Key Terms, Key Equations, Summaries, and Exercises (Chapter 16)

Identify each half-reaction below as either oxidation or reduction.

(a) Cl−⟶Cl2${\text{Cl}}^{\text{−}}⟶{\text{Cl}}_2$

(b) Mn2+⟶MnO2${\text{Mn}}^{\mathrm{2+}}⟶{\text{MnO}}_2$

(c) H2⟶H+${\text{H}}_2⟶{\text{H}}^{\text{+}}$

(d) NO3−⟶NO${\text{NO}}_3{}^{-}⟶\text{NO}$

Balance the equations below assuming they occur in an acidic solution.

(a) H2O2+Sn2+⟶H2O+Sn4+${\text{H}}_2{\text{O}}_2+{\text{Sn}}^{\mathrm{2+}}⟶{\text{H}}_2\text{O}+{\text{Sn}}^{\mathrm{4+}}$

(b) PbO2+Hg⟶Hg22++Pb2+${\text{PbO}}_2+\text{Hg}⟶{\text{Hg}}_2{}^{\mathrm{2+}}+{\text{Pb}}^{\mathrm{2+}}$

(c) Al+Cr2O72−⟶Al3++Cr3+$\text{Al}+{\text{Cr}}_2{\text{O}}_7{}^{\mathrm{2-}}⟶{\text{Al}}^{\mathrm{3+}}+{\text{Cr}}^{\mathrm{3+}}$

Balance the equations below assuming they occur in a basic solution.

(a) SO32−(𝑎𝑞)+Cu(OH)2(𝑠)⟶SO42−(𝑎𝑞)+Cu(OH)(𝑠)${\text{SO}}_3{}^{\mathrm{2-}}(aq)+{\text{Cu(OH)}}_2(s)⟶{\text{SO}}_4{}^{\mathrm{2-}}(aq)+\text{Cu(OH)}(s)$

(b) O2(𝑔)+Mn(OH)2(𝑠)⟶MnO2(𝑠)${\text{O}}_2(g)+{\text{Mn(OH)}}_2(s)⟶{\text{MnO}}_2(s)$

(c) NO3−(𝑎𝑞)+H2(𝑔)⟶NO(𝑔)${\text{NO}}_3{}^{\text{−}}(aq)+{\text{H}}_2(g)⟶\text{NO}(g)$

(d) Al(𝑠)+CrO42−(𝑎𝑞)⟶Al(OH)3(𝑠)+Cr(OH)4−(𝑎𝑞)$\text{Al}(s)+{\text{CrO}}_4{}^{\mathrm{2-}}(aq)⟶{\text{Al(OH)}}_3(s)+{\text{Cr(OH)}}_4{}^{\text{−}}(aq)$

Why don’t hydroxide ions appear in equations for half-reactions occurring in acidic solution?

Why must the charge balance in oxidation-reduction reactions?

Assuming the schematics below represent galvanic cells as written, identify the half-cell reactions occurring in each.

(a) Mg(𝑠)│Mg2+(𝑎𝑞)║Cu2+(𝑎𝑞)│Cu(𝑠)$\text{Mg}(s)│{\text{Mg}}^{\mathrm{2+}}(aq)║{\text{Cu}}^{\mathrm{2+}}(aq)│\text{Cu}(s)$

(b) Ni(𝑠)│Ni2+(𝑎𝑞)║Ag+(𝑎𝑞)│Ag(𝑠)$\text{Ni}(s)│{\text{Ni}}^{\mathrm{2+}}(aq)║{\text{Ag}}^{\text{+}}(aq)│\text{Ag}(s)$

Balance each reaction below, and write a cell schematic representing the reaction as it would occur in a galvanic cell.

