Kelter P., Mosher M., Scott A. Chemistry. The Practical Science

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78 Chapter 2 Atoms: A Quest for Understanding

chemical formula A representation, in symbols, that con-

veys the relative proportions of atoms of the different

elements in a substance. (p. 68)

chemical nomenclature A system of rules used to assign

a name to a particular substance. (p. 72)

coulomb The unit of electrical charge. (p. 55)

covalent bonds Bonds that atoms form by sharing elec-

trons. (p. 70)

Dalton’s atomic theory The theory, developed by John

Dalton, that all substances are composed of indivisi-

ble atoms. (p. 50)

diatomic elements Elements whose normal state is in the

form of molecules composed of two atoms attached

together (most notably H

,Cl

,Br

and I

). (p. 71)

electron One of the subatomic particles of which atoms

are composed. Electrons carry a charge of −1.

(p. 52)

empirical formula A chemical formula indicating the

ratio in which elements are found to be present in a

particular compound, regardless of the molecular

structure of the compound. (p. 71)

gamma ray High-energy electromagnetic radiation

emitted from the decay of a radioactive element.

(p. 54)

group One of the vertical columns in the periodic table

of the elements. (p. 64)

halogens The elements of Group VIIA of the periodic

table. (p. 66)

inert gases See noble gases. (p. 65)

ion An electrically charged entity that results when an

atom has gained or lost electrons. (p. 56)

ionic compound A compound composed of ions. (p. 68)

isotopes Forms of the same element that differ in the

number of neutrons within the nucleus. (p. 57)

law of combining volumes Scientiﬁc law stating that when

gases combine, they do so in small whole-number ra-

tios, provided that all the gases are at the same tem-

perature and pressure. (p. 50)

law of conservation of mass Scientiﬁc law stating that the

mass of chemicals present at the start of a chemical

reaction must equal the mass of chemicals present at

the end of the reaction. Although not strictly true,

this law is correct at the level of accuracy of all labo-

ratory balances. (p. 47)

law of deﬁnite composition Scientiﬁc law stating that any

particular chemical is always composed of its compo-

nents in a ﬁxed ratio, by mass. (p. 48)

law of multiple proportions Scientiﬁc law stating that

when the same elements can produce more than one

compound, the ratio of the masses of the element

that combine with a ﬁxed mass of another element

corresponds to a small whole number. (p. 49)

mass number (A) The total number of protons and neu-

trons in the nucleus of an atom. (p. 57)

mass spectrometer An instrument used to measure the

masses and abundances of isotopes, molecules, and

fragments of molecules. (p. 59)

mass spectrum The output of a mass spectrometer, in

the form of a chart listing the abundance and masses

of the isotopes, molecules, and fragments of mole-

cules present. (p. 60)

metalloids A small number of elements, found at the

boundary between metals and nonmetals in the peri-

odic table, that have properties intermediate between

those of metals and nonmetals. Also known as semi-

metals. (p. 65)

metals The largest category of elements in the periodic

table. Metals exhibit such characteristic properties as

ability to conduct electricity, shiny appearance, and

malleability. (p. 64)

molecular formula A chemical formula indicating the ac-

tual number of each type of atom present in one

molecule of a compound. (p. 70)

molecular ions Ions composed of two or more atoms at-

tached together. Also known as polyatomic ions.

(p. 75)

molecules Neutral compounds containing two or more

atoms attached together. (p. 70)

neutron One of the subatomic particles of which atoms

are composed, carrying no electrical charge and

found in the nucleus. (p. 54)

noble gases The unreactive elements in the rightmost

group of the periodic table (Group VIIIA). Also

known as inert gases. (p. 65)

nonmetals The elements on the right-hand side of the

periodic table. Nonmetals exhibit characteristic prop-

erties that are distinct from those of the metals.

