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Chapter 1 Lecture- Matter & Measurement

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Chapter one lecture slides for AP Chemistry
Matter and Measurement

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Chapter 1 Lecture- Matter & Measurement

  1. 1. Chemistry
  2. 2. Chemistry • is the study of properties of materials and changes that they undergo.
  3. 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. 4. The Atomic and Molecular Perspective of Chemistry
  5. 5. The Atomic and Molecular Perspective of Chemistry Chemistry involves the study of the properties and the behavior of matter.
  6. 6. The Atomic and Molecular Perspective of Chemistry Chemistry involves the study of the properties and the behavior of matter. Matter:
  7. 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. 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. 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. 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. 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. 12. Elements
  13. 13. Elements
  14. 14. Elements • are made up of unique atoms, the building blocks of matter.
  15. 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. 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. 17. Molecules
  18. 18. Molecules
  19. 19. Molecules • are combinations of atoms held together in specific shapes.
  20. 20. Molecules • are combinations of atoms held together in specific shapes. • Macroscopic (observable) properties of matter relate to submicroscopic realms of atoms.
  21. 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. 22. 1.2 Classifications of Matter Matter is classified by state (solid, liquid or gas) or by composition (element, compound or mixture).
  23. 23. States of Matter
  24. 24. States of Matter Solids, liquids and gases are the three forms of matter called the states of matter.
  25. 25. States of Matter Solids, liquids and gases are the three forms of matter called the states of matter.
  26. 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. 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. 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. 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. 30. States of Matter
  31. 31. States of Matter Properties described on the molecular level:
  32. 32. States of Matter Properties described on the molecular level: • gas: molecules far apart, move at high speeds, collide often.
  33. 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. 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. 35. Pure Substances
  36. 36. Pure Substances Pure substances:
  37. 37. Pure Substances Pure substances: • are matter with fixed compositions and distinct proportions.
  38. 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. 39. Mixtures
  40. 40. Mixtures • are a combination of two or more pure substances.
  41. 41. Mixtures • are a combination of two or more pure substances. • Each substance retains its own identity.
  42. 42. Elements
  43. 43. Elements • There are 116 known elements.
  44. 44. Elements • There are 116 known elements. • They vary in abundance.
  45. 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. 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. 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. 48. Compounds
  49. 49. Compounds • Compounds are combinations of elements.
  50. 50. Compounds • Compounds are combinations of elements. Example: The compound H2O is a combination of the elements H and O.
  51. 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. 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. 53. Law of Constant (Definite) Proportions
  54. 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).
  55. 55. Mixtures
  56. 56. Mixtures • A mixture is a combination of two or more pure substances.
  57. 57. 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.
  58. 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. • Mixtures have variable composition.
  59. 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. • Heterogeneous mixtures do not have uniform composition, properties, and appearance, e.g., sand.
  60. 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. • Homogeneous mixtures are uniform throughout, e.g., air; they are solutions.
  61. 61. 1.3 Properties of Matter
  62. 62. 1.3 Properties of Matter Each substance has a unique set of physical and chemical properties.
  63. 63. 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.).
  64. 64. 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).
  65. 65. 1.3 Properties of Matter
  66. 66. 1.3 Properties of Matter Properties may be categorized as intensive or extensive.
  67. 67. 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.).
  68. 68. 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.).
  69. 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.). • 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.
  70. 70. Physical and Chemical Changes
  71. 71. Physical and Chemical Changes • Physical change: substance changes physical appearance without altering its identity (e.g., changes of state).
  72. 72. 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).
  73. 73. Separation of Mixtures
  74. 74. Separation of Mixtures Key: separation techniques exploit differences in properties of the components.
  75. 75. Separation of Mixtures Key: separation techniques exploit differences in properties of the components. • Filtration: remove solid from liquid.
  76. 76. 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.
  77. 77. 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.
  78. 78. The Scientific Method
  79. 79. The Scientific Method The scientific method provides guidelines for the practice of science.
  80. 80. The Scientific Method The scientific method provides guidelines for the practice of science. • Collect data (observe, experiment, etc.).
  81. 81. 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.
  82. 82. 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.
  83. 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. • Test hypothesis, then refine it. • Bring all information together into a scientific law (concise statement or equation that summarizes tested hypotheses).