(a) Al(𝑠)+Zr4+(𝑎𝑞)⟶Al3+(𝑎𝑞)+Zr(𝑠)$\text{Al}(s)+{\text{Zr}}^{\mathrm{4+}}(aq)⟶{\text{Al}}^{\mathrm{3+}}(aq)+\text{Zr}(s)$

(b) Ag+(𝑎𝑞)+NO(𝑔)⟶Ag(𝑠)+NO3−(𝑎𝑞)(acidic solution)${\text{Ag}}^{\text{+}}(aq)+\text{NO}(g)⟶\text{Ag}(s)+{\text{NO}}_3{}^{\text{−}}(aq)(\text{acidic solution})$

(c) SiO32−(𝑎𝑞)+Mg(𝑠)⟶Si(𝑠)+Mg(OH)2(𝑠)(basic solution)${\text{SiO}}_3{}^{\mathrm{2-}}(aq)+\text{Mg}(s)⟶\text{Si}(s)+\text{Mg}{\left(\text{OH}\right)}_2(s)\text{(basic solution)}$

(d) ClO3−(𝑎𝑞)+MnO2(𝑠)⟶Cl−(𝑎𝑞)+MnO4−(𝑎𝑞)(basic solution)${\text{ClO}}_3{}^{\text{−}}(aq)+{\text{MnO}}_2(s)⟶{\text{Cl}}^{\text{−}}(aq)+{\text{MnO}}_4{}^{\text{−}}(aq)\text{(basic solution)}$

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 an anode or a cathode? Explain.

The masses of three electrodes (A, B, and C), each from three different galvanic cells, were measured before and after the cells were allowed to pass current for a while. The mass of electrode A increased, that of electrode B was unchanged, and that of electrode C decreased. Identify each electrode as active or inert, and note (if possible) whether it functioned as anode or cathode.

Calculate the standard cell potential for each reaction below, and note whether the reaction is spontaneous under standard state conditions.

(a) Mn(𝑠)+Ni2+(𝑎𝑞)⟶Mn2+(𝑎𝑞)+Ni(𝑠)$\text{Mn}(s)+{\text{Ni}}^{\mathrm{2+}}(aq)⟶{\text{Mn}}^{\mathrm{2+}}(aq)+\text{Ni}(s)$

(b) 3Cu2+(𝑎𝑞)+2Al(𝑠)⟶2Al3+(𝑎𝑞)+3Cu(𝑠)$3{\text{Cu}}^{\mathrm{2+}}(aq)+\text{2Al}(s)⟶{\text{2Al}}^{\mathrm{3+}}(aq)+\text{3Cu}(s)$

(c) Na(𝑠)+LiNO3(𝑎𝑞)⟶NaNO3(𝑎𝑞)+Li(𝑠)$\text{Na}(s)+{\text{LiNO}}_3(aq)⟶{\text{NaNO}}_3(aq)+\text{Li}(s)$

(d) Ca(NO3)2(𝑎𝑞)+Ba(𝑠)⟶Ba(NO3)2(𝑎𝑞)+Ca(𝑠)${\text{Ca(NO}}_3{)}_2(aq)+\text{Ba}(s)⟶{\text{Ba(NO}}_3{)}_2(aq)+\text{Ca}(s)$

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

Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and note whether the reaction is spontaneous under standard state conditions.
Pt(𝑠)│H2(𝑔)│H+(𝑎𝑞)║Br2(𝑎𝑞),Br−(𝑎𝑞)│Pt(𝑠)$\text{Pt}(s)│{\text{H}}_2(g)│{\text{H}}^{\text{+}}(aq)║{\text{Br}}_2(aq),{\text{Br}}^{\text{−}}(aq)│\text{Pt}(s)$

For each pair of standard free energy change and electron stoichiometry values below, calculate a corresponding standard cell potential.

(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.

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 is the standard cell potential for the battery?

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

(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?

Use the Nernst equation to explain the drop in voltage observed for some batteries as they discharge.

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∘Al3+/Al=−2.07 V)$(E_{{\text{Al}}^{\text{3+}}\text{/Al}}^{°}=\text{−2.07 V})$ is more easily oxidized than iron (𝐸∘Fe3+/Fe=−0.477 V),$(E_{{\text{Fe}}^{\text{3+}}\text{/Fe}}^{°}=\text{−0.477 V}),$ 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. What might 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?

If a 2.5 A current flows through a circuit for 35 minutes, how many coulombs of charge moved through the circuit?

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

Write the half-reactions and cell reaction occurring during electrolysis of each molten salt below.

(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?

(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.