(p. 65)

nuclide notation Shorthand notation used to represent

atoms, listing the symbol for the element, accompa-

nied by the atomic number and the mass number of

the atom. (p. 57)

period One of the horizontal rows in the periodic table

of the elements. (p. 64)

periodic table of the elements A table presenting all the

elements, in the form of horizontal periods and verti-

cal groups; shown on the inside front cover of this

book. (p. 56)

polyatomic ions Ions composed of two or more atoms

covalently attached together. Also known as molecu-

lar ions. (p. 75)

products The materials that are produced in a chemical

reaction. (p. 47)

proton One of the subatomic particles of which atoms

are composed, carrying an electrical charge of +1

and found in the nucleus. (p. 53)

radiation The particles and/or energy emitted during

radioactive decay. (p. 54)

radioactive decay The process by which an unstable nu-

cleus becomes more stable via the emission or ab-

sorption of particles and accompanying energy.

(p. 54)

radioactivity The emission of radioactive particles

and/or energy. (p. 54)

Focus Your Learning 79

Focus Your Learning

The answers to the odd-numbered problems and some selected

problems appear at the back of the book, as represented by the

blue numbering.

Section 2.1 Early Attempts to Explain Matter

Skill Review

1. Exploring the unseen world often relies on indirect reason-

ing. Useful deductions led early philosophers and scientists

to make conclusions about the possible existence of atoms.

How would the formation of ice crystals on the branches of a

bush support the ideas of Democritus?

2. Ancient practices, and some modern ones, involve burning

incense. How would entering a room in which incense had

recently been burned lead you to consider the world to be

made of small, invisible particles?

3. The scientists who ﬁrst used atoms to explain chemical

changes assumed that the small, unseen pieces of matter

could continually be rearranged into new combinations with

new properties. Show how that same principle operates when

the letters in the word dormitory are used to form several

other words.

4. Taking the analogy in Problem 3 a bit further, explain what

would have to happen in order for the exercise to illustrate

the law of conservation of mass.

5. In your own words, explain the law of conservation of mass.

6. In your own words, explain the law of deﬁnite composition.

7. Modern science is based on asking questions. In what way

does a researcher determining the age of Earth ask the same

basic questions posed by Democritus?

8. What connection exists between the exploration and obser-

vations of the United States space program and Democritus’

questions?

Chemical Applications and Practices

9. Some compounds absorb water from the atmosphere in such

a way that the water is chemically combined with the com-

pound. The absorbed water is called water of hydration, and

the compounds with absorbed water are known as hydrates.

If 14.7 g of calcium chloride hydrate, when heated, loses 3.6 g

of water, how many grams of water are lost when a 23.4-g

sample is heated? To which law(s) did you refer in obtaining

these results?

10. The white appearance of several magnesium-containing

compounds causes an apparent resemblance among the

compounds. One such compound was shown to contain

60.0% magnesium. Another sample of a magnesium-

containing compound was found. What quantity of magne-

sium, in grams, would have to be present in a 34.6-g sample

of the second compound in order for us to claim that it is the

same compound as the ﬁrst sample?

Section 2.2 Dalton’s Atomic Theory and Beyond

Skill Review

11. Lavoisier used carefully determined masses to make discover-

ies about changes that occurred in chemical reactions.

a. Which fundamental law did his work lead him to discover?

b. Suppose Lavoisier evaporated the water from 250.0 g of

seawater. If the remaining salt had a mass of 2.2 g, what

would you calculate as the mass of the water that had

evaporated?

12. A 200.0-g sample of a pure substance was composed of

132.9 g of copper and 67.1 g of sulfur. Another sample of the

same substance, this time with a mass of 150.0 g, was brought

to Joseph Proust to analyze. How many grams of copper and

sulfur would you expect to ﬁnd in the sample? What basic

law in chemistry are you employing in order to make your

determination?

13. Fill in the missing information for a compound of copper

and oxygen, assuming that it is composed of a ﬁxed ratio of

copper to oxygen of 3.97:1.00.

14. Fill in the missing information for a compound consisting of

sodium and oxygen, assuming that is composed of a ﬁxed

ratio of sodium to oxygen of 1.44:1.00.

Chemical Applications and Practices

15. Copper and oxygen combine in more than one ratio. In one

such compound it is found that 65 grams of copper combine

with 16 grams of oxygen. If you assumed, as John Dalton

might have, that the combination represented one atom of

copper combining with one atom of oxygen, what would you

predict as the copper-to-oxygen ratio in another compound

between the two?

