Review
Reflecting on Chapter 10
Summarize this chapter in the format of your
choice. Here are a few ideas to use as guidelines:
• Give two different definitions for the term
oxidation.
• Give two different definitions for the term
reduction.
• Define a half-reaction. Give an example of
an oxidation half-reaction and a reduction
half-reaction.
• Compare the half-reaction and oxidation number
methods of balancing equations.
• Practise balancing equations using both methods.
• Write an example of a balanced chemical
equation for a redox reaction. Assign oxidation
numbers to each element in the equation, then
explain how you know it is a redox reaction.
Reviewing Key Terms
For each of the following terms, write a sentence
that shows your understanding of its meaning.
ore
oxidation
reduction
oxidation-reduction
redox reaction
reaction
oxidizing agent
reducing agent
half-reaction
disproportionation
oxidation numbers
Knowledge/Understanding
1. For each reaction below, write a balanced
chemical equation by inspection.
(a) zinc metal with aqueous silver nitrate
(b) aqueous cobalt(II) bromide with aluminum
metal
(c) metallic cadmium with aqueous tin(II)
chloride
2. For each reaction in question 1, write the total
ionic and net ionic equations.
3. For each reaction in question 1, identify the
oxidizing agent and reducing agent.
4. For each reaction in question 1, write the two
half-reactions.
5. When a metallic element reacts with a
non-metallic element, which reactant is
oxidized?
reduced?
the oxidizing agent?
the reducing agent?
(a)
(b)
(c)
(d)
6. Use a Lewis structure to assign an oxidation
number to each element in the following
compounds.
(a) BaCl2
(b) CS2
(c) XeF4
7. Determine the oxidation number of each
element present in the following substances.
BaH2
Al4C3
KCN
LiNO2
(NH4)2C2O4
S8
AsO33−
VO2+
XeO3F−
S4O62−
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
8. Identify a polyatomic ion in which chlorine has
an oxidation number of +3.
9. Determine which of the following balanced
chemical equations represent redox reactions.
For each redox reaction, identify the oxidizing
agent and the reducing agent.
(a) 2C6H6 + 15O2 → 12CO2 + 6H2O
(b) CaO + SO2 → CaSO3
(c) H2 + I2 → 2HI
(d) KMnO4 + 5CuCl + 8HCl →
KCl + MnCl2 + 5CuCl2 + 4H2O
10. Determine which of the following balanced net
ionic equations represent redox reactions. For
each redox reaction, identify the reactant that
undergoes oxidation and the reactant that
undergoes reduction.
(a) 2Ag+(aq) + Cu(s) → 2Ag(s) + Cu2+(aq)
(b) Pb2+(aq) + S2−(aq) → PbS(s)
(c) 2Mn2+ + 5BiO3− + 14H+ →
2MnO4− + 5Bi3+ + 7H2O
Chapter 10 Oxidation-Reduction Reactions • MHR
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11. (a) Examples of molecules and ions composed
only of vanadium and oxygen are listed
below. In this list, identify molecules and
ions in which the oxidation number of
vanadium is the same.
V2O5
V2O3
VO2
VO
VO2+
VO2+
VO3−
VO43−
V3O93−
(b) Is the following reaction a redox reaction?
2NH4VO3 → V2O5 + 2NH3 + H2O
12. The method used to manufacture nitric acid
involves the following three steps.
Step 1 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)
Step 2 2NO(g) + O2(g) → 2NO2(g)
Step 3 3NO2(g) + H2O() → 2HNO3(aq) + NO(g)
(a) Which of these steps are redox reactions?
(b) Identify the oxidizing agent and the reducing
agent in each redox reaction.
13. In a synthesis reaction involving elements A
and B, the oxidation number of element A
increases. What happens to the oxidation
number of element B? How do you know?
14. Balance each of the following half-reactions.
(a) The reduction of iodine to iodide ions
(b) The oxidation of lead to lead(IV) ions
(c) The reduction of tetrachlorogold(III) ions,
AuCl4− ,
16. Use the oxidation number method to balance
each of the following equations.
(a) SiCl4 + Al → Si + AlCl3
(b) PH3 + O2 → P4O10 + H2O
(c) I2O5 + CO → I2 + CO2
(d) SO32− + O2 → SO42−
17. Complete and balance a net ionic equation
for each of the following disproportionation
reactions.
