Exercises

For the scenario in the previous question, how many electrons moved through the circuit?

For each of the following unbalanced half-reactions, determine whether an oxidation or reduction is occurring.

(a) ${\text{Cl}}^{\text{−}}⟶{\text{Cl}}_2$

(b) ${\text{Mn}}^{\mathrm{2+}}⟶{\text{MnO}}_2$

(c) ${\text{H}}_2⟶{\text{H}}^{\text{+}}$

(d) ${\text{NO}}_3{}^{-}⟶\text{NO}$

Balance the following in acidic solution:

(a) ${\text{H}}_2{\text{O}}_2+{\text{Sn}}^{\mathrm{2+}}⟶{\text{H}}_2\text{O}+{\text{Sn}}^{\mathrm{4+}}$

(b) ${\text{PbO}}_2+\text{Hg}⟶{\text{Hg}}_2{}^{\mathrm{2+}}+{\text{Pb}}^{\mathrm{2+}}$

(c) $\text{Al}+{\text{Cr}}_2{\text{O}}_7{}^{\mathrm{2-}}⟶{\text{Al}}^{\mathrm{3+}}+{\text{Cr}}^{\mathrm{3+}}$

Balance the following in basic solution:

(a) ${\text{SO}}_3{}^{\mathrm{2-}}(aq)+{\text{Cu(OH)}}_2(s)⟶{\text{SO}}_4{}^{\mathrm{2-}}(aq)+\text{Cu(OH)}(s)$

(b) ${\text{O}}_2(g)+{\text{Mn(OH)}}_2(s)⟶{\text{MnO}}_2(s)$

(c) ${\text{NO}}_3{}^{\text{−}}(aq)+{\text{H}}_2(g)⟶\text{NO}(g)$

(d) $\text{Al}(s)+{\text{CrO}}_4{}^{\mathrm{2-}}(aq)⟶{\text{Al(OH)}}_3(s)+{\text{Cr(OH)}}_4{}^{\text{−}}(aq)$

Why is it not possible for hydroxide ion (OH⁻) to appear in either of the half-reactions or the overall equation when balancing oxidation-reduction reactions in acidic solution?

Why must the charge balance in oxidation-reduction reactions?

Given the following cell notations, determine the species oxidized, species reduced, and the oxidizing agent and reducing agent, without writing the balanced reactions.

(a) $\text{Mg}(s)│{\text{Mg}}^{\mathrm{2+}}(aq)║{\text{Cu}}^{\mathrm{2+}}(aq)│\text{Cu}(s)$

(b) $\text{Ni}(s)│{\text{Ni}}^{\mathrm{2+}}(aq)║{\text{Ag}}^{\text{+}}(aq)│\text{Ag}(s)$

Balance the following reactions and write the reactions using cell notation. Ignore any inert electrodes, as they are never part of the half-reactions.

(a) $\text{Al}(s)+{\text{Zr}}^{\mathrm{4+}}(aq)⟶{\text{Al}}^{\mathrm{3+}}(aq)+\text{Zr}(s)$

(b) ${\text{Ag}}^{\text{+}}(aq)+\text{NO}(g)⟶\text{Ag}(s)+{\text{NO}}_3{}^{\text{−}}(aq)(\text{acidic solution})$

(c) ${\text{SiO}}_3{}^{\mathrm{2-}}(aq)+\text{Mg}(s)⟶\text{Si}(s)+\text{Mg}{\left(\text{OH}\right)}_2(s)\text{(basic solution)}$

(d) ${\text{ClO}}_3{}^{\text{−}}(aq)+{\text{MnO}}_2(s)⟶{\text{Cl}}^{\text{−}}(aq)+{\text{MnO}}_4{}^{\text{−}}(aq)\text{(basic solution)}$

From the information provided, use cell notation to describe the following systems:

(a) In one half-cell, a solution of Pt(NO₃)₂ forms Pt metal, while in the other half-cell, Cu metal goes into a Cu(NO₃)₂ solution with all solute concentrations 1 M.

(b) The cathode consists of a gold electrode in a 0.55 M Au(NO₃)₃ solution and the anode is a magnesium electrode in 0.75 M Mg(NO₃)₂ solution.

(c) One half-cell consists of a silver electrode in a 1 M AgNO₃ solution, and in the other half-cell, a copper electrode in 1 M Cu(NO₃)₂ is oxidized.

An active (metal) electrode was found to gain mass as the oxidation-reduction reaction was allowed to proceed. Was the electrode part of the anode or cathode? Explain.

The mass of three different metal electrodes, each from a different galvanic cell, were determined before and after the current generated by the oxidation-reduction reaction in each cell was allowed to flow for a few minutes. The first metal electrode, given the label A, was found to have increased in mass; the second metal electrode, given the label B, did not change in mass; and the third metal electrode, given the label C, was found to have lost mass. Make an educated guess as to which electrodes were active and which were inert electrodes, and which were anode(s) and which were the cathode(s).

For each reaction listed, determine its standard cell potential at 25 °C and whether the reaction is spontaneous at standard conditions.