Sodium Oxygen

6.0 grams grams

grams 50.0 grams

Copper Oxygen

17.0 grams grams

grams 11.5 grams

reactants The materials that combine in a chemical re-

action. (p. 47)

semimetals A small number of elements found at the

boundary between metals and nonmetals in the peri-

odic table. Also known as metalloids. (p. 65)

80 Chapter 2 Atoms: A Quest for Understanding

16. Using the ﬁrst mass ratio presented in Problem 15 and your

answer for the other ratio, determine how many grams of

oxygen would combine with 10.0 g of copper in both

compounds.

17. When the calcium is properly extracted, you can expect to

obtain 40 g of calcium from 100.0 g of calcium carbonate. A

10.0-g sample of pure calcium chloride yields 6.40 g of chlo-

rine. Which of these two compounds is a better source of

calcium?

18. Three magnesium compounds have the following percent-

ages of magnesium: Compound A: 28.6% magnesium;

Compound B: 60.0% magnesium; Compound C: 41.4%

magnesium. Why, according to Dalton, can there be such a

great difference in the percentages of magnesium?

19. Water and hydrogen peroxide are both clear liquids that are

made up of only hydrogen and oxygen. Water contains 2 g of

hydrogen for every 16 g of oxygen in water. In hydrogen

peroxide there are 2 g of hydrogen for every 32 g of oxygen.

Show how these data illustrate the law of multiple

proportions.

20. Oxygen exists naturally in two combined forms. The di-

atomic oxygen that we utilize in breathing consists of two

atoms combined into one unit.Ozone consists of three atoms

of oxygen joined into one unit. The relative masses of the oxy-

gen and ozone units are 32 and 48. Does this agree with the

law of multiple proportions? Explain.

Section 2.3 The Structure of the Atom

Skill Review

21. If the discoverers of the particles that make up the atom had

determined the charge on the electron to be −2, what would

be the charge on the neutron and proton?

22. Describe what might have been the results of Rutherford’s

experiment if the charge on the electron had been reversed

(that is, if it were +1 instead of −1).

23. Draw diagrams of the models of the atom that J. J. Thomson

and Rutherford, respectively, would have come up with if the

electron had a +1 charge.

24. What are two objections to the explanation that positive pro-

tons compose a nucleus that is surrounded by negative

electrons.

+1 electron

– charged nucleus

Section 2.4 Atoms and Isotopes

Skill Review

25. Fill in the information missing from the following table:

26. Fill in the information missing from the following table:

27. Indicate whether each of these phrases describes the proton,

neutron, or electron. (You may need to assign more than one

term to each phrase.)

a. Determines the identity of an atom.

b. Has about the same mass as a proton.

c. The number of particles can be changed without changing

the identity of the atom.

28. Indicate whether each of these phrases refers to protons, neu-

trons, and/or electrons. (You may need to assign more than

one term to each phrase.)

a. Is represented by the mass number.

b. Is represented by the atomic number.

c. Are always present in the same quantity in an atom with

0 charge.

29. What would you calculate for the mass, in grams, of an atom

that contained 1 proton and 1 electron? . . . 1 proton, 1 neu-

tron, and 1 electron?

30. What would you calculate for the mass, in grams, of an atom

that contained 2 protons, 2 neutrons, and 2 electrons? . . .

1 proton, 1 neutron, and 2 electrons?

31. What would be the charge on each of the particles listed in

Problem 29?

32. What would be the charge on each of the particles listed in

Problem 30?

33. Assume that an atom contained 1 proton and 1 electron.

By what percentage would the mass of this atom increase if

10 electrons were added to the atom?

34. Assume that an atom contained 1 proton and 1 electron.

By what percentage would the mass of this atom increase if

10 neutrons were added to the atom?

35. Which of the most abundant isotope in each pair contains

the greater number of neutrons? Assume the mass number of

the most abundant isotope can be found by rounding the av-

erage atomic mass to the nearest whole number. (Use the pe-

riodic table of the elements inside the front cover of your

textbook.)

a. iron or chromium c. cobalt or nickel

b. tellurium or iodine d. helium or neon

Isotope Protons Neutrons Electrons Charge

calcium-40 20 18

14 14 0

sulfur-32 16 14

Isotope Protons Neutrons Electrons Charge

carbon-12 6 6 0

13 14 10

chlorine-35 18 −1

Focus Your Learning 81

36. Which of the most abundant isotope in each pair contains

the greater number of neutrons? Round the average atomic

mass to the nearest whole number to obtain the mass num-

ber of the most abundant isotope. (Use the periodic table of

the elements inside the front cover of your textbook.)

a. zinc or sodium c. phosphorus or sulfur

b. hydrogen or sodium d. bromine or selenium

37. Arrange the following atoms in order of their quantity of

protons (least number to greatest number): iron, hydrogen,

calcium, ﬂuorine, aluminum, boron.