(a) NO2 → NO2− + NO3− (acidic conditions),
which is one of the reactions involved in
acid rain formation
(b) Cl2 → ClO− + Cl− (basic conditions), which
is one of the reactions involved in the
bleaching action of chlorine in basic solution
18. Balance each of the following net ionic
equations. Then include the named spectator
ions to write a balanced chemical equation.
Include the states.
(a) Co3+ + Cd → Co2+ + Cd2+ (spectator ions
NO3− )
(b) Ag+ + SO2 → Ag + SO42− (acidic conditions;
spectator ions NO3− )
(c) Al + CrO42− → Al(OH)3 + Cr(OH)3 (basic
conditions; spectator ions Na+ )
19. If possible, give an example for each.
(a) a synthesis reaction that is a redox reaction
(b) a synthesis reaction that is not a redox
(c)
(d)
to chloride ions and metallic gold
(d) C2H5OH → CH3COOH (acidic conditions)
(e) S8 → H2S (acidic conditions)
(f) AsO2− → AsO43− (basic conditions)
15. Use the half-reaction method to balance each of
the following equations.
(a) MnO2 + Cl− → Mn2+ + Cl2 (acidic
conditions)
(b) NO + Sn → NH2OH + Sn2+ (acidic
conditions)
(c) Cd2+ + V2+ → Cd + VO3− (acidic
conditions)
(d) Cr → Cr(OH)4− + H2 (basic conditions)
(e) S2O32− + NiO2 → Ni(OH)2 + SO32− (basic
conditions)
(f) Sn2+ + O2 → Sn4+ (basic conditions)
500 MHR • Unit 5 Electrochemistry
(e)
(f)
reaction
a decomposition reaction that is a redox
reaction
a decomposition reaction that is not a redox
reaction
a double displacement reaction that is a
redox reaction
a double displacement reaction that is not a
redox reaction
20. Give an example of a reaction in which sulfur
behaves as
(a) an oxidizing agent
(b) a reducing agent
21. Write a balanced equation for a synthesis
reaction in which elemental oxygen acts as
a reducing agent.
22. Phosphorus, P4(s), reacts with hot water to
form phosphine, PH3(g), and phosphoric acid.
(a) Write a balanced chemical equation for this
reaction.
(b) Is the phosphorus oxidized or reduced?
Explain your answer.
23. The thermite reaction, which is highly
exothermic, can be used to weld metals. In
the thermite reaction, aluminum reacts with
iron(III) oxide to form iron and aluminum
oxide. The temperature becomes so high that
the iron is formed as a liquid.
(a) Write a balanced chemical equation for the
reaction.
(b) Is the reaction a redox reaction? If so, identify
the oxidizing agent and the reducing agent.
27. Highly toxic phosphine gas, PH3 , is used in
industry to produce flame retardants. One
way to make phosphine on a large scale is
by heating elemental phosphorus with a
strong base.
(a) Balance the following net ionic equation
for the reaction under basic conditions.
P4 → H2PO2− + PH3
(b) Show that the reaction in part (a) is a
disproportionation reaction.
(c) Calculate the mass of phosphine that can
theoretically be made from 10.0 kg of
phosphorus by this method.
Communication
28. Explain why, in a redox reaction, the reducing
agent undergoes oxidation.
Inquiry
29. Explain why you would not expect sulfide ions
24. Iodine reacts with concentrated nitric acid
to form iodic acid, gaseous nitrogen dioxide,
and water.
(a) Write the balanced chemical equation.
(b) Calculate the mass of iodine needed to
produce 28.0 L of nitrogen dioxide at STP.
25. Describe a laboratory investigation you could
perform to decide whether tin or nickel is the
better reducing agent. Include in your description all the materials and equipment you would
need, and the procedure you would follow.
26. The following table shows the average
composition, by volume, of the air we inhale
and exhale, as part of a biochemical process
called respiration. (The values are rounded.)
Gas
Oxygen
Carbon dioxide
Nitrogen and other gases
Inhaled Air
(% by volume)
Exhaled Air
(% by volume)
21
16
0.04
4
79
80
How do the data indicate that at least one
redox reaction is involved in respiration?
to act as an oxidizing agent.
