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CHAPTER 7

Reactions in Aqueous Solution

Water is the most universal of all liquids. Water has a relatively large heat capacity and a relativel
y large liquid range, which means it can absorb the heat liberated by many reactions while still re
maining in the liquid state. Water is very polar and dissolves well both ionic solutes and solutes w
ith which it can hydrogen bond (this is especially important to the biochemical reactions of the liv
ing cell).

Driving forces are types of changes in a system that pull a reaction in the direction of product for
mation; driving forces discussed in Chapter 7 include: formation of a solid, formation of water, fo
rmation of a gas, and transfer of electrons.

precipitation

A reactant in aqueous solution is indicated with (aq). Formation of a solid is indicated with (s)

When an ionic solute such as NaCl (sodium chloride) is dissolved in water, the resulting solution
consists of separate, individual, discrete hydrated sodium ions (Na+) and separate, individual, disc
rete hydrated chloride ions (Cl–). There are no identifiable NaCl units in such a solution and the p
ositive and negative ions behave independently of one another.

Because each formula unit of MgCl2 contains two chloride ions for each magnesium ion, that rati
o will be preserved in the solution when MgCl2 is dissolved in water.

A substance is said to be a strong electrolyte if each unit of the substance produces separated, dist
inct ions when the substance is dissolved in water. NaCl and KNO3 are both strong electrolytes.

Chemists know that a solution contains independent ions because such a solution will readily allo
w an electrical current to pass through it. The simplest experiment that demonstrates this uses the
sort of light–bulb conductivity apparatus described in the text: if the light bulb glows strongly, the
n the solution must contain a lot of ions to be conducting the electricity well.

The solubility rules are general rules describing the solubility of common ionic substances in wat
er. They are based on countless observations of chemical compounds and reactions. For example,
Solubility Rule 1 says that “most nitrate salts are soluble”. So when we write an equation for the r
eaction of a nitrate salt with some other reagent, we know that any precipitate that forms will not i
nvolve the nitrate ion.

10.

(a); The precipitate BaSO4 will form.

11.

insoluble (Rule 6: most sulfide salts are insoluble.)

insoluble (Rule 5: most hydroxide compounds are insoluble)

soluble (Rule 2: most salts of Na+ are soluble; Rule 4: most sulfate salts are soluble.)

soluble (Rule 2: most salts of NH4+ are soluble.)

Chapter 7: Reactions in Aqueous Solution

insoluble (Rule 6: most carbonate salts are insoluble.)

insoluble (Rule 6: most phosphate salts are insoluble.)

insoluble (Exception to Rule 3)

insoluble (Exception to Rule 4)

12.

soluble (Rule 1: most nitrate salts are soluble.)

soluble (Rule 2: most salts of K+ are soluble.)

insoluble (Rule 4: most sulfate salts are soluble with PbSO4 as an exception.)

insoluble (Rule 5: most hydroxide compounds are insoluble.)

soluble (Rule 2: most salts of K+ are soluble.)

insoluble (Rule 3: most chloride salts are soluble with Hg2Cl2 as an exception.)

soluble (Rule 2: most salts of NH4+ are soluble.)

insoluble (Rule 6: most sulfide salts are insoluble.)

13.

Rule 2: Most salts of K+ are soluble.

Rule 1: Most nitrate salts are soluble.

Rule 2: Most salts of NH4+ are soluble.

Rule 4: Most sulfate salts are soluble.

Rule 3: Most chloride salts are soluble.

14.

Rule 5: Most hydroxide compounds are insoluble.

Rule 6: Most carbonate salts are insoluble.

Rule 6: Most phosphate salts are insoluble.

Rule 3: Exception to the rule for chloride salts.

Rule 4: Exception to the rule for sulfate salts.

15.

CuS: Rule 6 (most sulfide salts are insoluble).

Ba3(PO4)2: Rule 6 (most phosphate salts are insoluble).

AgCl: Rule 3 (exception to rule for chloride salts).

CoCO3: Rule 6 (most carbonate salts are insoluble).

CaSO4: Rule 4 (exception to rule for sulfate salts).

Hg2Cl2: Rule 3 (exception to rule for chloride salts).

16.

MnCO3: Rule 6 (most carbonates are only slightly soluble).

CaSO4: Rule 4 (exception for sulfates).

Hg2Cl2: Rule 3: (exception for chlorides).

soluble

Ni(OH)2: Rule 5 (most hydroxides are only slightly soluble).

Chapter 7: Reactions in Aqueous Solution

BaSO4: Rule 4 (exception for sulfates).

17.

The precipitates are marked in boldface type.

No precipitate: both (NH4)2SO4 and HCl are soluble.

NH4Cl(aq) + H2SO4(aq)  no precipitate

Rule 6: Most carbonate salts are only slightly soluble.

