Separation of Mixtures
Key: separation techniques exploit differences
in properties of the components.
• Filtration: remove solid from liquid.

78. Separation of Mixtures
Key: separation techniques exploit differences
in properties of the components.
• Filtration: remove solid from liquid.
• Distillation: boil off one or more
components of the mixture.

79. Separation of Mixtures
Key: separation techniques exploit differences
in properties of the components.
• Filtration: remove solid from liquid.
• Distillation: boil off one or more
components of the mixture.
• Chromatography: exploit solubility of
components.

80. The Scientific Method

81. The Scientific Method
The scientific method provides guidelines for the
practice of science.

82. The Scientific Method
The scientific method provides guidelines for the
practice of science.
• Collect data (observe, experiment, etc.).

83. The Scientific Method
The scientific method provides guidelines for the
practice of science.
• Collect data (observe, experiment, etc.).
• Look for patterns, try to explain them, and
develop a hypothesis or tentative explanation.

84. The Scientific Method
The scientific method provides guidelines for the
practice of science.
• Collect data (observe, experiment, etc.).
• Look for patterns, try to explain them, and
develop a hypothesis or tentative explanation.
• Test hypothesis, then refine it.

85. The Scientific Method
The scientific method provides guidelines for the
practice of science.
• Collect data (observe, experiment, etc.).
• Look for patterns, try to explain them, and
develop a hypothesis or tentative explanation.
• Test hypothesis, then refine it.
• Bring all information together into a scientific
law (concise statement or equation that
summarizes tested hypotheses).

86. The Scientific Method
The scientific method provides guidelines for the
practice of science.
• Collect data (observe, experiment, etc.).
• Look for patterns, try to explain them, and
develop a hypothesis or tentative explanation.
• Test hypothesis, then refine it.
• Bring all information together into a scientific
law (concise statement or equation that
summarizes tested hypotheses).
• Bring hypotheses and laws together into a
theory. A theory should explain general
principles.

87. 1.3
Steps in the Scientific Method

88. 1.3
Steps in the Scientific Method

89. 1.3
Steps in the Scientific Method

90. 1.3
Steps in the Scientific Method

91. 1.3
Steps in the Scientific Method

92. 1.3
Steps in the Scientific Method

93. Units of
Measurement

94. SI Units

95. SI Units
• Système International d’Unités

96. SI Units
• Système International d’Unités
• Uses a different base unit for each quantity

97. Metric System
Prefixes convert the base units into units that
are appropriate for the item being measured.

98. PRACTICE EXERCISE
(a) What decimal fraction of a second is
a picosecond, ps?
(b) Express the measurement 6.0 ×
103 m using a prefix to replace the
power of ten.
(c) Use exponential notation to express
3.76 mg in grams.

99. PRACTICE EXERCISE
(a) What decimal fraction of a second is
a picosecond, ps?
(b) Express the measurement 6.0 ×
103 m using a prefix to replace the
power of ten.
(c) Use exponential notation to express
3.76 mg in grams.
Answers: (a) 10–12 second, (b) 6.0 km, (c) 3.76 × 10–3 g

100. Volume

101. Volume
• The most commonly
used metric units for
volume are the liter (L)
and the milliliter (mL).

102. Volume
• The most commonly
used metric units for
volume are the liter (L)
and the milliliter (mL).
□ A liter is a cube 1 dm
long on each side.

103. Volume
• The most commonly
used metric units for
volume are the liter (L)
and the milliliter (mL).
□ A liter is a cube 1 dm
long on each side.
□ A milliliter is a cube 1 cm
long on each side.

104. Temperature:
A measure of the
average kinetic
energy of the
particles in a
sample.

105. Temperature

106. Temperature
• In scientific
measurements, the
Celsius and Kelvin
scales are most often
used.

107. Temperature
• In scientific
measurements, the
Celsius and Kelvin
scales are most often
used.
• The Celsius scale is
based on the properties
of water.

108. Temperature
• In scientific
measurements, the
Celsius and Kelvin
scales are most often
used.
• The Celsius scale is
based on the properties
of water.
□ 0°C is the freezing point
of water.

109. Temperature
• In scientific
measurements, the
Celsius and Kelvin
scales are most often
used.
• The Celsius scale is
based on the properties
of water.
□ 0°C is the freezing point
of water.
□ 100°C is the boiling
point of water.

