Key Terms, Key Equations, Summaries, and Exercises (Chapter 1)

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Chapter One

Key Terms, Key Equations, Summaries, and Exercises (Chapter 1)

Key Terms

accuracy: how closely a measurement aligns with a correct value

atom: smallest particle of an element that can enter into a chemical combination

Celsius (°C): unit of temperature; water freezes at 0 °C and boils at 100 °C on this scale

chemical change: change producing a different kind of matter from the original kind of matter

chemical property: behavior that is related to the change of one kind of matter into another kind of matter

chemistry: study of the composition, properties, and interactions of matter

compound: pure substance that can be decomposed into two or more elements

cubic centimeter (cm³ or cc): volume of a cube with an edge length of exactly 1 cm

cubic meter (m³): SI unit of volume

density: ratio of mass to volume for a substance or object

dimensional analysis: (also, factor-label method) versatile mathematical approach that can be applied to computations ranging from simple unit conversions to more complex, multi-step calculations involving several different quantities

element: substance that is composed of a single type of atom; a substance that cannot be decomposed by a chemical change

exact number: number derived by counting or by definition

extensive property: property of a substance that depends on the amount of the substance

Fahrenheit: unit of temperature; water freezes at 32 °F and boils at 212 °F on this scale

gas: state in which matter has neither definite volume nor shape

heterogeneous mixture: combination of substances with a composition that varies from point to point

homogeneous mixture: (also, solution) combination of substances with a composition that is uniform throughout

hypothesis: tentative explanation of observations that acts as a guide for gathering and checking information

intensive property: property of a substance that is independent of the amount of the substance

kelvin (K): SI unit of temperature; 273.15 K = 0 ºC

kilogram (kg): standard SI unit of mass; 1 kg = approximately 2.2 pounds

law: statement that summarizes a vast number of experimental observations, and describes or predicts some aspect of the natural world

law of conservation of matter: when matter converts from one type to another or changes form, there is no detectable change in the total amount of matter present

length: measure of one dimension of an object

liquid: state of matter that has a definite volume but indefinite shape

liter (L): (also, cubic decimeter) unit of volume; 1 L = 1,000 cm³

macroscopic domain: realm of everyday things that are large enough to sense directly by human sight and touch

mass: fundamental property indicating amount of matter

matter: anything that occupies space and has mass

meter (m): standard metric and SI unit of length; 1 m = approximately 1.094 yards

microscopic domain: realm of things that are much too small to be sensed directly

milliliter (mL): 1/1,000 of a liter; equal to 1 cm³

mixture: matter that can be separated into its components by physical means

molecule: bonded collection of two or more atoms of the same or different elements

physical change: change in the state or properties of matter that does not involve a change in its chemical composition

physical property: characteristic of matter that is not associated with any change in its chemical composition

plasma: gaseous state of matter containing a large number of electrically charged atoms and/or molecules

precision: how closely a measurement matches the same measurement when repeated

pure substance: homogeneous substance that has a constant composition

rounding: procedure used to ensure that calculated results properly reflect the uncertainty in the measurements used in the calculation

scientific method: path of discovery that leads from question and observation to law or hypothesis to theory, combined with experimental verification of the hypothesis and any necessary modification of the theory

second (s): SI unit of time

SI units (International System of Units): standards fixed by international agreement in the International System of Units (Le Système International d’Unités)

significant figures: (also, significant digits) all of the measured digits in a determination, including the uncertain last digit

solid: state of matter that is rigid, has a definite shape, and has a fairly constant volume

symbolic domain: specialized language used to represent components of the macroscopic and microscopic domains, such as chemical symbols, chemical formulas, chemical equations, graphs, drawings, and calculations

temperature: intensive property representing the hotness or coldness of matter

theory: well-substantiated, comprehensive, testable explanation of a particular aspect of nature

uncertainty: estimate of amount by which measurement differs from true value

unit: standard of comparison for measurements

unit conversion factor: ratio of equivalent quantities expressed with different units; used to convert from one unit to a different unit

volume: amount of space occupied by an object

weight: force that gravity exerts on an object

Key Equations

density = \frac{mass}{volume}

T_{°C} = \frac{5}{9} \times (T_{°F} - 32)

T_{°F} = (\frac{9}{5} \times T_{°C}) + 32

T_{K} = °C + 273.15

T_{°C} = K - 273.15

Summaries

1.1 Chemistry in Context

Chemistry deals with the composition, structure, and properties of matter, and the ways by which various forms of matter may be interconverted. Thus, it occupies a central place in the study and practice of science and technology. Chemists use the scientific method to perform experiments, pose hypotheses, and formulate laws and develop theories, so that they can better understand the behavior of the natural world. To do so, they operate in the macroscopic, microscopic, and symbolic domains. Chemists measure, analyze, purify, and synthesize a wide variety of substances that are important to our lives.

