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Chapter 3
Chemical Reactions
Chapter 3
Chemical Reactions
INSTRUCTOR’S NOTES
Chapter 3 is focused on chemical reactions, especially those that take place in aqueous solution. The chapter has a
heavy emphasis on chemistry at the molecular level. The images representing electrolytic behavior and chemical
reactions at the molecular level are intended to help students make the transition to visualize molecules instead of
chemical formulas.
We consider Chapter 3 an important early chapter in the book. We cover the material at this point in our course
because our students use the concepts in the laboratory. However, we emphasize that it is not crucial that this
material be covered at this time in the course. Some of the sections can be delayed, in particular, it would be easy to
cover the subject of redox reactions (Section 3.9) at some other point in the year, such as when covering
electrochemistry.
Chapter 3 introduces the properties of ionic compounds in aqueous solution and general guidelines to predict the
aqueous solubility of simple ionic compounds. Although one can readily identify many exceptions to the guidelines
in Figure 3.10, we do find them useful for students in a beginning course. We hope that the many photos of soluble
and insoluble compounds in this chapter will help students remember some of these guidelines. Demos can also
make this concept concrete for the students.
The section on net ionic equations (Section 3.6) can be a difficult one for students. We admit it does take time and
effort to help the students come to grips with this concept. However, it is a useful one, since (1) many reactions are
best seen in this fashion and (2) balancing reactions is made easier. Students can generally write balanced equations
for exchange reactions, but they sometimes have a difficult time turning them into net ionic equations. Time and
effort on everyone's part is needed here.
The introduction to acids and bases (Section 3.7) is somewhat brief at this stage, but it does enable one to use
common acids and bases in examples and in the laboratory with the knowledge that students have some familiarity
with them. New to this edition is the introduction of the Brønsted-Lowry definition of acids and bases, which
prepares students for the more advanced acid-base coverage later in the book, and the introduction of the hydronium
ion, H3O+ at this point in the text. Also, in this edition pH calculations and the pH scale have been moved to
Chapter 4.
Oxidation numbers have been introduced only as a way of telling if a redox reaction has occurred. Balancing redox
reactions is covered in Chapter 20.
Chapter 3 requires approximately three to four lectures.
47
Chapter 3
Chemical Reactions
SUGGESTED DEMONSTRATIONS
1.
Illustrations of Chemical Reactions and Balancing Equations

Burn sulfur in oxygen.

Make NO2 from a penny and concentrated HNO3.

Burn magnesium in air.

Decompose H2O2.

React various metals with HCl or H2O.

Wright, S. W. “A Method for Generating Oxygen from Consumer Chemicals,” Journal of Chemical
Education 2003, 80, 1158.

Keiter, R. L.; Gamage, C. P. “Combustion of White Phosphorous,” Journal of Chemical Education 2001,
78, 908.

Senkbeil, E. G. “Combustion Demonstration Using Updated Flame Tornado,” Journal of Chemical
Education 2000, 77, 1449.
2.
Electrolytes

We always illustrate the conductivity, and the difference between strong and weak electrolytes, using an
apparatus such as that in the Chapter Focus. We use solutions such as CuSO4, Na2CO3, vinegar, and pure
water. It is interesting to compare the behavior of these solutions with the conductivity of soda or juice.

Haworth, D. T.; Bartelt, M. R.; Kenney, M. J. “Solution Conductivity Apparatus,” Journal of Chemical
Education 1999, 76, 625.

3.
Cortel, A. “Fast Ionic Migration of Copper Chromate,” Journal of Chemical Education 2001, 78, 207.
Solubility of Ionic Compounds

It is very important to illustrate the solubility of salts. Our most recent demonstrations are outlined below.
Throughout the demonstrations the solubility guidelines are projected onto the screen in the lecture room.
Note: Precipitation reactions cannot be illustrated on an overhead projector. The precipitate is seen only
as a dark blob. However, even in a large lecture room, students can see the reaction if it is done in a large
flask (1-2 L).
(a) A gram or so of KI is dissolved in a few hundred milliliters of water in one flask, and Pb(NO3)2 is
dissolved in water in another flask. The students can clearly see the salts dissolve. On mixing the
solutions, though, a bright yellow precipitate of PbI 2 appears, clearly demonstrating one of the
exceptions to the general rule that halide salts are water-soluble.
(b) BaCl2 is dissolved in water while discussing halide solubility and mixed with a solution of CuSO 4
(which shows the usual solubility of sulfates). Precipitation of BaSO 4 on mixing the solutions then
demonstrates the exceptions to the guideline regarding sulfates.
(c) Our attempt at dissolving a piece of blackboard chalk fails, clearly showing that many carbonates are
not soluble.
48
Chapter 3

Chemical Reactions
Other solubility demonstrations include:
(a) Dissolve NaCl in water and then precipitate AgCl with AgNO3.
(b) Dissolve Na2CO3 and compare this with chalk.
(c) Show a sample of fool's gold (iron pyrite) to illustrate the insolubility of metal sulfides.
(d) Add NaOH to solutions of Fe(NO3)3 and CuSO4 to get insoluble hydroxides.
(e) deVos, W. “Using Large Glass Cylinders to Demonstrate Chemical Reactions,” Journal of Chemical
Education 1999, 76, 528.
(f) Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of
Wisconsin Press, 1983; Vol. 1, pp. 307-313.
4.
Acid-Base Reactions

Although is it easy for students to see precipitation reactions done in large flasks, we find that acid-base
reactions are difficult to demonstrate, since common ones are not spectacular. However, a suggested
demonstration would be to dissolve a precipitate of Ca(OH) 2. At least the solid would be seen to dissolve
on adding acid.
5.
Oxides

Show samples of metal oxides, especially those of iron, aluminum, boron, lead, and magnesium. Show a
sample of anodized aluminum or aluminum oxide sandpaper.