2K2CO3(aq) + SnCl4(aq)  Sn(CO3)2(s) + 4KCl(aq)

Rule 3: exception to rule for chlorides

2NH4Cl(aq) + Pb(NO3)2(aq)  PbCl2(s) + 2NH4NO3(aq)

Rule 5: Most hydroxide compounds are only slightly soluble.

CuSO4(aq) + 2KOH(aq)  Cu(OH)2(s) + K2SO4(aq)

Rule 6: Most phosphate salts are only slightly soluble.

Na3PO4(aq) + CrCl3(aq)  CrPO4(s) + 3NaCl(aq)

Rule 6: Most sulfide salts are only slightly soluble.

3(NH4)2S(aq) + 2FeCl3(aq)  Fe2S3(s) + 6NH4Cl(aq)

18.

The formulas of the precipitates are in boldface type.

Rule 6: Most carbonate salts are insoluble.

Na2CO3(aq) + CuSO4(aq)  Na2SO4(aq) + CuCO3(s)

Rule 3: Exception for chloride salts.

HCl(aq) + AgC2H3O2(aq)  HC2H3O2(aq) + AgCl(s)

No precipitate

Rule 6: Most sulfide salts are insoluble.

3(NH4)2S(aq) + 2FeCl3(aq)  6NH4Cl(aq) + Fe2S3(s)

Rule 4: Exception for sulfate salts

H2SO4(aq) + Pb(NO3)2(aq)  2HNO3(aq) + PbSO4(s)

Rule 6: Most phosphate salts are insoluble.

2K3PO4(aq) + 3CaCl2(aq)  6KCl(aq) + Ca3(PO4)2(s)

19.

Hint: when balancing equations involving polyatomic ions, especially in precipitation reactions, b
alance the polyatomic ions as a unit, not in terms of the atoms the polyatomic ions contain (e.g., tr
eat nitrate ion, NO3– as a single entity, not as one nitrogen and three oxygen atoms). When finishe
d balancing, however, be sure to count the individual number of atoms of each type on each side
of the equation.

Na2SO4(aq) + CaCl2(aq)  CaSO4(s) + NaCl(aq)

Balance sodium: Na2SO4(aq) + CaCl2(aq)  CaSO4(s) + 2NaCl(aq)

Balanced equation: Na2SO4(aq) + CaCl2(aq)  CaSO4(s) + 2NaCl(aq)

Co(C2H3O2)2(aq) + Na2S(aq)  CoS(s) + NaC2H3O2(aq)

Chapter 7: Reactions in Aqueous Solution

Balance acetate: Co(C2H3O2)2(aq) + Na2S(aq)  CoS(s) + 2NaC2H3O2(aq)

Balanced equation: Co(C2H3O2)2(aq) + Na2S(aq)  CoS(s) + 2NaC2H3O2(aq)

KOH(aq) + NiCl2(aq)  Ni(OH)2(s) + KCl(aq)

Balance hydroxide: 2KOH(aq) + NiCl2(aq)  Ni(OH)2(s) + KCl(aq)

Balance potassium: 2KOH(aq) + NiCl2(aq)  Ni(OH)2(s) + 2KCl(aq)

Balanced equation: 2KOH(aq) + NiCl2(aq)  Ni(OH)2(s) + 2KCl(aq)

20.

CaCl2 (aq) + AgNO3 (aq)  Ca(NO3)2 (aq) + AgCl (s)

balance chlorine: CaCl2 (aq) + AgNO3 (aq)  Ca(NO3)2 (aq) + 2AgCl (s)

balance silver: CaCl2 (aq) + 2AgNO3 (aq)  Ca(NO3)2 (aq) + 2AgCl (s)

balanced equation: CaCl2 (aq) + 2AgNO3 (aq)  Ca(NO3)2 (aq) + 2AgCl (s)

AgNO3(aq) + K2CrO4(aq)  Ag2CrO4(s) + KNO3(aq)

balance silver: 2AgNO3(aq) + K2CrO4(aq)  Ag2CrO4(s) + KNO3(aq)

balance nitrate ion: 2AgNO3(aq) + K2CrO4(aq)  Ag2CrO4(s) + 2KNO3(aq)

balanced equation: 2AgNO3(aq) + K2CrO4(aq)  Ag2CrO4(s) + 2KNO3(aq)

BaCl2(aq) + K2SO4(aq)  BaSO4(s) + KCl(aq)

balance potassium: BaCl2(aq) + K2SO4(aq)  BaSO4(s) + 2KCl(aq)

balanced equation: BaCl2(aq) + K2SO4(aq)  BaSO4(s) + 2KCl(aq)

21.

The products are determined by having the ions “switch partners.” For example, for a general rea
ction AB + CD  , the possible products are AD and CB if the ions switch partners. If either AD
or CB is insoluble, then a precipitation reaction has occurred. In the following reaction, the formu
la of the precipitate is given in boldface type.

(NH4)2SO4(aq) + Ba(NO3)2(aq)  2NH4NO3(aq) + BaSO4(s)

Rule 4: BaSO4 is a listed exception.