1.2 Phases and Classification of Matter

Matter is anything that occupies space and has mass. The basic building block of matter is the atom, the smallest unit of an element that can enter into combinations with atoms of the same or other elements. In many substances, atoms are combined into molecules. On earth, matter commonly exists in three states: solids, of fixed shape and volume; liquids, of variable shape but fixed volume; and gases, of variable shape and volume. Under high-temperature conditions, matter also can exist as a plasma. Most matter is a mixture: It is composed of two or more types of matter that can be present in varying amounts and can be separated by physical means. Heterogeneous mixtures vary in composition from point to point; homogeneous mixtures have the same composition from point to point. Pure substances consist of only one type of matter. A pure substance can be an element, which consists of only one type of atom and cannot be broken down by a chemical change, or a compound, which consists of two or more types of atoms.

1.3 Physical and Chemical Properties

All substances have distinct physical and chemical properties, and may undergo physical or chemical changes. Physical properties, such as hardness and boiling point, and physical changes, such as melting or freezing, do not involve a change in the composition of matter. Chemical properties, such flammability and acidity, and chemical changes, such as rusting, involve production of matter that differs from that present beforehand.

Measurable properties fall into one of two categories. Extensive properties depend on the amount of matter present, for example, the mass of gold. Intensive properties do not depend on the amount of matter present, for example, the density of gold. Heat is an example of an extensive property, and temperature is an example of an intensive property.

1.4 Measurements

Measurements provide quantitative information that is critical in studying and practicing chemistry. Each measurement has an amount, a unit for comparison, and an uncertainty. Measurements can be represented in either decimal or scientific notation. Scientists primarily use SI (International System) units such as meters, seconds, and kilograms, as well as derived units, such as liters (for volume) and g/cm³ (for density). In many cases, it is convenient to use prefixes that yield fractional and multiple units, such as microseconds (10⁻⁶ seconds) and megahertz (10⁶ hertz), respectively.

1.5 Measurement Uncertainty, Accuracy, and Precision

Quantities can be defined or measured. Measured quantities have an associated uncertainty that is represented by the number of significant figures in the quantity’s number. The uncertainty of a calculated quantity depends on the uncertainties in the quantities used in the calculation and is reflected in how the value is rounded. Quantities are characterized with regard to accuracy (closeness to a true or accepted value) and precision (variation among replicate measurement results).

1.6Mathematical Treatment of Measurement Results

Measurements are made using a variety of units. It is often useful or necessary to convert a measured quantity from one unit into another. These conversions are accomplished using unit conversion factors, which are derived by simple applications of a mathematical approach called the factor-label method or dimensional analysis. This strategy is also employed to calculate sought quantities using measured quantities and appropriate mathematical relations.

Exercises

1.1 Chemistry in Context

1.

Explain how you could experimentally determine whether the outside temperature is higher or lower than 0 °C (32 °F) without using a thermometer.

2.

Identify each of the following statements as being most similar to a hypothesis, a law, or a theory. Explain your reasoning.

(a) Falling barometric pressure precedes the onset of bad weather.

(b) All life on earth has evolved from a common, primitive organism through the process of natural selection.

(c) My truck’s gas mileage has dropped significantly, probably because it’s due for a tune-up.

3.

Identify each of the following statements as being most similar to a hypothesis, a law, or a theory. Explain your reasoning.

(a) The pressure of a sample of gas is directly proportional to the temperature of the gas.

(b) Matter consists of tiny particles that can combine in specific ratios to form substances with specific properties.

(c) At a higher temperature, solids (such as salt or sugar) will dissolve better in water.

4.

Identify each of the underlined items as a part of either the macroscopic domain, the microscopic domain, or the symbolic domain of chemistry. For any in the symbolic domain, indicate whether they are symbols for a macroscopic or a microscopic feature.

