In the Brønsted–Lowry concept of acids and bases, acid– base reactions are viewed as proton-transfer reactions. The stronger the acid, the weaker is its conjugate base. If we were to think of redox reactions in a similar way, what particle would be analogous to the proton? Would strong oxidizing agents be analogous to strong acids or strong bases? [Sections 20.1 and 20.2]

The diagram that follows represents a molecular view of a process occurring at an electrode in a voltaic cell.

(a) Does the process represent oxidation or reduction?

The diagram that follows represents a molecular view of a process occurring at an electrode in a voltaic cell.

(b) Is the electrode the anode or cathode?

The diagram that follows represents a molecular view of a process occurring at an electrode in a voltaic cell.

(c) Why are the atoms in the electrode represented by larger spheres than those in the solution? [Section 20.3]

Assume that you want to construct a voltaic cell that uses the following half-reactions: A2+1aq2 + 2 e- ¡ A1s2 Ered ° = -0.10 V B2+1aq2 + 2 e- ¡ B1s2 E°red = -1.10 V You begin with the incomplete cell pictured here in which the electrodes are immersed in water.

(a) What additions must you make to the cell for it to generate a standard emf?

Consider the following table of standard electrode potentials for a series of hypothetical reactions in an aqueous solution: reduction half-reaction E °(V) (c) Which substance(s) can oxidize C2+?

Consider the following voltaic cell:

(c) What is the change in the cell voltage when the ion concentrations in the cathode half-cell are increased by a factor of 10?

The electrodes in a silver oxide battery are silver oxide 1Ag2O2 and zinc (b) Which battery do you think has an energy density most similar to the silver oxide battery: a Li-ion battery, a nickel– cadmium battery, or a lead–acid battery? [Section 20.7]

Bars of iron are put into each of the three beakers as shown here. In which beaker—A, B, or C—would you expect the iron to show the most corrosion ? [Section 20.8]

Magnesium, the element, is produced commercially by electrolysis from a molten salt (the 'electrolyte') using a cell similar to the one shown here. (a) What is the most common oxidation number for Mg when it is part of a salt?

Magnesium, the element, is produced commercially by electrolysis from a molten salt (the 'electrolyte') using a cell similar to the one shown here. (b) Chlorine gas is evolved as voltage is applied in the cell. Knowing this, identify the electrolyte.

Magnesium, the element, is produced commercially by electrolysis from a molten salt (the 'electrolyte') using a cell similar to the one shown here. (c) Recall that in an electrolytic cell the anode is given the + sign and the cathode is given the – sign, which is the opposite of what we see in batteries. What half-reaction occurs at the anode in this electrolytic cell?

(b) On which side of an oxidation half-reaction do the electrons appear?

Indicate whether each of the following statements is true or false: (a) If something is oxidized, it is formally losing electrons.

Indicate whether each of the following statements is true or false: (b) For the reaction Fe³⁺(aq) + Co²⁺(aq) → Fe2⁺(aq) + Co³⁺(aq), Fe³⁺(aq) is the reducing agent and Co²⁺(aq) is the oxidizing agent.

Indicate whether each of the following statements is true or false: (c) If there are no changes in the oxidation state of the reactants or products of a particular reaction, that reaction is not a redox reaction.

Indicate whether each of the following statements is true or false: (a) If something is reduced, it is formally losing electrons. (b) A reducing agent gets oxidized as it reacts.

Indicate whether each of the following statements is true or false: (c) An oxidizing agent is needed to convert CO into CO₂.

For each of the following balanced oxidation–reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction. (a) I₂O₅(s) + 5 CO(g) → I₂(s) + 5 CO₂(g) (b) 2 Hg²⁺(aq) + N₂H₄(aq) → 2 Hg(l) + N₂(g) + 4 H⁺(aq) (c) 3 H₂S(aq) + 2 H⁺(aq) + 2 NO₃^-(aq) → 3 S(s) + 2 NO(g) + 4 H₂O(l)

For each of the following balanced oxidation–reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction. (a) 2 MnO₄^-(aq) + 3 S^2-(aq + 4 H₂O(l) → 3 S(s) + 2 MnO₂(s) + 8 OH^-(aq) (b) 4 H₂O₂(aq) + Cl₂O₇(g) + 2 OH^-(aq) → 2 ClO₂^-(aq) + 5 H₂O(l) + 4 O₂(g) (c) Ba²⁺(aq) + 2 OH^-(aq) + H₂O₂(aq) + 2 ClO₂(aq) → Ba(ClO₂)₂(s) + 2 H₂O(l) + O₂(g)

Indicate whether the following balanced equations involve oxidation–reduction. If they do, identify the elements that undergo changes in oxidation number. (a) PBr₃(l) + 3 H₂O(l) → H₃PO₃(aq) + 3 HBr(aq)

Indicate whether the following balanced equations involve oxidation–reduction. If they do, identify the elements that undergo changes in oxidation number. (b) NaI(aq) + 3 HOCl(aq) → NaIO₃(aq) + 3 HCl(aq) (c) 3 SO(1g) + 2 HNO₃(aq) + 2 H₂O(l) → 3 H₂SO₄(aq) + 2 NO(g)

Indicate whether the following balanced equations involve oxidation–reduction. If they do, identify the elements that undergo changes in oxidation number. (a) 2 AgNO₃(aq) + CoCl₂(aq) → 2 AgCl(s) + Co(NO₃)₂(aq)