H2S(aq) + NiSO4(aq)  H2SO4(aq) + NiS(s)

Rule 6: Most sulfide salts are only slightly soluble.

FeCl3(aq) + 3NaOH(aq)  3NaCl(aq) + Fe(OH)3(s)

Rule 5: Most hydroxide compounds are only slightly soluble.

22.

The precipitate is lead(II) phosphate. The balanced equation is:

23.

The net ionic equation for a reaction in solution indicates only those components that are directly
involved in the reaction. Other ions that may be present to balance charge, but which do not activ

Chapter 7: Reactions in Aqueous Solution

ely participate in the reaction are called spectator ions and are not indicated when writing the che
mical equation for the reaction.

24.

(e)

25.

The products are determined by having the ions in the two aqueous ionic reagents “switch
partners.” For example, for a general reaction AB + CD  , the possible products are AD and CB
if the ions switch partners. If either AD or CB is insoluble according to the solubility rules in
Table 7.1, then a precipitation reaction has occurred. Answers will vary for each student.

26.

Molecular:

Complete Ionic:

Net Ionic:

27.

Cu2+(aq) + CrO42–(aq)  CuCrO4(s)

Co3+(aq) + CrO42–(aq)  Co2(CrO4)3(s)

Ba2+(aq) + CrO42–(aq)  BaCrO4(s)

Fe3+(aq) + CrO42–(aq)  Fe2(CrO4)3(s)

28.

Ag+(aq) + Cl–(aq)  AgCl(s)

Pb2+(aq) + 2Cl–(aq)  PbCl2(s)

Hg22+(aq) + 2Cl–(aq)  Hg2Cl2(s)

29.

Ca2+(aq) + C2O42–(aq)  CaC2O4(s)

30.

Co2+(aq) + S2–(aq)  CoS(s)

2Co3+(aq) + 3S2–(aq)  Co2S3(s)

Fe2+(aq) + S2–(aq)  FeS(s)

2Fe3+(aq) + 3S2–(aq)  Fe2S3(s)

31.

Strong acids ionize completely in water. The strong acids are also strong electrolytes. Strong elec
trolytes dissociate completely in water.

32.

Strong bases fully produce hydroxide ions when dissolved in water. The strong bases are also str
ong electrolytes. Strong electrolytes dissociate completely in water.

33.

H+(aq) + OH–(aq)  H2O; formation of a water molecule

Chapter 7: Reactions in Aqueous Solution

34.

acids: HCl, H2SO4, HNO3, HClO4, HBr

bases: NaOH, KOH, RbOH, CsOH

35.

1000; 1000

36.

A salt is the ionic product remaining in solution when an acid neutralizes a base. For example, in
the reaction HCl (aq) + NaOH(aq)  NaCl(aq) + H2O(l) sodium chloride is the salt produced by
the neutralization reaction.

37.

Your textbook mentions four strong acids. You only had to give three of the following equations.

HCl(aq)  H+(aq) + Cl–(aq)

HNO3(aq)  H+(aq) + NO3–(aq)

H2SO4(aq)  H+(aq) + HSO4–(aq)

HClO4(aq)  H+(aq) + ClO4–(aq)

38.

HBr(aq)  H+(aq) + Br–(aq)

HClO4(aq)  H+(aq) + ClO4–(aq)

39.

The formulas of the salts are marked in boldface type. Remember that in an acid/base reaction in
aqueous solution, water is always one of the products: keeping this in mind makes predicting the
formula of the salt produced easy to do.

HCl(aq) + KOH(aq)  H2O(l) + KCl(aq)

RbOH(aq) + HNO3(aq)  H2O(l) + RbNO3(aq)

HClO4(aq) + NaOH(aq)  H2O(l) + NaClO4(aq)

HBr(aq) + CsOH(aq)  H2O(l) + CsBr(aq)

40.

In general, the salt formed in an aqueous acid–base reaction consists of the positive ion of the bas
e involved in the reaction, combined with the negative ion of the acid. The hydrogen ion of the str
ong acid combines with the hydroxide ion of the strong base to produce water, which is the other
product of the acid–base reactions.

H2SO4(aq) + 2KOH(aq)  K2SO4(aq) + 2H2O(l)

HNO3(aq) + NaOH(aq)  NaNO3(aq) + H2O(l)

2HCl(aq) + Ca(OH)2(aq)  CaCl2(aq) + 2H2O(l)

2HClO4(aq) + Ba(OH)2(aq)  Ba(ClO4)2(aq) + 2H2O(l)

41.

An oxidation–reduction reaction is one in which one species loses electrons (oxidation) and anoth
er species gains electrons (reduction). Electrons are transferred from the species being oxidized to
the species being reduced.

42.

Answer depends on student choice of example: Na(s) + Cl2(g)  2NaCl(s) is an example.

43.