38. Arrange these atoms in order of their quantity of protons

(least number to greatest number): sodium, magnesium,

copper, iodine, tin, carbon, lithium.

39. Arrange these atoms in order of their quantity of electrons

(least number to greatest number): iron, hydrogen, calcium,

ﬂuorine, aluminum, boron.

40. Arrange these atoms in order of their quantity of electrons

(least number to greatest number): sodium, magnesium,

copper, iodine, tin, carbon, lithium.

41. Using the periodic table, determine the ﬁrst element that

does not have the same number of protons as neutrons.

Round the average atomic mass to the nearest whole number

to obtain the mass number of the most abundant isotope.

42. Using the periodic table, indicate two atoms that have the

same number of protons as neutrons. Round the average

atomic mass to the nearest whole number to obtain the mass

number of the most abundant isotope.

43. In a neutral atom, the number of electrons must equal the

number of protons. However, if an atom had two more pro-

tons than electrons, what would you report as its charge?

Would this ion be called an anion or a cation?

44. If an atom had two more electrons than protons, what would

you report as its charge? Would this ion be called an anion or

a cation?

Chemical Applications and Practices

45. Several radioactive isotopes have medical applications. Using

the nuclide representation shown here, determine the num-

ber of protons, electrons, and neutrons in an atom of each of

these elements.

Radioactive sodium is used in blood circulation studies:

Radioactive iodine can be used to examine the thyroid gland:

131

Radioactive cobalt can be used to kill cancer cells:

Radioactive chromium can be used to measure total blood

volume:

Radioactive phosphorus is used in antileukemia therapy:

46. Write nuclide notation for each of the following isotopes:

carbon-12, carbon-13, phosphorus-32, chlorine-35,

chlorine-37, iron-55.

Section 2.5 Atomic Mass

Skill Review

47. If the atomic mass of

C had been set equal to 6.000000 amu,

what would be the mass of

O? . . . the mass of

48. If the atomic mass of oxygen-16 were set equal to 4.000000

amu, what would be the mass of carbon-12? . . . the mass of

copper-63?

49. Calculate the mass, in grams, of 1 atom of carbon-12. . . . of

nitrogen-14. . . . of ﬂuorine-19. (Assume that the mass num-

ber of these isotopes is equal to the atomic mass, rounded to

the nearest whole number.)

50. Calculate the mass, in grams, of a gross (144) of atoms of

carbon-12. . . . of nitrogen-14. . . . of ﬂuorine-19. (Assume

that the mass number of these isotopes is equal to the atomic

mass, rounded to the nearest whole number.)

Chemical Applications and Practices

51. The metal indium is used to make heat-conducting alloys. In-

dium typically is found with two isotopes. Based on the fol-

lowing information, what would you calculate as the percent

relative abundance of the two isotopes? Indium-113 has a

mass of 112.9043, and indium-115 has a mass of 114.9041.

52. The element neon is widely used in brightly colored elec-

tronic advertising signs. There are three common isotopes of

neon with masses of 21.99, 20.99, and 19.99. One of the three

isotopes makes up approximately 90% of the atoms in a neon

sample. Which isotope is most likely to be responsible for the

majority of neon’s atomic mass reported on the periodic

table? Sketch a diagram of the probable appearance of the

mass spectrum for neon that shows the masses and an ap-

proximation of the percentage abundance.

53. The sulfur-containing amino acids that help us make certain

proteins are critically important to our diet. Sulfur samples

typically consist of four isotopes. Use the following table to

determine the average atomic mass of sulfur.

Sketch a mass spectrum that illustrates the masses observed

for a natural sample of sulfur.