  84. 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. • 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.
  85. 85. 1.3 Steps in the Scientific Method
  86. 86. 1.3 Steps in the Scientific Method
  87. 87. 1.3 Steps in the Scientific Method
  88. 88. 1.3 Steps in the Scientific Method
  89. 89. 1.3 Steps in the Scientific Method
  90. 90. 1.3 Steps in the Scientific Method
  91. 91. Units of Measurement
  92. 92. SI Units
  93. 93. SI Units • Système International d’Unités
  94. 94. SI Units • Système International d’Unités • Uses a different base unit for each quantity
  95. 95. Metric System Prefixes convert the base units into units that are appropriate for the item being measured.
  96. 96. 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.
  97. 97. 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
  98. 98. Volume
  99. 99. Volume • The most commonly used metric units for volume are the liter (L) and the milliliter (mL).
  100. 100. 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.
  101. 101. 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.
  102. 102. Temperature: A measure of the average kinetic energy of the particles in a sample.
  103. 103. Temperature
  104. 104. Temperature • In scientific measurements, the Celsius and Kelvin scales are most often used.
  105. 105. Temperature • In scientific measurements, the Celsius and Kelvin scales are most often used. • The Celsius scale is based on the properties of water.
  106. 106. 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.
  107. 107. 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.
  108. 108. Temperature
  109. 109. Temperature • The Kelvin is the SI unit of temperature.
  110. 110. Temperature • The Kelvin is the SI unit of temperature. • It is based on the properties of gases.
  111. 111. Temperature • The Kelvin is the SI unit of temperature. • It is based on the properties of gases. • There are no negative Kelvin temperatures.
  112. 112. 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
  113. 113. Temperature
  114. 114. Temperature • The Fahrenheit scale is not used in scientific measurements.
  115. 115. Temperature • The Fahrenheit scale is not used in scientific measurements. • °F = 9/5(°C) + 32
  116. 116. Temperature • The Fahrenheit scale is not used in scientific measurements. • °F = 9/5(°C) + 32 • °C = 5/9(°F − 32)
  117. 117. Density: Physical property of a substance
  118. 118. Density: Physical property of a substance m d= V
  119. 119. 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.
  120. 120. 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
  121. 121. 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.
  122. 122. 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
  123. 123. 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?
  124. 124. 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:
  125. 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? 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
  126. 126. Uncertainty in Measurements Different measuring devices have different uses and different degrees of accuracy.
  127. 127. 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?
  128. 128. 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
  129. 129. Uncertainty in Measurement
  130. 130. Significant Figures
  131. 131. Significant Figures • The term significant figures refers to digits that were measured.
  132. 132. 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.
  133. 133. Significant Figures
  134. 134. Significant Figures 1. All nonzero digits are significant.
  135. 135. Significant Figures 1. All nonzero digits are significant. 2. Zeroes between two significant figures are themselves significant.
  136. 136. 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.
  137. 137. 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.
  138. 138. 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?
  139. 139. 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.
  140. 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? 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.
  141. 141. Significant Figures
  142. 142. Significant Figures • When addition or subtraction is performed, answers are rounded to the least significant decimal place.
  143. 143. 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.
  144. 144. 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.
  145. 145. 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.
  146. 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. Solution In multiplication & division, count sig figs. In addition & subtraction, count decimal places.
  147. 147. 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?
  148. 148. 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
  149. 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? 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
  150. 150. Accuracy versus Precision
  151. 151. Accuracy versus Precision • Accuracy refers to the proximity of a measurement to the true value of a quantity.
  152. 152. 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.
  153. 153. 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.)
  154. 154. 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.)
  155. 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.) 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:
  156. 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.) 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:
  157. 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: The answer can be given to only three significant figures, the number of significant figures in 115 lb.
  158. 158. 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.
  159. 159. 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
  160. 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. 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
  161. 161. SAMPLE EXERCISE 1.11 Converting Volume Units Earth’s oceans contain approximately 1.36 × 109 km3 of water. Calculate the volume in liters.
  162. 162. 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
  163. 163. 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

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