A driving force, in general, is an event that tends to help to convert the reactants of a process into
the products. Some elements (metals) tend to lose electrons, whereas other elements (nonmetals) t
end to gain electrons. A transfer of electrons from atoms of a metal to atoms of a nonmetal would

Chapter 7: Reactions in Aqueous Solution

be favorable and would result in a chemical reaction. A simple example of such a process is the re
action of sodium with chlorine: sodium atoms tend to each lose one electron (to form Na+), where
as chlorine atoms tend to each gain one electron (to form Cl–). The reaction of sodium metal with
chlorine gas represents a transfer of electrons from sodium atoms to chlorine atoms to form sodiu
m chloride.

44.

The aluminum atoms lose 3 electrons to become Al3+ ions. Fe3+ ions gain 3 electrons to become F
e atoms.

45.

Each calcium atom would lose two electrons. Each fluorine atom would gain one electron (so the
F2 molecule would gain two electrons). One calcium atom would be required to react with one flu
orine, F2, molecule. Calcium ions are charged 2+, fluoride ions are charged 1–.

46.

Each magnesium atom would lose two electrons. Each oxygen atom would gain two electrons (so
the O2 molecule would gain four electrons). Two magnesium atoms would be required to react wi
th each oxygen, O2, molecule. Magnesium ions are charged 2+, oxide ions are charged 2–.

47.

MgCl2 is made up of Mg2+ ions and Cl– ions. Magnesium atoms each lose two electrons to becom
e Mg2+ ions. Chlorine atoms each gain one electron to become Cl– ions (so each Cl2 molecule gai
ns two electrons to become two Cl– ions).

48.

AlBr3 is made up of Al3+ ions and Br– ions. Aluminum atoms each lose three electrons and bromin
e atoms each gain one electron (Br2 gains two electrons).

49.

Co(s) + Br2(l)  CoBr3(s)

Balance bromine: Co(s) + 3Br2(l)  2CoBr3(s)

Balance cobalt: 2Co(s) + 3Br2(l)  2CoBr3(s)

Balanced equation: 2Co(s) + 3Br2(l)  2CoBr3(s)

cobalt is oxidized, bromine is reduced

Al(s) + H2SO4(aq)  Al2(SO4)3(aq) + H2(g)

Balance sulfate ions: Al(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + H2(g)

Balance hydrogen: Al(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2(g)

Balance aluminum: 2Al(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2(g)

Balanced equation: 2Al(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 3H2(g)

aluminum is oxidized, hydrogen is reduced

Na(s) + H2O(l)  NaOH(aq) + H2(g)

Balance hydrogen: Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

Balance sodium: 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

Balanced equation: 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

sodium is oxidized, hydrogen is reduced

Cu(s) + O2(g)  Cu2O(s)

Balance copper: 2Cu(s) + O2(g)  Cu2O(s)

Chapter 7: Reactions in Aqueous Solution

Balance oxygen: 2Cu(s) +

2 O2(g)  Cu2O(s)

Balanced equation: 4Cu(s) + O2(g)  2Cu2O(s)

copper is oxidized, oxygen is reduced

50.

P4(s) + O2(g)  P4O10(s)

balance oxygen: P4(s) + 5O2(g)  P4O10(s)

balanced equation: P4(s) + 5O2(g)  P4O10(s)

MgO(s) + C(s)  Mg(s) + CO(g)

This equation is already balanced.

Sr(s) + H2O(l)  Sr(OH)2(aq) + H2(g)

balance oxygen: Sr(s) + 2H2O(l)  Sr(OH)2(aq) + H2(g)

balanced equation: Sr(s) + 2H2O(l)  Sr(OH)2(aq) + H2(g)

Co(s) + HCl(aq)  CoCl2(aq) + H2(g)

balance hydrogen: Co(s) + 2HCl(aq)  CoCl2(aq) + H2(g)

balanced equation: Co(s) + 2HCl(aq)  CoCl2(aq) + H2(g)

51.

In a double displacement reaction, two ionic solutes “switch partners” with the positive io
n from one combining with the negative ion from the other to form the precipitate:  for ex
ample, in the reaction AgNO3(aq) + HCl(aq)    AgCl(s) + HNO3(aq), silver ion from on
e solute combines with chloride ion from the other solute to form the precipitate. In a sing
le displacement reaction, one element replaces another from its compound: in other word
s, a single displacement reaction is typically an oxidation–reduction reaction also:  for ex
ample in the reaction Zn(s) + CuSO4(aq)    Cu(s) + ZnSO4(aq), zinc in the elemental for
m replaces copper in the copper compound, producing copper in the elemental form and a
zinc compound. Many other examples are possible.

examples of formation of water:

HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
H2SO4(aq) + 2KOH(aq)  2H2O(l) + K2SO4(aq)

examples of formation of a gaseous product:

Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
2KClO3(s)  2KCl(s) + 3O2(g)

52.