110. Temperature

111. Temperature
• The Kelvin is the SI
unit of temperature.

112. Temperature
• The Kelvin is the SI
unit of temperature.
• It is based on the
properties of gases.

113. Temperature
• The Kelvin is the SI
unit of temperature.
• It is based on the
properties of gases.
• There are no
negative Kelvin
temperatures.

114. Temperature
• The Kelvin is the SI
unit of temperature.
• It is based on the
properties of gases.
• There are no
negative Kelvin
temperatures.
• K = °C + 273.15

115. Temperature

116. Temperature
• The Fahrenheit
scale is not used in
scientific
measurements.

117. Temperature
• The Fahrenheit
scale is not used in
scientific
measurements.
• °F = 9/5(°C) + 32

118. Temperature
• The Fahrenheit
scale is not used in
scientific
measurements.
• °F = 9/5(°C) + 32
• °C = 5/9(°F − 32)

119. Density:
Physical property of a substance

120. Density:
Physical property of a substance
m
d=
V

121. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(a) Calculate the density of mercury
if 1.00 × 10 2 g occupies a volume of
7.36 cm3.

122. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(a) Calculate the density of mercury
if 1.00 × 10 2 g occupies a volume of
7.36 cm3.
Solution
(a) We are given mass and volume, so Equation 1.3
yields

123. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(b) Calculate the volume of 65.0 g of the
liquid methanol (wood alcohol) if its density
is 0.791 g/mL.

124. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(b) Calculate the volume of 65.0 g of the
liquid methanol (wood alcohol) if its density
is 0.791 g/mL.
Solution
(b) Solving Equation 1.3 for volume and then
using the given mass and density gives

125. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(c) What is the mass in grams of a cube of gold
(density = 19.32 g/ cm3) if the length of the cube
is 2.00 cm?

126. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(c) What is the mass in grams of a cube of gold
(density = 19.32 g/ cm3) if the length of the cube
is 2.00 cm?
Solution
(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:

127. SAMPLE EXERCISE 1.4 Determining Density and Using
Density to Determine Volume or Mass
(c) What is the mass in grams of a cube of gold
(density = 19.32 g/ cm3) if the length of the cube
is 2.00 cm?
Solution
(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:
Solving Equation 1.3 for mass and substituting the volume and
density of the cube, we have

128. Uncertainty in Measurements
Different measuring devices have different
uses and different degrees of accuracy.

129. PRACTICE EXERCISE
A balance has a precision of ± 0.001 g. A sample
that has a mass of about 25 g is placed on this
balance. How many significant figures should be
reported for this measurement?

130. PRACTICE EXERCISE
A balance has a precision of ± 0.001 g. A sample
that has a mass of about 25 g is placed on this
balance. How many significant figures should be
reported for this measurement?
Answer: five, as in the measurement 24.995 g

131. Uncertainty in
Measurement

132. Significant Figures

133. Significant Figures
• The term significant figures refers to
digits that were measured.

134. Significant Figures
• The term significant figures refers to
digits that were measured.
• When rounding calculated numbers, we
pay attention to significant figures so we
do not overstate the accuracy of our
answers.

135. Significant Figures

136. Significant Figures
1. All nonzero digits are significant.

137. Significant Figures
1. All nonzero digits are significant.
2. Zeroes between two significant figures
are themselves significant.

138. Significant Figures
1. All nonzero digits are significant.
2. Zeroes between two significant figures
are themselves significant.
3. Zeroes at the beginning of a number
are never significant.

139. Significant Figures
1. All nonzero digits are significant.
2. Zeroes between two significant figures
are themselves significant.
3. Zeroes at the beginning of a number
are never significant.
4. Zeroes at the end of a number are
significant if a decimal point is written
in the number.

140. SAMPLE EXERCISE 1.6 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?

141. SAMPLE EXERCISE 1.6 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 term does not add to the
number of significant figures.

142. SAMPLE EXERCISE 1.6 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 term does not add to the
number of significant figures.
(c) One. We assume that the zeros are not significant
when there is no decimal point shown. If the number
has more significant figures, a decimal point should be
employed or the number written in exponential
notation. Thus, 5000. has four significant figures,
whereas 5.00 × 103 has three.

143. Significant Figures

144. Significant Figures
• When addition or subtraction is
performed, answers are rounded to the
least significant decimal place.

145. Significant Figures
• When addition or subtraction is
performed, answers are rounded to the
least significant decimal place.
• When multiplication or division is
performed, answers are rounded to the
number of digits that corresponds to the
least number of significant figures in
any of the numbers used in the
calculation.

146. SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a
Calculated Quantity
The width, length, and height of a small box are 15.5 cm, 27.3
cm, and 5.4 cm, respectively. Calculate the volume of the box,
using the correct number of significant figures in your answer.

147. SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a
Calculated Quantity
The width, length, and height of a small box are 15.5 cm, 27.3
cm, and 5.4 cm, respectively. Calculate the volume of the box,
using the correct number of significant figures in your answer.
Solution
In multiplication & division, count sig figs.
In addition & subtraction, count decimal places.

148. SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a
Calculated Quantity
The width, length, and height of a small box are 15.5 cm, 27.3
cm, and 5.4 cm, respectively. Calculate the volume of the box,
using the correct number of significant figures in your answer.
Solution
In multiplication & division, count sig figs.
In addition & subtraction, count decimal places.

149. SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a
Calculated Quantity
A gas at 25°C fills a container whose volume is 1.05 × 103
cm3. The container plus gas have a mass of 837.6 g. The
container, when emptied of all gas, has a mass of 836.2 g.
What is the density of the gas at 25°C?

150. SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a
Calculated Quantity
A gas at 25°C fills a container whose volume is 1.05 × 103
cm3. The container plus gas have a mass of 837.6 g. The
container, when emptied of all gas, has a mass of 836.2 g.
What is the density of the gas at 25°C?
Solution
To calculate the density, we must know both the mass and the
volume of the gas.
mass of the gas = full - empty container:
(837.6 – 836.2) g = 1.4 g

151. SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a
Calculated Quantity
A gas at 25°C fills a container whose volume is 1.05 × 103
cm3. The container plus gas have a mass of 837.6 g. The
container, when emptied of all gas, has a mass of 836.2 g.
What is the density of the gas at 25°C?
Solution
To calculate the density, we must know both the mass and the
volume of the gas.
mass of the gas = full - empty container:
(837.6 – 836.2) g = 1.4 g

152. Accuracy versus Precision

153. Accuracy versus Precision
• Accuracy refers to the proximity of
a measurement to the true value
of a quantity.

154. Accuracy versus Precision
• Accuracy refers to the proximity of
a measurement to the true value
of a quantity.
• Precision refers to the proximity of
several measurements to each
other.

155. SAMPLE EXERCISE 1.9 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.)

156. SAMPLE EXERCISE 1.9 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.)

157. SAMPLE EXERCISE 1.9 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 lb to g, we look for a
relationship between these units of mass. From the back inside cover
we have 1 lb = 453.6 g. In order to cancel pounds and leave grams,
we write the conversion factor with grams in the numerator and
pounds in the denominator:

158. SAMPLE EXERCISE 1.9 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 lb to g, we look for a
relationship between these units of mass. From the back inside cover
we have 1 lb = 453.6 g. In order to cancel pounds and leave grams,
we write the conversion factor with grams in the numerator and
pounds in the denominator:

159. SAMPLE EXERCISE 1.9 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 lb to g, we look for a
relationship between these units of mass. From the back inside cover
we have 1 lb = 453.6 g. In order to cancel pounds and leave grams,
we write the conversion factor with grams in the numerator and
pounds in the denominator:
The answer can be given to only three significant
figures, the number of significant figures in 115 lb.

160. SAMPLE EXERCISE 1.10 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.

161. SAMPLE EXERCISE 1.10 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 On the back inside cover of the book, we find that
1 mi = 1.6093 km
1 km = 103 m
60 s = 1 min 60 min = 1 hr

162. SAMPLE EXERCISE 1.10 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 On the back inside cover of the book, we find that
1 mi = 1.6093 km
1 km = 103 m
60 s = 1 min 60 min = 1 hr

163. SAMPLE EXERCISE 1.11 Converting Volume Units
Earth’s oceans contain approximately 1.36 × 109 km3 of
water. Calculate the volume in liters.

164. SAMPLE EXERCISE 1.11 Converting Volume Units
Earth’s oceans contain approximately 1.36 × 109 km3 of
water. Calculate the volume in liters.
Solution
1 L = 10–3 m3
1 km = 103 m

165. SAMPLE EXERCISE 1.11 Converting Volume Units
Earth’s oceans contain approximately 1.36 × 109 km3 of
water. Calculate the volume in liters.
Solution
1 L = 10–3 m3
1 km = 103 m
Thus, converting from km3 to m3 to L, we have