Show samples of silica sand and quartz crystals.

Take a sample of dry ice and discuss some of the chemistry of CO2.

React hot sulfur with pure O2 (Shakhashiri, Vol. 2, pp. 184-189). This is especially effective in a very dark
room. (Fill the bottle with pure O2 before lecture and tightly stopper.)

Make NO2 from copper (a penny) and concentrated HNO3. (Place a penny in about 15 mL of concentrated
HNO3 in a 2-L flask and stopper lightly with a rubber stopper. Although a little NO 2 comes out of the flask,
in a large lecture room it is not noticeable.)
6.
Gas-Forming Reactions

Derr, H. R.; Lewis, T.; Derr, B. J. “Gas Me Up, or, A Baking Powder Diver,” Journal of Chemical
Education 2000, 77, 171.
7.
Oxidation-Reduction Chemistry

After introducing the ideas of redox and the concept of oxidation numbers, we do a demonstration on the
oxidation states of vanadium.
Directions: Dissolve 1 g of ammonium vanadate in 200 mL of water. (Warm to dissolve completely.)
(a) Add 3 M H2SO4 (20-50 mL), and the solution turns yellow owing to an acid-base reaction
VO3– + 2 H3O+  VO2+ + 3 H2O
(b) Add a handful of zinc chips. These reduce the V(V) to V(IV).
2 VO2+ + Zn + 4 H3O+  2 VO2+ (vanadyl, blue) + 6 H2O + Zn2+
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Chapter 3
Chemical Reactions
(c) The solution will slowly turn emerald green owing to another reduction step.
2 VO2+ (blue) + Zn + 4 H+  2 V3+ (green) + 2 H2O + Zn2+
(d) The solution will finally turn violet. We allow the reaction to continue throughout the lecture with a
stopper on the flask. This excludes air well enough that the violet, aqueous vanadium(II) ion can form.
(The vanadium(II) ion is readily oxidized by oxygen in air to V 3+. Therefore, to see the violet color air
must be excluded.)

Wellman, W. E.; Noble, M. E. “Greening the Blue Bottle,” Journal of Chemical Education 2003, 80, 537.

Volkovich, V. A.; Griffiths, T. R. “Catalytic Oxidation of Ammonia: A Sparkling Experiment,” Journal of
Chemical Education 2000, 77, 177.

Elsworth, J. F. “Entertaining Chemistry—Two Colorful Reactions,” Journal of Chemical Education 2000,
77, 484.

Eliason, R.; Lee, E. J.; Wakefield, D.; Bergren, A. “Improvement of Sugar-Chlorate Rocket
Demonstration,” Journal of Chemical Education 2000, 77, 1580.