A reaction must be an oxidation–reduction reaction if any of the oxidation numbers of the atoms i
n the equation change. Aluminum changes oxidation state from 0 in Al to +3 (oxidation) in Al2O3
and AlCl3; nitrogen changes oxidation state from –3 in NH4+ to +2 in NO (oxidation); chlorine ch
anges oxidation state from +7 in ClO4– to –1 in AlCl3 (reduction)

Chapter 7: Reactions in Aqueous Solution

53.

For each reaction, the type of reaction is first identified, followed by some of the reasoning that le
ads to this choice (there may be more than one way in which you can recognize a particular type
of reaction).

precipitation (from Table 7.1, BaSO4 is insoluble).

oxidation–reduction (Zn changes from the elemental to the combined state; hydrogen
changes from the combined to the elemental state).

precipitation (From Table 7.1, AgCl is insoluble.)

acid–base (HCl is an acid; KOH is a base; water and a salt are produced.)

oxidation–reduction (Cu changes from the combined to the elemental state; Zn changes
from the elemental to the combined state.)

acid–base (The H2PO4– ion behaves as an acid; NaOH behaves as a base; a salt and water
are produced.)

precipitation (From Table 7.1, CaSO4 is insoluble); acid–base [Ca(OH)2 is a base; H2SO4
is an acid; a salt and water are produced.]

oxidation–reduction (Mg changes from the elemental to the combined state; Zn changes
from the combined to the elemental state.)

precipitation (From Table 7.1, BaSO4 is insoluble.)

54.

oxidation–reduction (Oxygen changes from the combined state to the elemental state.)

oxidation–reduction (Zinc changes from the elemental to the combined state; hydrogen
changes from the combined to the elemental state.)

acid–base (H2SO4 is a strong acid and NaOH is a strong base; water and a salt are
formed.)

acid–base, precipitation (H2SO4 is a strong acid, and Ba(OH)2 is a base; water and a salt
are formed; an insoluble product forms.)

precipitation (From the Solubility Rules of Table 7.1, AgCl is only slightly soluble.)

precipitation (From the Solubility Rules of Table 7.1, Cu(OH)2 is only slightly soluble.)

oxidation–reduction (Chlorine and fluorine change from the elemental to the combined
state.)

oxidation–reduction (Oxygen changes from the elemental to the combined state.)

acid–base (HNO3 is a strong acid and Ca(OH)2 is a strong base; a salt and water are
formed.)

55.

A combustion reaction is typically a reaction in which an element or compound reacts with oxyge
n so quickly and with so much release of energy that a flame results. In addition to the carbon dio
xide and water chemical products, combustion reactions are a major source of heat energy.

56.

oxidation–reduction

Chapter 7: Reactions in Aqueous Solution

57.

A synthesis reaction represents the production of a given compound from simpler substances (eith
er elements or simpler compounds). For example,

O2(g) + 2F2(g)  2OF2(g)

represents a simple synthesis reaction. Synthesis reactions may often (but not necessarily always)
also be classified in other ways. For example, the reaction

C(s) + O2(g)  CO2(g)

could also be classified as an oxidation–reduction reaction or as a combustion reaction (a special
sub–classification of oxidation–reduction reaction that produces a flame). As another example,
the reaction

2Fe(s) + 3Cl2(g)  2FeCl3(s)

is a synthesis reaction that also is an oxidation–reduction reaction.

58.

A decomposition reaction is one in which a given compound is broken down into simpler compou
nds or constituent elements. The reactions

CaCO3(s)  CaO(s) + CO2(g)

2HgO(s)  2Hg(l) + O2(g)

both represent decomposition reactions. Such reactions often (but not necessarily always) may be
classified in other ways. For example, the reaction of HgO(s) is also an oxidation–reduction
reaction.

59.

Compounds like those in parts (a) and (b) of this problem, containing only carbon and hydrogen,
are called hydrocarbons. When a hydrocarbon is reacted with oxygen (O2), the hydrocarbon is al
most always converted to carbon dioxide and water vapor. Because water molecules contain an o
dd number of oxygen atoms, and O2 contains an even number of oxygen atoms, it is often difficul
t to balance such equations. For this reason, it is simpler to balance the equation using fractional c
oefficients if necessary, and then to multiply by a factor that will give whole number coefficients
for the final balanced equation.

C6H6 + O2  CO2 + H2O

Balance carbon: C6H6 + O2  6CO2 + H2O

Balance hydrogen: C6H6 + O2  6CO2 + 3H2O

Balance oxygen with fractional coefficient: C6H6 +

2 O2  6CO2 + 3H2O

Balanced equation: 2C6H6 + 15O2  12CO2 + 6H2O

C5H12 + O2  CO2 + H2O

Balance carbon: C5H12 + O2  5CO2 + H2O

Balance hydrogen: C5H12 + O2  5CO2 + 6H2O

Balance oxygen: C5H12 + 8O2  5CO2 + 6H2O

Balanced equation: C5H12 + 8O2  5CO2 + 6H2O

Chapter 7: Reactions in Aqueous Solution

C2H6O(l) + O2(g)  CO2 + H2O

Balance carbon: C2H6O(l) + O2(g)  2CO2 + H2O

Balance hydrogen: C2H6O(l) + O2(g)  2CO2 + 3H2O

Balance oxygen: C2H6O(l) + 3O2(g)  2CO2 + 3H2O

Balanced equation: C2H6O(l) + 3O2(g)  2CO2 + 3H2O

60.