30. Why can’t the oxidation number of an element
in a compound be greater than the number of
valence electrons in one atom of that element?
31. Explain why the historical use of the word
“reduction,” that is, the production of a metal
from its ore, is consistent with the modern
definitions of reduction.
32. Organic chemists sometimes describe redox
reactions in terms of the loss or gain of pairs of
hydrogen atoms. Examples include the addition
of hydrogen to ethene to form ethane, and the
elimination of hydrogen from ethanol to form
ethanal.
(a) Write a balanced equation for each reaction.
(b) Determine whether the organic reactant is
oxidized or reduced in each reaction.
(c) Write a definition of oxidation and a
definition of reduction based on an organic
reactant losing or gaining hydrogen.
(d) Would your definitions be valid for the
synthesis and decomposition of a metal
hydride? Explain your answer.
Chapter 10 Oxidation-Reduction Reactions • MHR
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Making Connections
33. The compound NaAl(OH)2CO3 is a component
of some common stomach acid remedies.
(a) Determine the oxidation number of each
element in the compound.
(b) Predict the products of the reaction of the
compound with stomach acid (hydrochloric
acid), and write a balanced chemical equation for the reaction.
(c) Were the oxidation numbers from part (a)
useful in part (b)? Explain your answer.
(d) What type of reaction is this?
(e) Check your medicine cabinet at home for
stomach acid remedies. If possible, identify
the active ingredient in each remedy.
34. Two of the substances on the head of a safety
match are potassium chlorate and sulfur. When
the match is struck, the potassium chlorate
decomposes to give potassium chloride and
oxygen. The sulfur then burns in the oxygen
and ignites the wood of the match.
(a) Write balanced chemical equations for the
decomposition of potassium chlorate and for
the burning of sulfur in oxygen.
(b) Identify the oxidizing agent and the reducing
agent in each reaction in part (a).
(c) Does any element in potassium chlorate
undergo disproportionation in the reaction?
Explain your answer.
(d) Research the history of the safety match to
determine when it was invented, why it was
invented, and what it replaced.
35. Ammonium ions, from fertilizers or animal
waste, are oxidized by atmospheric oxygen.
The reaction results in the acidification of soil
on farms and the pollution of ground water
with nitrate ions.
(a) Write a balanced net ionic equation for this
reaction.
(b) Why do farmers use fertilizers? What
alternative farming methods have you heard
of? Which farming method(s) do you
support, and why?
502 MHR • Unit 5 Electrochemistry
36. One of the most important discoveries in the
history of the chemical industry in Ontario was
accidental. Thomas “Carbide” Willson
(1860 – 1915) was trying to make the element
calcium from lime, CaO, by heating the lime
with coal tar. Instead, he made the compound
calcium carbide, CaC2 . This compound reacts
with water to form a precipitate of calcium
hydroxide and gaseous ethyne (acetylene).
Willson’s discovery led to the large-scale use
of ethyne in numerous applications.
(a) Was Willson trying to perform a redox
reaction? How do you know? Why do you
not need to know the substances in coal tar
to answer this question?
(b) Write a balanced chemical equation for the
reaction of calcium carbide with water. Is
this reaction a redox reaction?
(c) An early use of Willson’s discovery was in
car headlights. Inside a headlight, the
reaction of calcium carbide and water
produced ethyne, which was burned to
produce light and heat. Write a balanced
chemical equation for the complete combustion of ethyne. Is this reaction a redox
reaction?
(d) Research the impact of Willson’s discovery
on society, from his lifetime to the present
day.