Maya, H. D.; Neves, E. A. “A Further Demonstration of Sulfite-Induced Redox Cycling of Metal Ions
Initiated by Shaking,” Journal of Chemical Education 1999, 76, 930.
50
Chapter 3
Chemical Reactions
SOLUTIONS TO STUDY QUESTIONS
3.1
(a) 2 Al(s) + Fe2O3(s)  Al2O3(s) + 2 Fe()
So much energy is released as heat in this reaction that the iron formed is in the liquid state.
(b) C(s) + H2O(g)  CO(g) + H2(g)
(c) SiCl4() + 2 Mg (s)  Si(s) + 2 MgCl2(s)
3.2
(a) N2(g) + 3 H2(g)  2 NH3(g)
(b) 2 H2(g) + CO(g)  CH3OH()
(c) 2 S(s) + 3 O2(g) + 2 H2O()  2 H2SO4()
3.3
(a) 4 Cr(s) + 3 O2(g)  2 Cr2O3(s)
(b) Cu2S(s) + O2(g)  2 Cu(s) + SO2(g)
(c) C6H5CH3() + 9 O2(g)  4 H2O() + 7 CO2(g)
3.4
(a) 2 Cr(s) + 3 Cl2(g)  2 CrCl3(s)
(b) SiO2(s) + 2 C(s)  Si(s) + 2 CO(g)
(c) 3 Fe(s) + 4 H2O(g)  Fe3O4(s) + 4 H2(g)
3.5
(a) Fe2O3(s) + 3 Mg(s)  3 MgO(s) + 2 Fe(s)
iron(III) oxide, magnesium, magnesium oxide, iron
(b) AlCl3(s) + 3 NaOH(aq)  Al(OH)3(s) + 3 NaCl(aq)
aluminum chloride, sodium hydroxide, aluminum hydroxide, sodium chloride
(c) 2 NaNO3(s) + H2SO4(aq)  Na2SO4(s) + 2 HNO3(aq)
sodium nitrate, sulfuric acid, sodium sulfate, nitric acid
(d) NiCO3(s) + 2 HNO3(aq)  Ni(NO3)2(aq) + CO2(g) + H2O()
nickel(II) carbonate, nitric acid, nickel(II) nitrate, carbon dioxide, water
3.6
(a) SF4(g) + 2 H2O()  SO2(g) + 4 HF()
sulfur tetrafluoride, water, sulfur dioxide, hydrogen fluoride
(b) 4 NH3(aq) + 5 O2(g)  4 NO(g) + 6 H2O()
ammonia, oxygen, nitrogen monoxide, water
(c) BF3(g) + 3 H2O()  3 HF(aq) + H3BO3(aq)
boron trifluoride, water, hydrogen fluoride, hydrogen borate (boric acid)
3.7
The reaction involving HCl is more product-favored at equilibrium.
3.8
H3BO3(aq) + H2O()
H3O+(aq) + H2BO3-(aq)
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Chapter 3
Chemical Reactions
H3PO4(aq) + H2O()
H3O+(aq) + H2PO4-(aq)
The phosphoric acid dissociation is more product-favored than the dissociation of the boric acid, producing
more ions in solutions and the greater current.
3.9
Electrolytes are compounds whose aqueous solutions conduct electricity. Substances whose solutions are
good electrical conductors are strong electrolytes (such as sodium chloride), poor electrical conductors are
weak electrolytes (such as acetic acid).
3.10
Hydrochloric acid, HCl, and nitric acid, HNO3, are strong electrolytes. Acetic acid, HO2CCH3, is a weak
electrolyte. Sodium hydroxide, NaOH, and potassium hydroxide, KOH, are strong electrolytes. Ammonia,
NH3, is a weak electrolyte.
3.11
(a) CuCl2
(b) AgNO3
(c) K2CO3, KI, KMnO4
3.12
(a) Ba(NO3)2
(b) Na2SO4, NaClO4, NaCH3CO2
(c) KBr, Al2Br6
3.13
3.14
3.15
3.16
3.17
(a)
K+ and OH– ions
(c)
Li+ and NO3– ions
(b)
K+ and SO42– ions
(d)
NH4+ and SO42– ions
(a)
K+ and I– ions
(c)
K+ and HPO42– ions
(b)
Mg2+ and CH3CO2– ions
(d)
Na+ and CN– ions
(a)
soluble; Na+ and CO32– ions
(c)
insoluble
(d)
soluble; Ba2+ and Br– ions
2+
SO42–
(b)
soluble; Cu and
ions
(a)
soluble; Ni2+ and Cl– ions
(c)
soluble; Pb2+ and NO3– ions
(b)
soluble; Cr2+ and NO3– ions
(d)
insoluble
CdCl2(aq) + 2 NaOH(aq)  Cd(OH)2(s) + 2 NaCl(aq)
Cd2+(aq) + 2 OH–(aq)  Cd(OH)2(s)
3.18
Ni(NO3)2(aq) + Na2CO3(aq)  NiCO3(s) + 2 NaNO3(aq)
Ni2+(aq) + CO32–(aq)  NiCO3(s)
3.19
(a)
NiCl2(aq) + (NH4)2S(aq)  NiS(s) + 2 NH4Cl(aq)
Ni2+(aq) + S2–(aq)  NiS(s)
(b)