Isotope Atomic Mass Abundance

S-32 31.97 95.0%

S-33 32.97 0.76%

S-34 33.97 4.22%

S-36 35.97 0.014%

Neon gas discharge tube.

82 Chapter 2 Atoms: A Quest for Understanding

54. Silver has many uses in our world, only one of which is as a

precious metal in coinage. Use the following table to deter-

mine the average atomic mass of silver.

Sketch a mass spectrum that illustrates the masses observed

for a natural sample of silver.

55. Cerium has four naturally occurring isotopes. Use the fol-

lowing table to determine the average atomic mass of cerium.

Sketch a mass spectrum that illustrates the masses observed

for a natural sample of cerium.

56. A hypothetical element was just discovered to have three

isotopes. Use the following table to determine the average

atomic mass of this hypothetical element.

Sketch a mass spectrum that illustrates the masses observed

for a natural sample of the element.

Section 2.6 The Periodic Table

Skill Review

57. Use the descriptions given in this chapter of the organization

of the periodic table to identify each of these elements.

a. The element that is in the second column and in the

second period.

b. The transition element that is in the seventh group and in

the fourth row.

c. The third most massive noble gas.

58. Use the descriptions given in this chapter of the organization

of the periodic table to identify each of these elements.

a. The lightest alkali metal.

b. The semimetal that is in the third group.

c. The halogen that is a liquid at room temperature.

59. How many elements are classiﬁed as halogens?

60. Which column contains the alkali metals? . . . the chalcogens?

61. Identify the name of the group for each of these elements.

sulfur iodine helium

beryllium francium

Isotope Atomic Mass Abundance

A-100 99.754 35.25%

A-102 101.688 25.75%

A-103 102.599 39.00%

Isotope Atomic Mass Abundance

Ce-136 135.907 0.19%

Ce-138 137.906 0.25%

Ce-140 139.905 88.48%

Ce-142 141.909 11.08%

Isotope Atomic Mass Abundance

Ag-107 106.9051 51.84%

Ag-109 108.9048 48.16%

62. Identify the name of the group for each of these elements.

lithium barium neon

oxygen chlorine

63. Which of these elements would you predict to be shiny?

gold silver lead

silicon carbon iodine

64. Which of these elements would you predict to be brittle?

gold silver lead

silicon carbon iodine

Section 2.7 Ionic Compounds

Skill Review

65. One of the ﬁrst considerations in writing ionic formulas is

to determine which element is to be the cation and which

is to be the anion. Decide which element in each of these

pairs is more likely to become the positive ion.

Ca or Br S or Al Cl or Al

Sr or N I or Be

66. Use the information from the previous question to deter-

mine what formula is most likely to result from the combina-

tion of each of these pairs of elements.

Ca and Br S and Al Cl and Al

Sr and N I and Be

67. Write the formulas of the ionic compounds that form when

chlorine combines individually with lithium, beryllium,

sodium, calcium, and aluminum.

68. Write the formulas of the ionic compounds that form

when oxygen combines individually with lithium, beryllium,

sodium, calcium, and aluminum.

69. Ionic compounds typically dissociate into ions as they dis-

solve in water. For each of these compounds, predict the ions,

including their charges, that will be produced when the ionic

compound dissolves.

MgBr

FeCl

KI Na

70. Ionic compounds typically dissociate into ions as they dis-

solve in water. For each of these compounds, predict the ions,

including their charges, that will be produced when the ionic

compound dissolves.

CaCl

TiO

ZnO NiCl

Section 2.8 Molecules

Skill Review

71. The two major classiﬁcations of compounds are ionic and

molecular. Label the descriptions given below as characteris-

tic of an ionic compounds or of a molecule

a. Results from a metal combining with a nonmetal.

b. Persists as a complete entity when dissolved in water.

c. Formula can often be determined by the identity of the

atoms that combine.

Focus Your Learning 83

72. The two major classiﬁcations of compounds are ionic and

molecular. Label the descriptions given below as characteris-

tic of an ionic compounds or of a molecule.

a. Results as atoms share their electrons.

b. Formula represents a ratio of ions involved rather than

actual numbers of atoms combined.

c. Results from the combination of two nonmetals.

d. Separates into individual charged particles when dissolved

in water.