(a) The mass of a lead pipe is 14 lb.

(b) The mass of a certain chlorine atom is 35 amu.

(c) A bottle with a label that reads Al contains aluminum metal.

(d) Al is the symbol for an aluminum atom.

5.

Identify each of the underlined items as a part of either the macroscopic domain, the microscopic domain, or the symbolic domain of chemistry. For those in the symbolic domain, indicate whether they are symbols for a macroscopic or a microscopic feature.

(a) A certain molecule contains one H atom and one Cl atom.

(b) Copper wire has a density of about 8 g/cm³.

(c) The bottle contains 15 grams of Ni powder.

(d) A sulfur molecule is composed of eight sulfur atoms.

6.

According to one theory, the pressure of a gas increases as its volume decreases because the molecules in the gas have to move a shorter distance to hit the walls of the container. Does this theory follow a macroscopic or microscopic description of chemical behavior? Explain your answer.

7.

The amount of heat required to melt 2 lbs of ice is twice the amount of heat required to melt 1 lb of ice. Is this observation a macroscopic or microscopic description of chemical behavior? Explain your answer.

1.2 Phases and Classification of Matter

8.

Why is an object’s mass, rather than its weight, used to indicate the amount of matter it contains?

9.

What properties distinguish solids from liquids? Liquids from gases? Solids from gases?

10.

How does a heterogeneous mixture differ from a homogeneous mixture? How are they similar?

11.

How does a homogeneous mixture differ from a pure substance? How are they similar?

12.

How does an element differ from a compound? How are they similar?

13.

How do molecules of elements and molecules of compounds differ? In what ways are they similar?

14.

How does an atom differ from a molecule? In what ways are they similar?

15.

Many of the items you purchase are mixtures of pure compounds. Select three of these commercial products and prepare a list of the ingredients that are pure compounds.

16.

Classify each of the following as an element, a compound, or a mixture:

(a) copper

(b) water

(c) nitrogen

(d) sulfur

(e) air

(f) sucrose

(g) a substance composed of molecules each of which contains two iodine atoms

(h) gasoline

17.

Classify each of the following as an element, a compound, or a mixture:

(a) iron

(b) oxygen

(c) mercury oxide

(d) pancake syrup

(e) carbon dioxide

(f) a substance composed of molecules each of which contains one hydrogen atom and one chlorine atom

(g) baking soda

(h) baking powder

18.

A sulfur atom and a sulfur molecule are not identical. What is the difference?

19.

How are the molecules in oxygen gas, the molecules in hydrogen gas, and water molecules similar? How do they differ?

20.

Why are astronauts in space said to be “weightless,” but not “massless”?

21.

Prepare a list of the principal chemicals consumed and produced during the operation of an automobile.

22.

Matter is everywhere around us. Make a list by name of fifteen different kinds of matter that you encounter every day. Your list should include (and label at least one example of each) the following: a solid, a liquid, a gas, an element, a compound, a homogenous mixture, a heterogeneous mixture, and a pure substance.

23.

When elemental iron corrodes it combines with oxygen in the air to ultimately form red brown iron(III) oxide called rust. (a) If a shiny iron nail with an initial mass of 23.2 g is weighed after being coated in a layer of rust, would you expect the mass to have increased, decreased, or remained the same? Explain. (b) If the mass of the iron nail increases to 24.1 g, what mass of oxygen combined with the iron?

24.

As stated in the text, convincing examples that demonstrate the law of conservation of matter outside of the laboratory are few and far between. Indicate whether the mass would increase, decrease, or stay the same for the following scenarios where chemical reactions take place:

(a) Exactly one pound of bread dough is placed in a baking tin. The dough is cooked in an oven at 350 °F releasing a wonderful aroma of freshly baked bread during the cooking process. Is the mass of the baked loaf less than, greater than, or the same as the one pound of original dough? Explain.

(b) When magnesium burns in air a white flaky ash of magnesium oxide is produced. Is the mass of magnesium oxide less than, greater than, or the same as the original piece of magnesium? Explain.

(c) Antoine Lavoisier, the French scientist credited with first stating the law of conservation of matter, heated a mixture of tin and air in a sealed flask to produce tin oxide. Did the mass of the sealed flask and contents decrease, increase, or remain the same after the heating?