3 Mn(NO3)2(aq) + 2 Na3PO4(aq)  Mn3(PO4)2(s) + 6 NaNO3(aq)
3 Mn2+(aq) + 2 PO43–(aq)  Mn3(PO4)2(s)
52
Chapter 3
3.20
Chemical Reactions
(a) Pb(NO3)2(aq) + 2 KBr(aq)  PbBr2(s) + 2 KNO3(aq)
Pb2+(aq) + 2 Br–(aq)  PbBr2(s)
(b) Ca(NO3)2(aq) + 2 KF(aq)  CaF2(s) + 2 KNO3(aq)
Ca2+(aq) + 2 F–(aq)  CaF2(s)
(c) Ca(NO3)2(aq) + Na2C2O4(aq)  CaC2O4(s) + 2 NaNO3(aq)
Ca2+(aq) + C2O42–(aq)  CaC2O4(s)
3.21
HNO3(aq) + H2O()  H3O+(aq) + NO3–(aq)
3.22
HClO4(aq) + H2O()  H3O+(aq) + ClO4–(aq)
3.23
H2C2O4(aq) + H2O()  H3O+(aq) + HC2O4–(aq)
HC2O4–(aq) + H2O()  H3O+(aq) + C2O42–(aq)
3.24
H3PO4(aq) + H2O()  H3O+(aq) + H2PO4–(aq)
H2PO4–(aq) + H2O()  H3O+(aq) + HPO42–(aq)
HPO42–(aq) + H2O()  H3O+(aq) + PO43–(aq)
3.25
MgO(s) + H2O()  Mg(OH)2(s)
3.26
SO3(g) + H2O()  H2SO4(aq)
3.27
(a) 2 CH3CO2H(aq) + Mg(OH)2(s)  Mg(CH3CO2)2(aq) + 2 H2O()
acetic acid, magnesium hydroxide, magnesium acetate, water
(b) HClO4(aq) + NH3(aq)  NH4ClO4(aq)
perchloric acid, ammonia, ammonium perchlorate
3.28
(a) H3PO4(aq) + 3 KOH(aq)  K3PO4(aq) + 3 H2O()
phosphoric acid, potassium hydroxide, potassium phosphate, water
(b) H2C2O4(aq) + Ca(OH)2(s)  CaC2O4(s) + 2 H2O()
oxalic acid, calcium hydroxide, calcium oxalate, water
3.29
Ba(OH)2(aq) + 2 HNO3(aq)  Ba(NO3)2(aq) + 2 H2O()
3.30
2 Al(OH)3(s) + 3 H2SO4(aq)  Al2(SO4)3(aq) + 6 H2O()
3.31
HNO3(aq) + H2O()
NO3-(aq) + H3O+(aq)
Brønsted acids: HNO3(aq) and H3O+(aq)
Brønsted bases: H2O() and NO3-(aq)
HNO3 is a strong acid; therefore, this reaction is product-favored at equilibrium.
53
Chapter 3
3.32
Chemical Reactions
C6H5CO2H(aq) + H2O()
C6H5CO2-(aq) + H3O+(aq)
Brønsted acids: C6H5CO2H (aq) and H3O+(aq)
Brønsted bases: H2O() and C6H5CO2-(aq)
3.33
3.34
H2O() + HBr(aq)
H3O+(aq) + Br-(aq)
H2O() + NH3(aq)
OH-(aq) + NH4+(aq)
H2PO4-(aq) + CO32-(aq)
HPO4-(aq) + HCO3-(aq)
HPO4-(aq) + CH3CO2H(aq)
3.35
H2PO4-(aq) + CH3CO2-(aq)
(a) (NH4)2CO3(aq) + Cu(NO3)2  CuCO3(s) + 2 NH4NO3(aq)
CO32–(aq) + Cu2+(aq)  CuCO3(s)
(b) Pb(OH)2(s) + 2 HCl(aq)  PbCl2(s) + 2 H2O()
Pb(OH)2(s) + 2 H3O+(aq) + 2 Cl–(aq)  PbCl2(s) + 4 H2O()
(c) BaCO3(s) + 2 HCl(aq)  BaCl2(aq) + H2O() + CO2(g)
BaCO3(s) + 2 H3O+(aq)  Ba2+(aq) + 3 H2O() + CO2(g)
(d) 2 CH3CO2H(aq) + Ni(OH)2(s)  Ni(CH3CO2)2(aq) + 2 H2O()
2 CH3CO2H(aq) + Ni(OH)2(s)  Ni2+(aq) + 2 CH3CO2–(aq) + 2 H2O()
3.36
(a) Zn(s) + 2 HCl(aq)  H2(g) + ZnCl2(aq)
Zn(s) + 2 H3O+(aq)  H2(g) + Zn2+(aq) + 2 H2O()
(b) Mg(OH)2(s) + 2 HCl(aq)  MgCl2(aq) + 2 H2O()
Mg(OH)2(s) + 2 H3O+(aq)  Mg2+(aq) + 4 H2O()
(c) 2 HNO3(aq) + CaCO3(s)  Ca(NO3)2(aq) + H2O() + CO2(g)
2 H3O+(aq) + CaCO3(s)  Ca2+(aq) + 3 H2O() + CO2(g)
(d) (NH4)2S(aq) + FeCl2(aq)  2 NH4Cl(aq) + FeS(s)
S2–(aq) + Fe2+(aq)  FeS(s)
3.37
(a) AgNO3(aq) + KI(aq)  AgI(s) + KNO3(aq)
Ag+(aq) + I–(aq)  AgI(s)
(b) Ba(OH)2(aq) + 2 HNO3(aq)  Ba(NO3)2(aq) + 2 H2O()
OH–(aq) + H3O+(aq)  2 H2O()
(c) 2 Na3PO4(aq) + 3 Ni(NO3)2(aq)  Ni3(PO4)2(s) + 6 NaNO3(aq)
2 PO43–(aq) + 3 Ni2+(aq)  Ni3(PO4)2(s)
54
Chapter 3
3.38
Chemical Reactions
(a) 2 NaOH(aq) + FeCl2(aq)  Fe(OH)2(s) + 2 NaCl(aq)
2 OH–(aq) + Fe2+(aq)  Fe(OH)2(s)
(b) BaCl2(aq) + Na2CO3(aq)  BaCO3(s) + 2 NaCl(aq)
Ba2+(aq) + CO32–(aq)  BaCO3(s)
(c) 3 NH3(aq) + H3PO4(aq)  (NH4)3PO4(aq)
3 NH3(aq) + H3PO4(aq)  3 NH4+(aq) + PO43–(aq)
3.39
(a) HNO2(aq) + OH-(aq)  H2O() + NO2-(aq)
(b) Ca(OH)2(s) + 2 H3O+(aq)  4H2O() + Ca2+(aq)