73. Which of these compounds are not molecular compounds?

a. CO d. PCl

b. BrO

e. WO

c. K

74. Which of these compounds are molecular compounds?

a. SO

d. Na

b. NiO e. SiF

c. BCl

75. In these examples, the molecular and empirical formulas of

several compounds are listed. In each case, decide which is

the empirical formula.

a. Hydrogen peroxide is a useful disinfectant. (H

or HO)

b. Butane is used as a heating fuel. (C

or C

)

c. Teﬂon is made from tetraﬂuoroethylene. (C

or CF

)

d. Glucose is a sweet-tasting molecule. (CH

O or C

)

76. In each of these examples, determine the correct empirical

formula.

a. Acetic acid is found in vinegar: C

b. Cane sugar is called sucrose: C

c. Paradichlorobenzene is used as a moth repellent: C

d. Benzene is a component of gasoline: C

Section 2.9 Naming Compounds

Skill Review

77. Provide the correct name for each of these binary molecular

compounds.

PCl

CCl

78. Provide the correct name for each of these binary molecular

compounds.

PBr

79. Provide the correct name for each of these binary ionic

compounds.

O CaBr

AlCl

BaS

80. Provide the correct name for each of these binary ionic

compounds.

CaO MgCl

LiF

81. Provide the correct name for each of these ionic compounds.

CaBr

Fe(NO

)

CaSO

Cl NaCl

82. Provide the correct name for each of these ionic compounds.

KHCO

Ca(CN)

Co(OH)

Cu(NO

)

MgSO

83. Provide the correct formula for each of these compounds.

copper(II) hydroxide

chromium(III) oxide

sulfur hexachloride

carbon tetraiodide

aluminum hydroxide

magnesium sulfate

sodium sulﬁte

ammonium hydroxide

boron tribromide

sodium acetate

84. Provide the correct formula for each of these compounds.

hydrogen chloride

cobalt(VI) ﬂuoride

dinitrogen tetroxide

calcium perchlorate

barium nitrate

lead(IV) chloride

potassium dichromate

sodium bicarbonate

lithium hydroxide

titanium(IV) carbonate

85. Using the metal ion Mn

, write the correct formula for a

combination with each of the following: sulfate; chloride;

nitrite; carbonate; and bisulﬁte.

86. Using the metal ion Cr

, write the correct formula for a

combination with each of the following: sulfate; chloride;

nitrite; carbonate; and bisulﬁte.

87. Write the correct formula for oxalate compounds of the

following metals: Mn

;Cu

;Fe

;Mn

; and Ti

88. Write the correct formula for nitrate compounds of the

following metals: Mn

;Cu

;Fe

;Mn

; and Ti

89. Name these compounds.

(NH

)

NaHCO

Cu(HSO

)

Ca(OH)

KMnO

Mg(CN)

LiClO

90. Name these compounds.

Zn(CN)

CrO

HPO

Ba(NO

)

SrSO

KClO

91. Determine the charge on the metal cation in these

compounds.

V(NO

)

TiSO

W(C

)

AgOH Ru(HCO

)

92. Determine the charge on the metal cation in these

compounds.

Cr(NO

)

MnO

Pd(C

)

AuCl

Co(OH)

Chemical Applications and Practices

93. A lab accident caused part of the label on a bottle of a chem-

ical to be obscured. All that is visible is “Fe(NO

.” Is this

enough information to name the compound? What possible

suggestions could you make to determine the identity of the

compound?

84 Chapter 2 Atoms: A Quest for Understanding

97. When hydrogen gas combines with chlorine gas, a corrosive

gas known as hydrogen chloride forms. According to

Dalton, the volume ratio between the reactants and prod-

ucts should be 1:1:1. However, Gay-Lussac would claim a

ratio of 1:1:2 (assuming that all gases are at the same tem-

perature and pressure). Gay-Lussac’s claim was proved to be

valid. Which law is being used here? Explain why this expla-

nation works better than Dalton’s.

98. If the mass of a metal nut were 35.0 g and the mass of a

matching threaded bolt were 7.0 g, a combined nut and bolt

would have a mass of 42.0 g.

a. What would be the mass of the threaded bolts needed to

attach to 175.0 g of nuts?

b. If someone decided to attach two bolts to a nut, what

would be the resulting mass of the unit?

c. What would be the mass of the threaded bolts needed to

attach to 175.0 g of nuts under the conditions stated in

part b?

d. What is the mass ratio between bolts connected to nuts in

the ﬁrst unit (part a) and bolts connected to nuts in the

second unit (part b)?