Answers to Practice Problems and Short Answers to
Section Review Questions
Practice Problems: 1. Zn(s) + Fe2+(aq) → Zn2+(aq) + Fe(s)
2.(a) 3Mg(s) + 2Al3+(aq) → 3Mg2+(aq) + 2Al(s)
(b) 2Ag+(aq) + Cd(s) → 2Ag(s) + Cd2+(aq)
3.(a) Mg oxidized, Al3+ reduced
(b) Cd oxidized, Ag+ reduced
4.(a) Al3+ oxidizing agent, Mg reducing agent
(b) Ag+ oxidizing agent, Cd reducing agent
5. Al(s) → Al3+(aq) + 3e− , Fe3+(aq) + 3e− → Fe(s)
6.(a) Fe(s) → Fe2+(aq) + 2e− , Cu2+(aq) + 2e− → Cu(s)
(b) Cd(s) → Cd2+(aq) + 2e− ,
−
Ag+(aq) + 1e− → Ag(s)
+ 2e , Pb2+(aq) + 2e− → Pb(s)
(b) Ag(s) → Ag + e , Au3+(aq) + 3e− → Au(s)
(c) Zn(s) → Zn2+(aq) + 2e− , Fe3+(aq) + 3e− → Fe(s)
8. Hg22+(aq) → Hg() + Hg2+(aq) , Hg22+(aq) + 2e− → 2Hg() ,
Hg22+(aq) → 2Hg2+(aq) + 2e−
9.(a) +3 (b) 0 (c) +6 (d) +5 (e) 0 (f) +2
7.(a) Sn(s) →
Sn2+(aq)
+
−
+1 ; S, +4 ; O, −2
+1 ; O, −2 (c) H, +1 ; P, +5 ; O, −2
11.(a) +2 (b) −1 12.(a) Al, +3 ; H, +1 ; C, +4 ; O, −2
(b) N, −3 ; H, +1 ; P, +5 ; O, −2 (c) K, +1 ; H, +1 ; I, +7 ; O, −2
10.(a) H,
(b) H,
13.(a) yes (b) no
H2O2 oxidizing agent, Fe2+ reducing agent
15. ClO2− oxidized, Br2 reduced 16. yes
17. Ce4+ + e− → Ce3+ 18. 2Br− → Br2 + 2e−
19.(a) O2 + 2H+ + 2e− → H2O2 (b) 2H2O → O2 + 4H+ + 4e−
(c) 2NO3− + 12H+ + 10e− → N2 + 6H2O
20.(a) ClO3− + 6H+ + 6e− → Cl− + 3H2O
(b) NO + 2H2O → NO3− + 4H+ + 3e−
(c) Cr2O72−− + 14H+ + 6e− → 2Cr3+ + 7H2O
21. Cr2+ → Cr3+ + e− 22. O2 + 4e− → 2O2−
23.(a) Al + 4OH− → Al(OH)4− + 3e−
(b) CN− + 2OH− → CNO− + H2O + 2e−
(c) MnO4− + 2H2O + 3e− → MnO2 + 4OH−
(d) CrO42− + 4H2O + 3e− → Cr(OH)3 + 5OH−
(e) 2CO32− + 2H2O + 2e− → C2O42− + 4OH−
24.(a) FeO42− + 8H+ + 3e− → Fe3+ + 4H2O
(b) ClO2− + 2H2O + 4e− → Cl− + 4OH−
25.(a) 2Na + F2 → 2NaF
ox: Na → Na+ + e−
red: F2 + 2e− → 2F−
(b) 3Mg + N2 → Mg3N2
ox: Mg → Mg2+ + 2e−
red: N2 + 6e− → 2N3−
(c) 2HgO → 2Hg + O2
ox: 2O2− → O2 + 4e−
red: Hg2+ + 2e− → Hg
26. 2Cu2+ + 5I− → 2CuI + I3−
27.(a) MnO4− + 5Ag + 8H+ → Mn2+ + 5Ag+ + 4H2O
oxidizing agent, MnO4− ; reducing agent, Ag
(b) Hg + 2NO3− + 4Cl− + 4H+ → HgCl42− + 2NO2 + 2H2O
oxidizing agent, NO3− ; reducing agent, Hg
(c) AsH3 + 4Zn2+ + 4H2O → H3AsO4 + 4Zn + 8H+
oxidizing agent, Zn2+ ; reducing agent, AsH3
(d) 3I3− + 3H2O → 8I− + IO3− + 6H+
oxidizing agent, I3− ; reducing agent, I3−
28.(a) 3CN− + 2MnO4− + H2O → 3CNO− + 2MnO2 + 2OH−