3.40
(a) Ag+(aq) + I-(aq)  AgI(s)
(b) Ba2+(aq) + CO32-(aq)  BaCO3(s)
3.41
FeCO3(s) + 2 HNO3(aq)  Fe(NO3)2(aq) + CO2(g) + H2O()
iron(II) nitrate, carbon dioxide, water
3.42
MnCO3(s) + 2 HCl(aq)  MnCl2(aq) + CO2(g) + H2O()
manganese(II) chloride, carbon dioxide, water
3.43
(NH4)2S(s) + 2 HBr(aq)  2 NH4Br(aq) + H2S(g)
ammonium sulfide, hydrobromic acid, ammonium bromide, hydrogen sulfide
3.44
Na2SO3(aq) + 2 CH3CO2H(aq)  2 NaCH3CO2(aq) + SO2(g) + H2O()
sodium sulfite, acetic acid, sodium acetate, sulfur dioxide, water
3.45
3.46
3.47
(a) Br is +5 and O is –2
(d)
Ca is +2 and H is –1
(b) C is +3 and O is –2
(e)
H is +1, Si is +4, and O is –2
(c) F is –1
(f)
H is +1, S is +6, and O is –2
(a) P is +5 and F is –1
(d)
N is +5 and O is –2
(b) H is +1, As is +5, and O is –2
(e)
P is +5, O is –2, and Cl is –1
(c) U is +4 and O is –2
(f)
Xe is +6 and O is –2
(a) oxidation-reduction reaction
Oxidation number of Zn changes from 0 to +2, while that of N changes from +5 to +4
(b) acid-base reaction
(c) oxidation-reduction reaction
Oxidation number of Ca changes from 0 to +2, while that of H changes from +1 to 0
3.48
(a) precipitation reaction
(b) oxidation-reduction reaction
Oxidation number of Ca changes from 0 to +2, while that of O changes from 0 to –2
55
Chapter 3
Chemical Reactions
(c) oxidation-reduction reaction
Oxidation number of Fe changes from +2 to +3, while that of O changes from 0 to –2
3.49
(a) C2H4 is oxidized and is the reducing agent; O2 is reduced and is the oxidizing agent
(b) Si is oxidized and is the reducing agent; Cl2 is reduced and is the oxidizing agent
3.50
(a) Cr2O72– is reduced and is the oxidizing agent; Sn2+ is oxidized and is the reducing agent.
(b) FeS is oxidized and is the reducing agent; NO3– is reduced and is the oxidizing agent.
3.51
(a) Ba(OH)2(aq) + 2 HCl(aq)  BaCl2(aq) + 2 H2O()
acid-base reaction
(b) 2 HNO3(aq) + CoCO3(s)  Co(NO3)2(aq) + H2O() + CO2(g)
gas-forming reaction
(c) 2 Na3PO4(aq) + 3 Cu(NO3)2(aq)  Cu3(PO4)2(s) + 6 NaNO3(aq)
precipitation reaction
3.52
(a) K2CO3(aq) + Cu(NO3)2(aq)  CuCO3(s) + 2 KNO3(aq)
precipitation reaction
(b) Pb(NO3)2(aq) + 2 HCl(aq)  PbCl2(s) + 2 HNO3(aq)
precipitation reaction
(c) MgCO3(s) + 2 HCl(aq)  MgCl2(aq) + CO2(g) + H2O()
gas-forming reaction
3.53
(a) MnCl2(aq) + Na2S(aq)  MnS(s) + 2 NaCl(aq)
precipitation reaction
Mn2+(aq) + S2–(aq)  MnS(s)
(b) K2CO3(aq) + ZnCl2(aq)  ZnCO3(s) + 2 KCl(aq)
precipitation reaction
CO32–(aq) + Zn2+(aq)  ZnCO3(s)
3.54
(a) Fe(OH)3(s) + 3 HNO3(aq)  Fe(NO3)3(aq) + 3 H2O()
acid-base reaction
Fe(OH)3(s) + 3 H3O+(aq)  Fe3+(aq) + 6 H2O()
(b) FeCO3(s) + 2 HNO3(aq)  Fe(NO3)2(aq) + CO2(g) + H2O()
gas-forming reaction
FeCO3(s) + 2 H3O+(aq)  Fe2+(aq) + CO2(g) + 3 H2O()
3.55
(a) CuCl2(aq) + H2S(aq)  CuS(s) + 2 HCl(aq)
precipitation reaction
(b) H3PO4(aq) + 3 KOH(aq)  3 H2O() + K3PO4(aq)
56
Chapter 3
Chemical Reactions
acid-base reaction
(c) Ca(s) + 2 HBr(aq)  H2(g) + CaBr2(aq)
oxidation-reduction reaction, gas-forming reaction
(d) MgCl2(aq) + 2 NaOH(aq)  Mg(OH)2(s) + 2 NaCl(aq)
precipitation reaction
3.56
(a) NiCO3(s) + H2SO4(aq)  NiSO4(aq) + H2O() + CO2(g)
gas-forming reaction
(b) Co(OH)2(s) + 2 HBr(aq)  CoBr2(aq) + 2 H2O()
acid-base reaction
(c) AgCH3CO2(aq) + NaCl(aq)  AgCl(s) + NaCH3CO2(aq)