99. Robert Millikan determined the electrical charge for an

electron, and its mass was easily calculated by using the

charge-to-mass ratio. The following analogy applies

roughly the same logic: Suppose you were buying a box of

golf balls. What information would you need in order to

know the price of each golf ball?

100. A tiny oil droplet is given an electrical charge by an electric

current. The charge causes the oil droplet to be attracted to

an oppositely charged metal plate above it. What two forces

must be equal in order for an oil droplet to be suspended

(neither rising nor falling) in the apparatus?

94. The same lab accident caused part of the label on another

bottle to be obscured as well. The visible part reads “NH

C.”

Which of the following would you suggest as the most likely

name of the compound in this bottle? Explain why you reject

the others: ammonium carbonate, ammonium oxalate,

ammonium chloride, ammonium copper(II).

Comprehensive Problems

95. What were the major objections to Dalton’s atomic theory?

Rewrite Dalton’s atomic theory to take into account our cur-

rent understanding of matter.

96. Match each chemical statement on the left-hand side of the

following list with a correlating statement from Dalton’s

atomic theory on the right.

Chemistry Atomic Theory

(1) Elements are made up of

tiny particles called atoms.

(2) Atoms of a given element

are identical; atoms of

different elements differ in

some fundamental way.

(3) Compounds are formed

when atoms combine; each

compound always has the

same numbers and types of

atoms.

(4) Chemical reactions consist

of the reorganization of

the atoms involved; the

atoms themselves are not

changed.

a. Graphite and

diamond have

different properties,

but both are made

of only atoms of

carbon.

b. 2H

+ O

→ 2H

c. It takes 6.02 ×10

atoms of iron to make

55.85 g.

d. It takes 6.02 × 10

atoms of gold to

make 197 g.

Charged

plate (+)

Charged oil drop

Charged

plate (–)

Window

Viewing chamber

Microscope

Oil

Problem 100

then measured, and the weights were found to be as follows

(assume these are exact weights):

32 oz 28 oz

18 oz 22 oz

36 oz 14 oz

a. What is the mass of each bag in grams?

b. Judging on the basis of the masses recorded in ounces, and

the simplifying assumption that the tomatoes are the same

size, what would you estimate as the weight of just one

tomato?

c. Another bag had an exact weight of 20 oz; how does that

inﬂuence your answer to part a? What is the chemical law

that you are using to answer this question?

d. Describe the process that you could follow to identify the

volume of each of the tomatoes.

107. In this chapter, we discussed determining the age of objects

using radioactive isotopes. Scientists are constantly reﬁning

these techniques to correct for the differences between con-

ditions long ago and those at the present time. What condi-

tions might have changed that could require correcting? In

each case, what would be the nature of the correction (that

is, whether the real age is greater or less) and why?

108. Scientists now think that much of the universe is composed

of “dark matter.” Investigate some scientiﬁc journals, such as

Scientiﬁc American, to learn enough about dark matter to

describe it. On the basis of your research, discuss the exper-

iments that compelled scientists to draw this conclusion.

Why is dark matter useful to understand?

2.2

2.0

1.8

1.6

1.4

1.2

1.0

0 10,000 20,000 30,000 40,000

Age (years before present)

Atmospheric C-14/C-12 ratio

relative to present day

101. How much mass was lost from the combination of protons,

neutrons, and electrons when

Cu (62.9296011 amu) was

formed? What happened to this mass?

102. If the atoms in a 2.500 g sample of aluminum were lined up

end to end, how long would the line of iron be in inches?

(Assume the aluminum atoms are spheres with a radius of

143 pm.)

103. If a sample of carbon atoms contains 7.45  10

atoms,

how many gross of atoms are there in the sample (1 gross 

144 atoms)? If each of those atoms is 154 pm in diameter,

would they stretch to 1 mile if they were lined up so that

they were touching end to end?