Compounds like those in this problem, containing only carbon and hydrogen, are called hydrocar
bons. When a hydrocarbon is reacted with oxygen (O2), the hydrocarbon is almost always convert
ed to carbon dioxide and water vapor. Because water molecules contain an odd number of oxygen
atoms, and O2 contains an even number of oxygen atoms, it is often difficult to balance such equa
tions. For this reason, it is simpler to balance the equation using fractional coefficients if necessar
y, and then to multiply by a factor that will give whole number coefficients for the final balanced
equation.

C3H8(g) + O2(g)  CO2(g) + H2O(g)

balance carbon: C3H8(g) + O2(g)  3CO2(g) + H2O(g)

balance hydrogen: C3H8(g) + O2(g)  3CO2(g) + 4H2O(g)

balance oxygen: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

balanced equation: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

C2H4(g) + O2(g)  CO2(g) + H2O(g)

balance carbon: C2H4(g) + O2(g)  2CO2(g) + H2O(g)

balance hydrogen: C2H4(g) + O2(g)  2CO2(g) + 2H2O(g)

balance oxygen: C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(g)

balanced equation: C2H4(g) + 3O2(g)  2CO2(g) + 2H2O(g)

C4H10(g) + O2(g)  CO2(g) + H2O(g)

balance carbon: C4H10(g) + O2(g)  4CO2(g) + H2O(g)

balance hydrogen: C4H10(g) + O2(g)  4CO2(g) + 5H2O(g)

balance oxygen: C4H10(g) +

2 O2(g)  4CO2(g) + 5H2O(g)

balanced equation: 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)

61.

Specific examples will depend on the students’ input. A typical combustion reaction is represente
d by the reaction of methane (CH4) with oxygen gas

CH4(g) + 2O2(g)  CO2(g) + 2H2O(g).

62.

A reaction in which small molecules or atoms combine to make a larger molecule is called a synt
hesis reaction. An example would be the synthesis of sodium chloride from the elements

2Na(s) + Cl2(g)  2NaCl(s).

A reaction in which a molecule is broken down into simpler molecules or atoms is called a
decomposition reaction. An example would be the decomposition of sodium hydrogen carbonate
when heated.

2NaHCO3(s)  Na2CO3(s) + CO2(g) + H2O(g).

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Chapter 7: Reactions in Aqueous Solution

Specific examples will depend on the students’ input.

63.

CaO(s) + H2O(l)  Ca(OH)2(s)

4Fe(s) + 3O2(g)  2Fe2O3(s)

P2O5(s) + 3H2O(l)  2H3PO4(aq)

64.

8Fe(s) + S8(s)  8FeS(s)

4Co(s) + 3O2(g)  2Co2O3(s)

Cl2O7(g) + H2O(l)  2HClO4(aq)

65.

CaSO4(s)  CaO(s) + SO3(g)

Li2CO3(s)  Li2O(s) + CO2(g)

2LiHCO3(s)  Li2CO3(s) + H2O(g) + CO2(g)

C6H6(l)  6C(s) + 3H2(g)

4PBr3(l)  P4(s) + 6Br2(l)

66.

2Al(s) + 3Br2(l)  2AlBr3(s)

Zn(s) + 2HClO4(aq)  Zn(ClO4)2(aq) + H2(g)

3Na(s) + P(s)  Na3P(s)

CH4(g) + 4Cl2(g)  CCl4(l) + 4HCl(g)

Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)

67.

A molecular equation uses the normal, uncharged formulas for the compounds involved. The com
plete ionic equation shows the compounds involved broken up into their respective ions (all ions
present are shown). The net ionic equation shows only those ions that combine to form a precipita
te, a gas, or a nonionic product such as water. The net ionic equation shows most clearly the speci
es that are combining with each other.

68.

(c); (a), (b), and (d) will form insoluble compounds with Pb2+. For (e), a compound will not form
between Na+ and Pb2+.

69.

2Fe3+(aq) + 3CO32–(aq)  Fe2(CO3)3(s)

Hg22+(aq) + 2 Cl–(aq)  Hg2Cl2(s)

no precipitate

Cu2+(aq) + S2–(aq)  CuS(s)

Pb2+(aq) + 2Cl–(aq)  PbCl2(s)

Ca2+(aq) + CO32–(aq)  CaCO3(s)

Au3+(aq) + 3OH–(aq)  Au(OH)3(s)

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Chapter 7: Reactions in Aqueous Solution

70.

The formulas of the salts are indicated in boldface type.