precipitation reaction
(d) NiO(s) + CO(g)  Ni(s) + CO2(g)
oxidation-reduction reaction
3.57
(a) Ca(OH)2(s) + 2 HBr(aq)  2H2O() + CaBr2(aq)
(b) MgCO3(s) + 2HNO3(aq)  Mg(NO3)2(aq) + CO2(g) + H2O()
(c) BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq)
(d) NH3(aq) + H2O()  NH4+(aq) + OH-(aq)
3.58
(a) 2 NH3(aq) + H2SO4(aq)  (NH4)2SO4(aq)
(b) CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O()
(c) 2AgNO3(aq) + BaCl2(aq)  Ba(NO3)2(aq) + 2AgCl(s)
(d) HClO4(aq) + H2O()  H3O+(aq) + ClO4-(aq)
3.59
(a) CO2(g) + 2 NH3(g)  NH2CONH2(s) + H2O()
(b) UO2(s) + 4 HF(aq)  UF4(s) + 2 H2O()
UF4(s) + F2(g)  UF6(s)
(c) TiO2(s) + 2 Cl2(g) + 2 C(s)  TiCl4() + 2 CO(g)
TiCl4() + 2 Mg(s)  Ti(s) + 2 MgCl2(s)
3.60
(a) Ca3(PO4)2(s) + 2 H2SO4(aq)  Ca(H2PO4)2(aq) + 2 CaSO4(s)
(b) 2 NaBH4(s) + H2SO4(aq)  B2H6(g) + 2 H2(g) + Na2SO4(aq)
(c) WO3(s) + 3 H2(g)  W(s) + 3 H2O()
(d) (NH4)2Cr2O7(s)  N2(g) + 4 H2O() + Cr2O3(s)
57
Chapter 3
3.61
Chemical Reactions
(a) One possible answer is NaBr
(b) One possible answer is Cu(OH)2
(c) One possible answer is CaCO3
(d) One possible answer is KNO3
(e) One possible answer is CH3COOH
3.62
(a) One possible answer is NaCH3CO2
(b) One possible answer is NiS
(c) One possible answer is NaOH
(d) One possible answer is PbCl2
(e) One possible answer is NaOH
3.63
Cu(NO3)2 and CuCl2 are soluble in water, CuCO3 and Cu3(PO4)2 are insoluble in water
3.64
NO3–, nitrate ion, and ClO4–, perchlorate ion
3.65
Nitrate ions are spectator ions in this acid-base reaction
2 H3O+(aq) + Mg(OH)2(s)  4 H2O() + Mg2+(aq)
3.66
(a) CuS(s)
copper(II) sulfide
Cu2+(aq) + S2–(aq)  CuS(s)
(b) CaCO3(s)
calcium carbonate
Ca2+(aq) + CO32–(aq)  CaCO3(s)
(c) AgI(s)
silver iodide
Ag+(aq) + I–(aq)  AgI(s)
3.67
(a) Cl2, chlorine, has been reduced and NaBr, sodium bromide, has been oxidized
(b) Cl2, chlorine, is the oxidizing agent and NaBr, sodium bromide, is the reducing agent
3.68
Oxidizing agents: HNO3, Cl2, O2, KMnO4
Reducing agent: Na
3.69
(a) MgCO3(s) + 2 H3O+(aq)  CO2(g) + Mg2+(aq) + 3 H2O()
Chloride ion, Cl–, is the spectator ion in this reaction.
(b) gas-forming reaction
3.70
(a) (NH4)2S(aq) + Hg(NO3)2(aq)  HgS(s) + 2 NH4NO3(aq)
(b) ammonium sulfide, mercury(II) nitrate, mercury(II) sulfide, ammonium nitrate
(c) precipitation reaction
3.71
(a) H2O, NH3, NH4+, and OH– (and a trace of H3O+); weak Brønsted base
(b) H2O, CH3CO2H, CH3CO2–, and H3O+ (and a trace of OH–); weak Brønsted acid
(c) H2O, Na+, and OH– (and a trace of H3O+); strong Brønsted base
(d) H2O, H3O+, and Br– (and a trace of OH–); strong Brønsted acid
58
Chapter 3
3.72
Chemical Reactions
Some possible answers:
(a) Water soluble:
Water insoluble:
(b) Water soluble:
Water insoluble:
3.73
CuCl2, copper(II) chloride
Cu(NO3)2, copper(II) nitrate
CuCO3, copper(II) carbonate
CuS, copper(II) sulfide
BaBr2, barium bromide
Ba(CH3CO2)2, barium acetate
BaSO4, barium sulfate
BaCrO4, barium chromate
(a) K2CO3(aq) + 2 HClO4(aq)  2 KClO4(aq) + CO2(g) + H2O()
gas-forming reaction
potassium carbonate, perchloric acid, potassium perchlorate, carbon dioxide, water
CO32–(aq) + 2 H3O+(aq)  CO2(g) + 3 H2O()
(b) FeCl2(aq) + (NH4)2S(aq)  FeS(s) + 2 NH4Cl(aq)
precipitation reaction
iron(II) chloride, ammonium sulfide, iron(II) sulfide, ammonium chloride
Fe2+(aq) + S2–(aq)  FeS(s)