25.

Yeast converts glucose to ethanol and carbon dioxide during anaerobic fermentation as depicted in the simple chemical equation here:

(a) If 200.0 g of glucose is fully converted, what will be the total mass of ethanol and carbon dioxide produced?

(b) If the fermentation is carried out in an open container, would you expect the mass of the container and contents after fermentation to be less than, greater than, or the same as the mass of the container and contents before fermentation? Explain.

(c) If 97.7 g of carbon dioxide is produced, what mass of ethanol is produced?

1.3 Physical and Chemical Properties

26.

Classify the six underlined properties in the following paragraph as chemical or physical:

Fluorine is a pale yellow gas that reacts with most substances. The free element melts at −220 °C and boils at −188 °C. Finely divided metals burn in fluorine with a bright flame. Nineteen grams of fluorine will react with 1.0 gram of hydrogen.

27.

Classify each of the following changes as physical or chemical:

(a) condensation of steam

(b) burning of gasoline

(c) souring of milk

(d) dissolving of sugar in water

(e) melting of gold

28.

Classify each of the following changes as physical or chemical:

(a) coal burning

(b) ice melting

(c) mixing chocolate syrup with milk

(d) explosion of a firecracker

(e) magnetizing of a screwdriver

29.

The volume of a sample of oxygen gas changed from 10 mL to 11 mL as the temperature changed. Is this a chemical or physical change?

30.

A 2.0-liter volume of hydrogen gas combined with 1.0 liter of oxygen gas to produce 2.0 liters of water vapor. Does oxygen undergo a chemical or physical change?

31.

Explain the difference between extensive properties and intensive properties.

32.

Identify the following properties as either extensive or intensive.

(a) volume

(b) temperature

(c) humidity

(d) heat

(e) boiling point

33.

The density (d) of a substance is an intensive property that is defined as the ratio of its mass (m) to its volume (V).

Considering that mass and volume are both extensive properties, explain why their ratio, density, is intensive.

1.4 Measurements

34.

Is one liter about an ounce, a pint, a quart, or a gallon?

35.

Is a meter about an inch, a foot, a yard, or a mile?

36.

Indicate the SI base units or derived units that are appropriate for the following measurements:

(a) the length of a marathon race (26 miles 385 yards)

(b) the mass of an automobile

(c) the volume of a swimming pool

(d) the speed of an airplane

(e) the density of gold

(f) the area of a football field

(g) the maximum temperature at the South Pole on April 1, 1913

37.

Indicate the SI base units or derived units that are appropriate for the following measurements:

(a) the mass of the moon

(b) the distance from Dallas to Oklahoma City

(c) the speed of sound

(d) the density of air

(e) the temperature at which alcohol boils

(f) the area of the state of Delaware

(g) the volume of a flu shot or a measles vaccination

38.

Give the name and symbol of the prefixes used with SI units to indicate multiplication by the following exact quantities.

(a) 10³

(b) 10⁻²

(c) 0.1

(d) 10⁻³

(e) 1,000,000

(f) 0.000001

39.

Give the name of the prefix and the quantity indicated by the following symbols that are used with SI base units.

(a) c

(b) d

(c) G

(d) k

(e) m

(f) n

(g) p

(h) T

40.

A large piece of jewelry has a mass of 132.6 g. A graduated cylinder initially contains 48.6 mL water. When the jewelry is submerged in the graduated cylinder, the total volume increases to 61.2 mL.

(a) Determine the density of this piece of jewelry.

(b) Assuming that the jewelry is made from only one substance, what substance is it likely to be? Explain.

41.

Visit this PhET density simulation and select the Same Volume Blocks.

(a) What are the mass, volume, and density of the yellow block?

(b) What are the mass, volume and density of the red block?

(c) List the block colors in order from smallest to largest mass.

(d) List the block colors in order from lowest to highest density.

(e) How are mass and density related for blocks of the same volume?

42.

Visit this PhET density simulation and select Custom Blocks and then My Block.

(a) Enter mass and volume values for the block such that the mass in kg is less than the volume in L. What does the block do? Why? Is this always the case when mass < volume?

(b) Enter mass and volume values for the block such that the mass in kg is more than the volume in L. What does the block do? Why? Is this always the case when mass > volume?

