1. Determine the oxidation number of each element present in the following substances.
a. BaH2
d. LiNO2
g. AsO33b. Al4C3
e. (NH4)2C2O4
h. XeO3Fc. KCN
f. S8
2. Examples of molecules and ions of vanadium and oxygen are listed below. In this list, identify molecules
and ions in which the oxidation number of vanadium is the same.
a. V2O5
d. VO5
g. VO3b. V2O3
e. VO2+
h. VO43c. VO2
f. VO2+
i. V3O933. Identify a polyatomic ion in which chlorine has an oxidation number of +3.
4. We stated at the beginning of this unit that oxidation originally meant “to combine with oxygen”. Explain why a
metal that combines with the element oxygen undergoes oxidation as we now define it. What happens to the
oxygen in this reaction? Write a balanced chemical equation for a reaction that illustrates your answer.
5. Write two different definitions for a redox reaction.
6. a) On which side of an oxidation half-reaction are the electrons? Why?
b) On which side of a reduction half-reaction are the electrons?
7. Explain why, in a redox reaction, the reducing agent undergoes oxidation.
8. Explain why, in a redox reaction, the oxidizing agent undergoes reduction.
9. In a combination reaction, does metallic lithium act as an oxidizing agent or a reducing agent? Explain. (Hint: what is
the oxidation number/ charge of lithium in compounds.)
10. 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?
11. Explain why you would not expect sulphide ions to act as an oxidazing agent.
12. When an element combines with another element, is the reaction a redox reaction? Explain your answer.
13. When a metallic element reacts with a non-metallic element, which reactant is
a. Oxidized?
c. The oxidizing agent?
b. Reduced?
d. The reducing agent?
14. The element potassium is made industrially by the single displacement reaction of molten sodium with molten
potassium chloride.
a) Write a net ionic equation for the reaction, assuming that all reactants and products are in the liquid state.
b) Identify the oxidizing and the reducing agent in the reaction.
c) Explain why the reaction is carried out in the liquid state and not in aqueous solution. (Hint: Find out how alkali
metals react with water.)
15. Nickel and copper are two metals that have played a role in the economy of Newfoundland and Labrador. Nickel and
copper ores usually contain the metals as sulfides, such as NiS and Cu2S. Do the extractions of these pure elemental
metals from their ores involve redox reactions? Explain your reasoning.
16. Determine whether each of the following reactions is a redox reaction?
a. H2 + I2 → 2HI
b. 2NaHCO3 → 2Na2CO3 + H2O + CO2
c. 2 HBr + Ca(OH)2 → CaBr2 + 2 H2O
d. PCl5 → PCl3 + Cl2
e. 2C6H6 + 15 O2 → 12 CO2 + 6 H2O
f. CaO + SO2 → CaSO3
g. KMnO4 + 5 CuCl + 8 HCl → KCl + MnCl2 + 5 CuCl2 + 4 H2O
h. 2 NH4 VO3 → V2O5 + 2 NH3 + H2O
17. Write balanced half-reactions from the net ionic equation for the reaction between solid aluminum and
aqueous iron(III) sulfate. The sulphate ions are spectator ions, and are not included.
Al(s) + Fe3+(aq) → Al3+(aq) + Fe (s)
18. Write balanced half-reactions from the following net ionic equations.
a) Fe (s) + Cu 2+(aq) → Fe2+(aq) + Cu (s)
b) Cd (s) + 2Ag +(aq) → Cd2+(aq) + 2Ag (s)
19. Write balance net ionic equations for the following reactions.
a) Sn (s) + PbCl2 (aq) → SnCl2 (aq) ) + Pb (s)
b) Au(NO3)3 (aq) + 3 Ag (s) → 3 AgNO3 (aq) + Au (s)
c) Fe2(SO4)3 (aq) + 3 Zn (s) → 3 ZnSO4 (aq) + 2Fe (s)
20. Write balanced half-reactions for the following equations.
a) Sn (s) + PbCl2 (aq) → SnCl2 (aq) ) + Pb (s)
b) Au(NO3)3 (aq) + 3 Ag (s) → 3 AgNO3 (aq) + Au (s)
c) Fe2(SO4)3 (aq) + 3 Zn (s) → 3 ZnSO4 (aq) + 2Fe (s)
21. Determine whether each reaction is a redox reaction. If so, identify the oxidizing agent and the reducing
agent.
a) H2O2 + 2 Fe(OH)2 → 2 Fe(OH)3
b) PCl3 + 3 H2O → H3PO3 + 3HCl
22. For the following balanced net ionic equation, identify the reactant that undergoes oxidation and the
reactant that undergoes reduction.
a) Br2 + 2ClO2- → 2Br - + 2 ClO2
b) 2 Ag+ (aq) + Cu (s) → 2 Ag (s) + Cu 2+ (aq)
c) Pb2+ (aq) + S2-(aq) → PbS (s)
d) 2Mn2+(aq) + 5BiO3- + 14 H+ → 2 MnO4- + 5 Bi3+ + 7 H2O
23. The method used to manufacture nitric acid involves the following three steps.
Step 1:
4 NH3 + 5 O2 → 4NO + 6 H2O
Step 2:
2 NO + O2 → 2 NO2
Step 3:
3 NO2 + H2O → 2 HNO3 + NO
a) Which of these steps are redox reactions?
b) Identify the oxidizing agent and the reducing agent in each redox reaction.
24. Balance following half-reactions by adding appropriate number of electrons. State if it is an oxidation half
reaction or a reduction half reaction.
a) Na → Na+
f) NO2 + H2O → NO3- + 2H+
b) S2→ S+6
g) ClO4- + 8H+ → Cl- + 4H2O
c) 2Cl- → Cl2
h) MnO4- + 8H+ → Mn2+ + 4H2O
d) Cr2+ → Cr3+
i) Al + 4 OH→ Al(OH)4 e) O22- → 2O-2
25. Write a net ionic equation for a reaction in which:
a) Fe2+ acts as an oxidizing agent
b) Al acts as a reducing agent
c) Au3+ acts as an oxidizing agent
d) Cu acts as a reducing agent
e) Sn2+ acts as an oxidizing agent and as a reducing agent