HNO3(aq) + KOH(aq)  H2O(l) + KNO3(aq)

H2SO4(aq) + Ba(OH)2(aq)  2H2O(l) + BaSO4(s)

HClO4(aq) + NaOH(aq)  H2O(l) + NaClO4(aq)

2HCl(aq) + Ca(OH)2(aq)  2H2O(l) + CaCl2(aq)

71.

For each cation, the precipitates that form with the anions listed in the right–hand column are give
n below. If no formula is listed, it should be assumed that the anion does not form a precipitate wi
th the particular cation. See Table 7.1 for the Solubility Rules.

Ag+ ion:

AgCl, Ag2CO3, AgOH, Ag3PO4, Ag2S, Ag2SO4

Ba2+ ion:

BaCO3, Ba(OH)2, Ba3(PO4)2, BaS, BaSO4

Ca2+ ion:

CaCO3, Ca(OH)2, Ca3(PO4)2, CaS, CaSO4

Fe3+ ion:

Fe2(CO3)3, Fe(OH)3, FePO4, Fe2S3

Hg22+ ion:

Hg2Cl2, Hg2CO3, Hg2(OH)2, (Hg2)3(PO4)2, Hg2S

Na+ ion:

all common salts are soluble

Ni2+ ion:

NiCO3, Ni(OH)2, Ni3(PO4)2, NiS

Pb2+ ion:

PbCl2, PbCO3, Pb(OH)2, Pb3(PO4)2, PbS, PbSO4

72.

73.

iron(III) hydroxide, Fe(OH)3. Rule 5: Most hydroxide salts are only slightly soluble.

nickel(II) sulfide, NiS. Rule 6: Most sulfide salts are only slightly soluble.

silver chloride, AgCl. Rule 3: Although Most chloride salts are soluble, AgCl is a listed
exception

barium carbonate, BaCO3. Rule 6: Most carbonate salts are only slightly soluble.

mercury(I) chloride or mercurous chloride, Hg2Cl2. Rule 3: Although Most chloride salts
are soluble, Hg2Cl2 is a listed exception

barium sulfate, BaSO4. Rule 4: Although Most sulfate salts are soluble, BaSO4 is a listed
exception

74.

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Chapter 7: Reactions in Aqueous Solution

75.

Rule 3: Ag+(aq) + Cl–(aq)  AgCl(s)

Rule 6: 3Ca2+(aq) + 2PO43–(aq)  Ca3(PO4)2(s)

Rule 3: Pb2+(aq) + 2Cl–(aq)  PbCl2(s)

Rule 6: Fe3+(aq) + 3OH–(aq)  Fe(OH)3(s)

76.

Molecular:

Complete Ionic:

Net Ionic:

77.

potassium hydroxide and perchloric acid

cesium hydroxide and nitric acid

potassium hydroxide and hydrochloric acid

sodium hydroxide and sulfuric acid

78.

Aluminum atoms lose 3 electrons to become Al3+ ions. Iodine atoms gain 1 electron each to beco
me I– ions.

79.

Fe2S3 is made up of Fe3+ and S2– ions. Iron atoms each lose three electrons to become Fe3+ ions. Su
lfur atoms each gain two electrons to become S2– ions.

80.

Na + O2  Na2O2

Balance sodium: 2Na + O2  Na2O2

Balanced equation: 2Na(s) + O2(g)  Na2O2(s)

Fe(s) + H2SO4(aq)  FeSO4(aq) + H2(g)

Equation is already balanced!

Al2O3  Al + O2

Balance oxygen: 2Al2O3  Al + 3O2

Balance aluminum: 2Al2O3  4Al + 3O2

Balanced equation: 2Al2O3(s)  4Al(s) + 3O2(g)

Fe + Br2  FeBr3

Balance bromine: Fe + 3Br2  2FeBr3

Balance iron: 2Fe + 3Br2  2FeBr3

Balanced equation: 2Fe(s) + 3Br2(l)  2FeBr3(s)

Zn + HNO3  Zn(NO3)2 + H2

Balance nitrate ions: Zn + 2HNO3  Zn(NO3)2 + H2

Balanced equation: Zn(s) + 2HNO3(aq)  Zn(NO3)2(aq) + H2(g)

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Chapter 7: Reactions in Aqueous Solution

81.

oxidation–reduction (Mg changes from the elemental state to the combined state in
MgSO4; hydrogen changes from the combined to the elemental state.)

acid–base (HClO4 is a strong acid and RbOH is a strong base; water and a salt are
produced.)

oxidation–reduction (Both Ca and O2 change from the elemental to the combined state.)

acid–base (H2SO4 is a strong acid and NaOH is a strong base; water and a salt are
produced.)

precipitation (From the Solubility Rules of Table 7.1, PbCO3 is insoluble.)

precipitation (From the Solubility Rules of Table 7.1, CaSO4 is insoluble.)

acid–base (HNO3 is a strong acid and KOH is a strong base; water and a salt are
produced.)

precipitation (From the Solubility Rules of Table 7.1, NiS is insoluble.)

oxidation–reduction (both Ni and Cl2 change from the elemental to the combined state).