oxidizing agent, MnO4− ; reducing agent, CN−
(b) H2O2 + 2ClO2 + 2OH− → 2ClO2− + O2 + 2H2O
oxidizing agent, ClO2 ; reducing agent, H2O2
(c) 6ClO− + 2CrO2− + 2H2O → 3Cl2 + 2CrO42− + 4OH−
oxidizing agent, ClO− ; reducing agent, CrO2−
(d) 2Al + NO2− + H2O + OH− → NH3 + 2AlO2−
oxidizing agent, NO2− ; reducing agent, Al
14.(a)
CS2 + 3O2 → CO2 + 2SO2
B2O3 + 6Mg → 3MgO + Mg3B2
(b) 8H2S + 8H2O2 → S8 + 16H2O
31.(a) Cr2O72− + 6Fe2+ + 14H+ → 2Cr3+ + 6Fe3+ + 7H2O
(b) I2 + 10NO3− + 8H+ → 2IO3− + 10NO2 + 4H2O
(c) 2PbSO4 + 2H2O → Pb + PbO2 + 2SO42− + 4H+
32.(a) 3Cl− + 2CrO42− + H2O → 3ClO− + 2CrO2− + 2OH−
(b) 3Ni + 2MnO4− + H2O → 3NiO + 2MnO2 + 2OH−
(c) I− + 6Ce4+ + 6OH− → IO3− + 6Ce3+ + 3H2O
Section Review 10.1: 1.(a) yes;
2AgNO3(aq) + Cd(s) → 2Ag(s) + Cd(NO3)2(aq) ;
2Ag+(aq) + Cd(s) → 2Ag(s) + Cd2+(aq) ;
Ag+(aq) + e− → Ag(s) ; Cd(s) → Cd2+(aq) + 2e−
(b) no (c) yes; 2Al(s) + 3HgCl2(aq) → 2AlCl3(aq) + 3Hg() ;
2Al(s) + 3Hg2+(aq) → 2Al3+(aq) + 3Hg() ;
Al(s) → Al3+(aq) + 3e− ; Hg2+(aq) + 2e− → Hg()
29.
30.(a)
2.(a) right side (b) left side 4. reducing agent
K+ + Na() → Na+ + K()
(b) oxidizing agent, K+ ; reducing agent, Na()
10.2: 2.(a) yes (b) no (c) no (d) yes
6.(a) +2 (b) +5 and −1
7.(a) N2 + 3H2 → 2NH3 (0, 0, −3 , +1 ); N2 is oxidizing agent,
H2 is reducing agent (b) no
8.(a) 2C + O2 → 2CO , Fe2O3 + 3CO → 2Fe + 3CO2
6.(a)
(b) carbon—reducing agent, oxygen—oxidizing agent;
carbon monoxide—reducing agent,
iron(III) oxide—oxidizing agent
10.3: 1.(a) 2C + 2e− → C22− , reduction
(b) 2S2O32− → S4O62− + 2e− , oxidation
(c) 4AsO43− + 20H+ + 8e− → As4O6 + 10H2O, reduction
(d) Br2 + 12OH− → 2BrO3− + 6H2O + 10e− , oxidation
3Co3+ + Au → 3Co2+ + Au3+
(b) 3Cu + 2NO3− + 8H+ → 3Cu2+ + 2NO + 4H2O
(c) 3NO3− + 8Al + 5OH− + 2H2O → 3NH3 + 8AlO2−
3. 2Ag+ + HCHO + H2O → 2Ag + HCOOH + 2H+
4.(b) 2NH3 + 8OH− → N2O + 7H2O + 8e−
5.(a) 2N2H4 + N2O4 → 3N2 + 4H2O
(b) N2O4 is oxidizing agent, N2H4 is reducing agent
(c) 2NH3 + ClO− → N2H4 + Cl− + H2O
2.(a)
6.(a) Larissa was right
3Zn + SO42− + 8H+ → 3Zn2+ + S + 4H2O
10.4: 1. no, not a redox reaction
2.(a) CH3COOH + 2O2 → 2CO2 + 2H2O
(b) O2 + 2H2SO3 → 2HSO4− + 2H+
3.(a) 8NH3 + 3Cl2 → 6NH4Cl + N2
(b) 3Mn3O4 + 8Al → 4Al2O3 + 9Mn
(b)
Chapter 10 Oxidation-Reduction Reactions • MHR
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