(c) How would (a) and (b) be different if the liquid in the tank were ethanol instead of water?

(d) How would (a) and (b) be different if the liquid in the tank were mercury instead of water?

43.

Visit this PhET density simulation and select Mystery Blocks.

(a) Pick one of the Mystery Blocks and determine its mass, volume, density, and its likely identity.

(b) Pick a different Mystery Block and determine its mass, volume, density, and its likely identity.

(c) Order the Mystery Blocks from least dense to most dense. Explain.

1.5 Measurement Uncertainty, Accuracy, and Precision

44.

Express each of the following numbers in scientific notation with correct significant figures:

(a) 711.0

(b) 0.239

(c) 90743

(d) 134.2

(e) 0.05499

(f) 10000.0

(g) 0.000000738592

45.

Express each of the following numbers in exponential notation with correct significant figures:

(a) 704

(b) 0.03344

(c) 547.9

(d) 22086

(e) 1000.00

(f) 0.0000000651

(g) 0.007157

46.

Indicate whether each of the following can be determined exactly or must be measured with some degree of uncertainty:

(a) the number of eggs in a basket

(b) the mass of a dozen eggs

(c) the number of gallons of gasoline necessary to fill an automobile gas tank

(d) the number of cm in 2 m

(e) the mass of a textbook

(f) the time required to drive from San Francisco to Kansas City at an average speed of 53 mi/h

47.

Indicate whether each of the following can be determined exactly or must be measured with some degree of uncertainty:

(a) the number of seconds in an hour

(b) the number of pages in this book

(c) the number of grams in your weight

(d) the number of grams in 3 kilograms

(e) the volume of water you drink in one day

(f) the distance from San Francisco to Kansas City

48.

How many significant figures are contained in each of the following measurements?

(a) 38.7 g

(b) 2 $\times$

10¹⁸ m

(c) 3,486,002 kg

(d) 9.74150 $\times$

10⁻⁴ J

(e) 0.0613 cm³

(f) 17.0 kg

(g) 0.01400 g/mL

49.

How many significant figures are contained in each of the following measurements?

(a) 53 cm

(b) 2.05 $\times$

10⁸ m

(c) 86,002 J

(d) 9.740 $\times$

10⁴ m/s

(e) 10.0613 m³

(f) 0.17 g/mL

(g) 0.88400 s

50.

The following quantities were reported on the labels of commercial products. Determine the number of significant figures in each.

(a) 0.0055 g active ingredients

(b) 12 tablets

(c) 3% hydrogen peroxide

(d) 5.5 ounces

(e) 473 mL

(f) 1.75% bismuth

(g) 0.001% phosphoric acid

(h) 99.80% inert ingredients

51.

Round off each of the following numbers to two significant figures:

(a) 0.436

(b) 9.000

(c) 27.2

(d) 135

(e) 1.497 $\times$

10⁻³

(f) 0.445

52.

Round off each of the following numbers to two significant figures:

(a) 517

(b) 86.3

(c) 6.382 $\times$

10³

(d) 5.0008

(e) 22.497

(f) 0.885

53.

Perform the following calculations and report each answer with the correct number of significant figures.

(a) 628 $\times$

342

(b) (5.63 $\times$

10²) $\times$

(7.4 $\times$

10³)

(c) $\frac{28.0}{13.483}$

(d) 8119 $\times$

0.000023

(e) 14.98 + 27,340 + 84.7593

(f) 42.7 + 0.259

54.

Perform the following calculations and report each answer with the correct number of significant figures.

(a) 62.8 $\times$

34

(b) 0.147 + 0.0066 + 0.012

(c) 38 $\times$

95 $\times$

1.792

(d) 15 – 0.15 – 0.6155

(e) $8.78 \times (\frac{0.0500}{0.478})$

(f) 140 + 7.68 + 0.014

(g) 28.7 – 0.0483

(h) $\frac{(88.5 - 87.57)}{45.13}$

55.

Consider the results of the archery contest shown in this figure.

(a) Which archer is most precise?

(b) Which archer is most accurate?

(c) Who is both least precise and least accurate?

56.

Classify the following sets of measurements as accurate, precise, both, or neither.