(c) Fe(NO3)2(aq) + Na2CO3(aq)  FeCO3(s) + 2 NaNO3(aq)
precipitation reaction
iron(II) nitrate, sodium carbonate, iron(II) carbonate, sodium nitrate
Fe2+(aq) + CO32–(aq)  FeCO3(s)
(d) 3 NaOH(aq) + FeCl3(aq)  3 NaCl(aq) + Fe(OH)3(s)
precipitation reaction
sodium hydroxide, iron(III) chloride, sodium chloride, iron(III) hydroxide
Fe3+(aq) + 3 OH–(aq)  Fe(OH)3(s)
3.74
(a) Pb(NO3)2(aq) + 2 KOH(aq)  Pb(OH)2(s) + 2 KNO3(aq)
Pb2+(aq) + 2 OH–(aq)  Pb(OH)2(s)
(b) Cu(NO3)2(aq) + Na2CO3(aq)  CuCO3(s) + 2 NaNO3(aq)
Cu2+(aq) + CO32–(aq)  CuCO3(s)
3.75
(a) The NaOH will dissolve leaving Ca(OH)2 solid.
(b) The MgCl2 will dissolve leaving MgF2 solid.
(c) The KI will dissolve leaving AgI solid.
(d) The NH4Cl will dissolve leaving solid PbCl2 solid.
3.76
(a) CuCl2. The others are not soluble.
(b) HCl, H2SO4. The oxalic and phosphoric acids will dissolve but are weak electrolytes.
3.77
(a) NH3. NaOH and Ba(OH)2 are strong bases.
(b) CH3CO2H, HF. Na3PO4 and HNO3 are strong electrolytes.
59
Chapter 3
3.78
Chemical Reactions
(a) 2 HX(aq) + 2 OH-(aq)  2H2O() + 2 X-(aq)
(b) Zn(s) + 2 H3O+(aq)  Zn2+(aq) + H2(g) + H2O()
3.79
S2-(aq) + 2 H3O+(aq)  H2S(g) + 2 H2O()
2 H2O() + H2S(g) + Pb2+(aq)  PbS(s) + 2 H3O+(aq)
3.80
2HI(g)
H2(g) + I2(g)
The HI(g) decomposes to produce H2(g) and I2(g). The two product gases in turn react with each other to
produce the initial reactant HI(g). The amount of HI decreases until just enough H 2 and I2 is produced so
that they react to generate HI at exactly the same rate the HI decomposes. The concentrations of the three
gases then remain constant although the forward and reverse reactions keep proceeding, indicating that the
system has reached equilibrium.
3.81
(a) Reactants: Na (+1), I (–1), H (+1), S (+6), O (–2), Mn (+4)
Products: Na (+1), S (+6), O (–2), Mn (+2), I (0), H (+1)
(b) Oxidizing agent MnO2, NaI is oxidized; Reducing agent NaI, MnO2 is reduced
(c) product-favored
(d) sodium iodide, sulfuric acid, manganese(IV) oxide, sodium sulfate, manganese(II) sulfate, iodine, water
3.82
The first reaction is an oxidation-reduction reaction and a precipitation reaction. The second reaction is an
acid-base reaction and a gas-forming reaction.
Acids: H2S
Bases: NaHCO3
Oxidizing agents: Ag2S
3.83
Reducing agents: Al
One possibility:
MgCO3(s) + 2 HCl(aq)  MgCl2(aq) + H2O() + CO2(g)
Use a gas-forming reaction to form magnesium chloride, water, and carbon dioxide (which escapes from
solution as a gas). Heat the solution to evaporate the water, giving solid magnesium chloride.
3.84
One possibility:
3 Ba(OH)2(aq) + 2 H3PO4(aq)  Ba3(PO4)2(s) + 6 H2O()
Use an acid-base reaction to form barium phosphate and water. Filter the solid to isolate barium phosphate.
3.85
C6H12O6 is oxidized and is the reducing agent.
Ag+ is reduced and is the oxidizing agent.
3.86
Measure the conductivity of the solution and compare it that of a soluble compound such as sodium
chloride. A conductivity apparatus that uses a light bulb will glow only dimly if the solid dissolves to a
small extent but the light bulb will glow brightly if the solid dissolves to a great extent. The dissolving
process for PbCl2 is reactant-favored.
60
Chapter 3
3.87
Chemical Reactions
Measure the conductivity of the solution and compare it that of a strong acid such as hydrochloric acid. A