104. During the analysis of an unknown sample of metal, a

researcher places a 16.07 g chunk of the metal into a gradu-

ated cylinder containing 20.3 mL of water. The volume of

the mixture was then recorded to be 22.3 mL. What is the

density of the metal?

105. A new isotope of beryllium is created such that it contains

9 neutrons.

a. How many protons and electrons would an atom of this

isotope contain?

b. What is the mass number for this new isotope?

c. Use the information in Table 2.3 to determine the mass of

the isotope in kilograms. (Assume that no mass is lost in

the creation of the new isotope from the individual parts.)

Thinking Beyond the Calculation

106. In the produce section of a grocery store, bags of small

tomatoes were being packaged. The weight of each bag was

Problem 107

Focus Your Learning

Introducing

Quantitative

Chemistry

The use of quantitative chemical

equipment is important in the

biomedical ﬁeld.

Contents and Selected Applications

3.1 Formula Masses

Chemical Encounters: Adrenaline in the Body

3.2 Counting by Weighing

3.3 Working with Moles

Chemical Encounters: Cholesterol in the Blood

3.4 Percentages by Mass

Chemical Encounters: Nitrogen in Fertilizers

3.5 Finding the Formula

Chemical Encounters: A Focus on Chocolate

3.6 Chemical Equations

Chemical Encounters: Preparation of Aspirin

3.7 Working with Equations

Chemical Encounters: Cleaning the Air on Manned Spacecraft

Go to college.hmco.com/pic/kelterMEE for online learning resources.

A sudden, loud noise, the shock of some

dramatic twist in a horror ﬁlm, losing control of

your car on an icy road—any one of these can cause the

rapid physical responses we attribute to the release of

adrenaline, including increasing heart rate, stronger heartbeat,

sweating, and shifting of some blood ﬂow from the skin to the

internal organs. Adrenaline, also known as epinephrine (Figure 3.1),

is normally present in your blood at approximately

2 × 10

–8

g/L and is secreted from the adrenal glands

near your kidneys. In times of extreme stress, your

adrenaline level increases a thousandfold. The powerful

effects of adrenaline demonstrate the vital signiﬁcance of

changing quantities of chemicals. The chemistry of the

body provides many other examples. For instance, taking

600 mg of aspirin can effectively relieve pain. Taking

100 times that amount, 60 g, would kill you. Quantities in

chemistry are crucial.

Up to this point, we have been dealing with individual

atoms and molecules. However, in the lab we typically

measure quantities in liters, grams, and other “macro-sized”

units. What is the relationship between “nanoworld” and

“macroworld” quantities? Speciﬁc questions arise: How

many molecules of adrenaline will be in each liter of blood?

How can we actually measure the level of adrenaline or any other chemical in

blood? These are the kinds of quantitative issues we explore in this chapter.

3.1 Formula Masses

In addition to being naturally present in the body, adrenaline is administered as a drug to

stimulate the heart, to alleviate allergic reactions, and even to help break up fat cells during

liposuction. As you might expect, control over the amounts administered is vital.

To make, use, or detect speciﬁc quantities of adrenaline, we need information about the

mass of its molecules, so we may know how many molecules are in any given mass that is

administered to a patient.

In Chapter 2 we looked at the method for determining the masses of atoms, and we saw

that each element has a given mass (expressed in atomic mass units), which is the average

mass of one atom of the element. We can work out the average

molecular mass (or molecular

weight) of a molecule, such as adrenaline, by adding together the average atomic masses of all

the atoms present in the molecule. To simplify our calculations, we will use atomic masses

rounded to two decimals, except for hydrogen, which we will round to three decimals.

Formula of each adrenaline molecule = C

In the formula we have

9 carbon atoms @ 12.01 amu = 108.09

13 hydrogen atoms @ 1.008 amu = 13.104

1 nitrogen atom @ 14.01 amu = 14.01

3 oxygen atoms @ 16.00 amu = 48.00

Average molecular mass of adrenaline = 183.204 = 183.2 amu

FIGURE 3.1

Adrenaline (epinephrine) has the chemical formula

. In this “ball-and-stick” representation,

carbon atoms are charcoal gray, hydrogen atoms are

light blue, oxygen atoms are red, and nitrogen atoms

are dark blue.

Application

HEMICAL ENCOUNTERS:

Adrenaline in the Body

Aspirin