82.

(a), (b), and (c); All three statements are true.

83.

4FeO(s) + O2(g)  2Fe2O3(s)

2CO(g) + O2(g)  2CO2(g)

H2(g) + Cl2(g)  2HCl(g)

16K(s) + S8(s)  8K2S(s)

6Na(s) + N2(g)  2Na3N(s)

84.

2NaHCO3(s)  Na2CO3(s) + H2O(g) + CO2(g)

2NaClO3(s)  2NaCl(s) + 3O2(g)

2HgO(s)  2Hg(l) + O2(g)

C12H22O11(s)  12C(s) + 11H2O(g)

2H2O2(l)  2H2O(l) + O2(g)

85.

For simplicity, the physical states of the substances are omitted.

2Ba + O2  2BaO

Ba + S  BaS

Ba + Cl2  BaCl2

3Ba + N2  Ba3N2

Ba + Br2  BaBr2

4K + O2  2K2O

2K + S  K2S

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Chapter 7: Reactions in Aqueous Solution

2K + Cl2  2KCl

6K + N2  2K3N

2K + Br2  2KBr

2Mg + O2  2MgO

Mg + S  MgS

Mg + Cl2  MgCl2

3Mg + N2  Mg3N2

Mg + Br2  MgBr2

4Rb + O2  2Rb2O

2Rb + S  Rb2S

2Rb + Cl2  2RbCl

6Rb + N2  2Rb3N

2Rb + Br2  2RbBr

2Ca + O2  2CaO

Ca + S  CaS

Ca + Cl2  CaCl2

3Ca + N2  Ca3N2

Ca + Br2  CaBr2

4Li + O2  2Li2O

2Li + S  Li2S

2Li + Cl2  2LiCl

6Li + N2  2Li3N

2Li + Br2  2LiBr

86.

87.

For simplicity, the physical states of the substances are omitted.

Mg + Cl2  MgCl2

Ca + Cl2  CaCl2

Sr + Cl2  SrCl2

Ba + Cl2  BaCl2

Mg + Br2  MgBr2

Ca + Br2  CaBr2

Sr + Br2  SrBr2

Ba + Br2  BaBr2

2Mg + O2  2MgO

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Chapter 7: Reactions in Aqueous Solution

2Ca + O2  2CaO

2Sr + O2  2SrO

2Ba + O2  2BaO

88.

one

two

three

89.

two; O + 2e–  O2–

one; F + e–  F–

three; N + 3e–  N3–

one; Cl + e–  Cl–

two; S + 2e–  S2–

90.

False. The balanced molecular equation is: Ba(OH)2(aq) + H2SO4(aq)  BaSO4(s) + 2H2O(l). The
complete ionic equation is: Ba2+(aq) + 2OH–(aq) + 2H+(aq) + SO42–(aq)  BaSO4(s) + 2H2O(l). T
he net ionic equation includes all species that take part in the chemical reaction. The OH– and H+ i
ons form water so they are also included in the net ionic equation. Thus the complete ionic equati
on and net ionic equation are the same.

91.

2I4O9(s)  2I2O6(s) + 2I2(s) + 3O2(g)

oxidation–reduction, decomposition

Mg(s) + 2AgNO3(aq)  Mg(NO3)2(aq) + 2Ag(s)

oxidation–reduction, single–displacement

SiCl4(l) + 2Mg(s)  2MgCl2(s) + Si(s)

oxidation–reduction, single–displacement

CuCl2(aq) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2AgCl(s)

precipitation, double–displacement

2Al(s) + 3Br2(l)  2AlBr3(s)

oxidation–reduction, synthesis

92.

93.

2Zn(s) + O2(g)  2ZnO(s)

4Al(s) + 3O2(g)  2Al2O3(s)

2Fe(s) + O2(g)  2FeO(s); 4Fe(s) + 3O2(g)  2Fe2O3(s)

2Cr(s) + O2(g)  2CrO(s); 4Cr(s) + 3O2(g)  2Cr2O3(s)

2Ni(s) + O2(g)  2NiO(s)

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Chapter 7: Reactions in Aqueous Solution

94.

PbCl2; PbSO4; Pb3(PO4)2; AgCl; Ag3PO4

95.

You have your choice of reactions here as long as they illustrate the correct reaction type. Those l
isted below are only examples:

C(s) + O2(g)  CO2(g)

AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)

C(s) + O2(g)  CO2(g)

Note that some examples are repeated: a given reaction may sometimes be classified as more than
one type of reaction. The reaction between carbon and oxygen gas, for example, is at the same
time a combustion reaction (burning in oxygen), an oxidation–reduction reaction (the oxidation
numbers of carbon and oxygen both change) and a synthesis reaction (two smaller entities unite
to form a larger, more complex molecule)

96.

PbSO4; AgCl; none

97.

Sr3(PO4)2; Ag2CO3; none; AgCl; PbCl2

107