conductivity apparatus that uses a light bulb will glow only dimly if the acid is weak but the light bulb will
glow brightly if the acid is strong. The fact that lactic acid is an electrolyte indicates that the reaction
proceeds in the forward direction. To prove the reaction is reversible, add a large amount of strong acid (a
source of H3O+) to a solution containing lactic acid dissolved in water. The additional H 3O+ will cause the
reaction to proceed in the reverse direction, causing the precipitation of lactic acid.
3.88
One possible answer:
BaCO3(s) + 2 HCl(aq)  BaCl2(aq) + H2O() + CO2(g)
3.89
Possible answers:
(a) BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2 NaCl(aq)
(b) BaCO3(s) + H2SO4(aq)  BaSO4(s) + H2O() + CO2(g)
3.90
Possible answers:
(a) Zn(OH)2(s) + 2 HCl(aq)  ZnCl2(aq) + 2 H2O()
(b) ZnCO3(s) + 2 HCl(aq)  ZnCl2(aq) + H2O() + CO2(g)
(c) Zn(s) + Cl2(g)  ZnCl2(s)
3.91
(20.315 g Ni)(1 mol/58.6934 g Ni) = 0.34612 mol Ni
(33.258 g C)(1 mol/12.0107 g C) = 2.7690 mol C
(4.884 g H)(1 mol/1.00794 g H) = 4.846 mol H
(22.151 g O)(1 mol/15.9994 g O) = 1.3845 mol O
(19.392 g N)(1 mol/14.0067 g N) = 1.3845 mol N
2.7690 mol C/0.34612 mol Ni = 8.0001 C:Ni
4.846 mol H/0.34612 mol Ni = 14.00 H:N
1.3845 mol O/0.34612 mol Ni = 4.0001 O:Ni
1.3845 mol N/0.34612 mol Ni = 4.0001 N:Ni
NiC8H14N4O4
3.92
(73.945 g Tb)(1 mol/158.925 g Tb) = 0.46528 mol Tb
100.000 g TbxOy – 73.945 g Tb = 26.055 g O
(26.055 g O)(1 mol/15.9994 mol O) = 1.6285 mol O
1.6285 mol O/0.46528 mol Tb = 3.5000 O:Tb
Tb2O7
61
Chapter 3
Chemical Reactions
4 Tb(S) + 7 O2(g)  2 Tb2O7(s)
3.93
(a) As2S3: As +3, S -2; HNO3: N +5; H3AsO4: As +5; NO: N +2; S: S 0
(b) (16.199 g As)(1 mol/74.9216 g As) = 0.21621 mol As
(69.964 g Ag)(1 mol/107.868 g Ag) = 0.64861 mol Ag
100.000 g AgxAsOy – 16.199 g As – 69.964 g Ag = 13.837 g O
(13.837 g O)(1 mol/15.9994 g O) = 0.86484 mol O
0.64861 mol Ag/0.21621 mol As = 2.9999 Ag:As
0.86484 mol O/0.21621 mol As = 4.0000 O:As
Ag3AsO4
SOLUTIONS TO APPLYING CHEMICAL PRINCIPLES: SUPERCONDUCTORS
1.
mol La = 63.43 g La ·
1 mol
= 0.4566 mol La
138.91 g
mol Ba = 5.085 g Ba ·
1 mol
= 0.03703 mol Ba
137.33 g
mol Cu = 15.69 g Cu ·
1 mol
= 0.2469 mol Cu
63.546 g
mol O = 15.80 g O ·
1 mol
= 0.9876 mol O
15.9994 g
Divide each value above by the moles of Cu since one mole of the superconductor contains one mole of
Cu:
0.4566/0.2469 = 1.849 mol La
0.03703/0.2469 = 0.1500 mol Ba
0.2469/0.2469 = 1.000 mol Cu
0.9876/0.2469 = 4.000 mol O
Formula = La2-xBaxCuO4 = La1.85Ba0.15CuO4
x = 0.15
2.
62
YBa2Cu3O6.93
Chapter 3
Chemical Reactions
88.91 g
1.000 mol Y ·
= 88.91 g Y
1 mol
2.000 mol Ba ·
137.33 g
= 274.7 g Ba
1 mol
3.000 mol Cu ·
63.546 g
= 190.6 g Cu
1 mol
6.930 mol O ·
15.9994 g
= 110.9 g O
1 mol
Total mass = 665.1 g
%Y = 88.91 g/665.1g · 100 % = 13.37 %
% Ba = 274.7 g/665.1 g · 100 % = 41.30 %
%Cu = 190.6 g/665.1 g · 100% = 28.66 %
%O = 110.9 g.665.1 g · 100 % = 16.67 %
3.
One formula unit of the compound contains two Cu2+ ions and one Cu3+ ion.
4.
Y2O3(s) + 4 BaCO3(s) + 6 CuO(s)  2 YBa2Cu3O6.5(s) + 4 CO2(g)
The compound is YBa2Cu3O7-x where x = 0.5.
5.
1.00 g YBCO(1 mol YBCO/658.20 g)(0.43 mol O/1 mol YBCO)(1 mol O 2/2 mol O)(32.00 g O2/mol O2) =
0.011 g O2
63