(a) $\text{Mn}(s)+{\text{Ni}}^{\mathrm{2+}}(aq)⟶{\text{Mn}}^{\mathrm{2+}}(aq)+\text{Ni}(s)$

(b) $3{\text{Cu}}^{\mathrm{2+}}(aq)+\text{2Al}(s)⟶{\text{2Al}}^{\mathrm{3+}}(aq)+\text{3Cu}(s)$

(c) $\text{Na}(s)+{\text{LiNO}}_3(aq)⟶{\text{NaNO}}_3(aq)+\text{Li}(s)$

(d) ${\text{Ca(NO}}_3{)}_2(aq)+\text{Ba}(s)⟶{\text{Ba(NO}}_3{)}_2(aq)+\text{Ca}(s)$

Determine the overall reaction and its standard cell potential at 25 °C for the reaction involving the galvanic cell made from a half-cell consisting of a silver electrode in 1 M silver nitrate solution and a half-cell consisting of a zinc electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?

Determine the overall reaction and its standard cell potential at 25 °C for these reactions. Is the reaction spontaneous at standard conditions? Assume the standard reduction for Br₂(l) is the same as for Br₂(aq).
$\text{Pt}(s)│{\text{H}}_2(g)│{\text{H}}^{\text{+}}(aq)║{\text{Br}}_2(aq),{\text{Br}}^{\text{−}}(aq)│\text{Pt}(s)$

For the ΔG° values given here, determine the standard cell potential for the cell.

(a) 12 kJ/mol, n = 3

(b) −45 kJ/mol, n = 1

Determine ΔG and ΔG° for each of the reactions in the previous problem.

What are the desirable qualities of an electric battery?

Consider a battery made from one half-cell that consists of a copper electrode in 1 M CuSO₄ solution and another half-cell that consists of a lead electrode in 1 M Pb(NO₃)₂ solution.

(a) What are the reactions at the anode, cathode, and the overall reaction?

(b) What is the standard cell potential for the battery?

(c) Most devices designed to use dry-cell batteries can operate between 1.0 and 1.5 V. Could this cell be used to make a battery that could replace a dry-cell battery? Why or why not.

(d) Suppose sulfuric acid is added to the half-cell with the lead electrode and some PbSO₄(s) forms. Would the cell potential increase, decrease, or remain the same?

An inventor proposes using a SHE (standard hydrogen electrode) in a new battery for smartphones that also removes toxic carbon monoxide from the air:
$\begin{array}{l}\underset{¯}{\begin{array}{l}\text{Anode: CO}(g)+{\text{H}}_2\text{O}(l)⟶{\text{CO}}_2(g)+{\text{2H}}^{\text{+}}(aq)+{\text{2e}}^{\text{−}}{E}_{\text{anode}}^{°}=\text{−0.53 V}\\ \text{Cathode:}2{\text{H}}^{\text{+}}(aq)+{\text{2e}}^{\text{−}}⟶{\text{H}}_2(g)E_{\text{cathode}}^{°}=\text{0 V}\end{array}}\\ \text{Overall: CO}(g)+{\text{H}}_2\text{O}(l)⟶{\text{CO}}_2(g)+{\text{H}}_2(g)E_{\text{cell}}^{°}=\text{+0.53 V}\end{array}$

Would this make a good battery for smartphones? Why or why not?

Explain what happens to battery voltage as a battery is used, in terms of the Nernst equation.

Which member of each pair of metals is more likely to corrode (oxidize)?

(a) Mg or Ca

(b) Au or Hg

(c) Fe or Zn

(d) Ag or Pt

Aluminum $(E_{{\text{Al}}^{\text{3+}}\text{/Al}}^{°}=\text{−2.07 V})$ is more easily oxidized than iron $(E_{{\text{Fe}}^{\text{3+}}\text{/Fe}}^{°}=\text{−0.477 V}),$ and yet when both are exposed to the environment, untreated aluminum has very good corrosion resistance while the corrosion resistance of untreated iron is poor. Explain this observation.

Suppose you have three different metals, A, B, and C. When metals A and B come into contact, B corrodes and A does not corrode. When metals A and C come into contact, A corrodes and C does not corrode. Based on this information, which metal corrodes and which metal does not corrode when B and C come into contact?

Identify the reaction at the anode, reaction at the cathode, the overall reaction, and the approximate potential required for the electrolysis of the following molten salts. Assume standard states and that the standard reduction potentials in Appendix L are the same as those at each of the melting points. Assume the efficiency is 100%.

(a) CaCl₂

(b) LiH

(c) AlCl₃

(d) CrBr₃

How long would it take to reduce 1 mole of each of the following ions using the current indicated? Assume the voltage is sufficient to perform the reduction.

(a) Al³⁺, 1.234 A

(b) Ca²⁺, 22.2 A

(c) Cr⁵⁺, 37.45 A

(d) Au³⁺, 3.57 A

An irregularly shaped metal part made from a particular alloy was galvanized with zinc using a Zn(NO₃)₂ solution. When a current of 2.599 A was used, it took exactly 1 hour to deposit a 0.01123-mm layer of zinc on the part. What was the total surface area of the part? The density of zinc is 7.140 g/cm³. Assume the efficiency is 100%.