(a) Checking for consistency in the weight of chocolate chip cookies: 17.27 g, 13.05 g, 19.46 g, 16.92 g

(b) Testing the volume of a batch of 25-mL pipettes: 27.02 mL, 26.99 mL, 26.97 mL, 27.01 mL

(c) Determining the purity of gold: 99.9999%, 99.9998%, 99.9998%, 99.9999%

1.6 Mathematical Treatment of Measurement Results

57.

Write conversion factors (as ratios) for the number of:

(a) yards in 1 meter

(b) liters in 1 liquid quart

(c) pounds in 1 kilogram

58.

Write conversion factors (as ratios) for the number of:

(a) kilometers in 1 mile

(b) liters in 1 cubic foot

(c) grams in 1 ounce

59.

The label on a soft drink bottle gives the volume in two units: 2.0 L and 67.6 fl oz. Use this information to derive a conversion factor between the English and metric units. How many significant figures can you justify in your conversion factor?

60.

The label on a box of cereal gives the mass of cereal in two units: 978 grams and 34.5 oz. Use this information to find a conversion factor between the English and metric units. How many significant figures can you justify in your conversion factor?

61.

Soccer is played with a round ball having a circumference between 27 and 28 in. and a weight between 14 and 16 oz. What are these specifications in units of centimeters and grams?

62.

A woman’s basketball has a circumference between 28.5 and 29.0 inches and a maximum weight of 20 ounces (two significant figures). What are these specifications in units of centimeters and grams?

63.

How many milliliters of a soft drink are contained in a 12.0-oz can?

64.

A barrel of oil is exactly 42 gal. How many liters of oil are in a barrel?

65.

The diameter of a red blood cell is about 3 $\times$

10⁻⁴ in. What is its diameter in centimeters?

66.

The distance between the centers of the two oxygen atoms in an oxygen molecule is 1.21 $\times$

10⁻⁸ cm. What is this distance in inches?

67.

Is a 197-lb weight lifter light enough to compete in a class limited to those weighing 90 kg or less?

68.

A very good 197-lb weight lifter lifted 192 kg in a move called the clean and jerk. What was the mass of the weight lifted in pounds?

69.

Many medical laboratory tests are run using 5.0 μL blood serum. What is this volume in milliliters?

70.

If an aspirin tablet contains 325 mg aspirin, how many grams of aspirin does it contain?

71.

Use scientific (exponential) notation to express the following quantities in terms of the SI base units in Table 1.3:

(a) 0.13 g

(b) 232 Gg

(c) 5.23 pm

(d) 86.3 mg

(e) 37.6 cm

(f) 54 μm

(g) 1 Ts

(h) 27 ps

(i) 0.15 mK

72.

Complete the following conversions between SI units.

(a) 612 g = ________ mg

(b) 8.160 m = ________ cm

(c) 3779 μg = ________ g

(d) 781 mL = ________ L

(e) 4.18 kg = ________ g

(f) 27.8 m = ________ km

(g) 0.13 mL = ________ L

(h) 1738 km = ________ m

(i) 1.9 Gg = ________ g

73.

Gasoline is sold by the liter in many countries. How many liters are required to fill a 12.0-gal gas tank?

74.

Milk is sold by the liter in many countries. What is the volume of exactly 1/2 gal of milk in liters?

75.

A long ton is defined as exactly 2240 lb. What is this mass in kilograms?

76.

Make the conversion indicated in each of the following:

(a) the men’s world record long jump, 29 ft 4¼ in., to meters

(b) the greatest depth of the ocean, about 6.5 mi, to kilometers

(c) the area of the state of Oregon, 96,981 mi², to square kilometers

(d) the volume of 1 gill (exactly 4 oz) to milliliters

(e) the estimated volume of the oceans, 330,000,000 mi³, to cubic kilometers.

(f) the mass of a 3525-lb car to kilograms

(g) the mass of a 2.3-oz egg to grams

77.

Make the conversion indicated in each of the following:

(a) the length of a soccer field, 120 m (three significant figures), to feet

(b) the height of Mt. Kilimanjaro, at 19,565 ft, the highest mountain in Africa, to kilometers

(c) the area of an 8.5- × 11-inch sheet of paper in cm²

(d) the displacement volume of an automobile engine, 161 in.³, to liters

(e) the estimated mass of the atmosphere, 5.6 × 10¹⁵ tons, to kilograms

(f) the mass of a bushel of rye, 32.0 lb, to kilograms

(g) the mass of a 5.00-grain aspirin tablet to milligrams (1 grain = 0.00229 oz)

78.

Many chemistry conferences have held a 50-Trillion Angstrom Run (two significant figures). How long is this run in kilometers and in miles? (1 Å = 1 $\times$

10⁻¹⁰ m)

79.

A chemist’s 50-Trillion Angstrom Run (see Exercise 1.78) would be an archeologist’s 10,900 cubit run. How long is one cubit in meters and in feet? (1 Å = 1 $\times$

10⁻⁸ cm)

80.

The gas tank of a certain luxury automobile holds 22.3 gallons according to the owner’s manual. If the density of gasoline is 0.8206 g/mL, determine the mass in kilograms and pounds of the fuel in a full tank.

81.

As an instructor is preparing for an experiment, he requires 225 g phosphoric acid. The only container readily available is a 150-mL Erlenmeyer flask. Is it large enough to contain the acid, whose density is 1.83 g/mL?

82.

To prepare for a laboratory period, a student lab assistant needs 125 g of a compound. A bottle containing 1/4 lb is available. Did the student have enough of the compound?

83.

A chemistry student is 159 cm tall and weighs 45.8 kg. What is her height in inches and weight in pounds?

84.

In a recent Grand Prix, the winner completed the race with an average speed of 229.8 km/h. What was his speed in miles per hour, meters per second, and feet per second?

85.

Solve these problems about lumber dimensions.

(a) To describe to a European how houses are constructed in the US, the dimensions of “two-by-four” lumber must be converted into metric units. The thickness $\times$

width $\times$

length dimensions are 1.50 in. $\times$

3.50 in. $\times$

8.00 ft in the US. What are the dimensions in cm $\times$

cm $\times$

m?

(b) This lumber can be used as vertical studs, which are typically placed 16.0 in. apart. What is that distance in centimeters?

86.

The mercury content of a stream was believed to be above the minimum considered safe—1 part per billion (ppb) by weight. An analysis indicated that the concentration was 0.68 parts per billion. What quantity of mercury in grams was present in 15.0 L of the water, the density of which is 0.998 g/ml? $(1 ppb Hg = \frac{1 ng Hg}{1 g water})$

87.

Calculate the density of aluminum if 27.6 cm³ has a mass of 74.6 g.

88.

Osmium is one of the densest elements known. What is its density if 2.72 g has a volume of 0.121 cm³?

89.

Calculate these masses.

(a) What is the mass of 6.00 cm³ of mercury, density = 13.5939 g/cm³?

(b) What is the mass of 25.0 mL octane, density = 0.702 g/cm³?

90.

Calculate these masses.

(a) What is the mass of 4.00 cm³ of sodium, density = 0.97 g/cm³ ?

(b) What is the mass of 125 mL gaseous chlorine, density = 3.16 g/L?

91.

Calculate these volumes.

(a) What is the volume of 25 g iodine, density = 4.93 g/cm³?

(b) What is the volume of 3.28 g gaseous hydrogen, density = 0.089 g/L?

92.

Calculate these volumes.

(a) What is the volume of 11.3 g graphite, density = 2.25 g/cm³?

(b) What is the volume of 39.657 g bromine, density = 2.928 g/cm³?

93.

Convert the boiling temperature of gold, 2966 °C, into degrees Fahrenheit and kelvin.

94.

Convert the temperature of scalding water, 54 °C, into degrees Fahrenheit and kelvin.

95.

Convert the temperature of the coldest area in a freezer, −10 °F, to degrees Celsius and kelvin.

96.

Convert the temperature of dry ice, −77 °C, into degrees Fahrenheit and kelvin.

97.

Convert the boiling temperature of liquid ammonia, −28.1 °F, into degrees Celsius and kelvin.

98.

The label on a pressurized can of spray disinfectant warns against heating the can above 130 °F. What are the corresponding temperatures on the Celsius and kelvin temperature scales?

99.

The weather in Europe was unusually warm during the summer of 1995. The TV news reported temperatures as high as 45 °C. What was the temperature on the Fahrenheit scale?

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Key Terms

Key Equations

Summaries

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

1.6Mathematical Treatment of Measurement Results

Exercises

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 Measurement Results

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