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Homework 1: Electrochemistry

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    8768

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    Q1.1

    Determine the oxidation states of the underlined elements in

    1. Cl2
    2. NaH
    3. H2CO
    4. S2O3 2-
    5. KMnO4
    6. FeCl3
    7. N2
    8. H2SO4
    9. HClO2
    10. CuSO4

    Q1.2

    Balance the following equations in both acidic and basic environments:

    1. \(H_2(g) + O_2(g) \rightarrow H_2O(l)\)
    2. \(Cr_2O_7^{2-}(aq) + C_2H_5OH (l) \rightarrow Cr^{3+}(aq) + CO_2(g)\)
    3. \(Fe^{2+}(aq) + MnO_4^-(aq) \rightarrow Fe^{3+}(aq) + Mn^{2+}(aq)\)
    4. \(Zn (s) + NO_3^-(aq) \rightarrow Zn^{2+}(aq) + NO(g)\)
    5. \(Al (s) + H_2O (l) + O_2 (g) \rightarrow [Al(OH)_4]^-(aq)\)

    Q1.3

    Using the table of Standard Reduction Potential and the following clues, suggest two possible elements for the metal \(\ce{M}\).

    1. \(\ce{M}\) will be oxidized by \(\ce{F2(g)}\), but will be reduced by \(\ce{Fe^2+(aq)}\)
    2. \(\ce{M}\) will reduce \(\ce{Br2(l)}\), but will oxidize \(\ce{I2(s)}\)

    Q1.4

    Write cell reactions for the electrochemical cells diagrammed here, and calculate \(\mathrm{E^\circ_{cell}}\) for each reaction.

    1. \(\mathrm{Al(s)| Al^{3+}|| Zn^{2+}(aq)| Zn(s)}\)
    2. \(\mathrm{Pt (s)| Fe^{2+}(aq), Fe^{3+}(aq)|| Cu^{2+}(aq)| Cu(s)}\)
    3. \(\mathrm{Pt (s)| Cr^{3+}(aq), Cr_2O_7^{2-}(aq)|| Ag^+(aq)| Ag(s)}\)
    4. \(\mathrm{O_2^-(aq)| O_2(g)|| H^+(aq)|H_2(g)| C(s)}\)

    Q1.5

    Use the table of Standard Reduction Potential to predict whether these reactions will happen spontaneously.

    1. \(\mathrm{2Ag^+(aq) + Cu(s) \rightarrow 2Ag(s) + Cu^{2+}(aq)}\)
    2. \(\mathrm{Fe^{3+}(aq) + Na(s) \rightarrow Fe^{2+}(aq) + Na^+(aq)}\)
    3. \(\mathrm{Zn^{2+}(aq) + 2I^-(aq) \rightarrow Zn(s) + I_2(l)}\)

    Q1.6

    For the electrochemical cells shown below under standard conditions, write the cell reactions and find \(\mathrm{E^\circ_{cell}}\) using the Standard Reduction Potential table. Hint: Water is not added to cell diagrams.

    1. \(\mathrm{Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)}\)
    2. \(\mathrm{Cu(s) | Cu^{2+}(aq) || Fe^{2+}(aq),\,Fe^{3+}(aq) | Pt(s)}\)
    3. \(\mathrm{PbO_2(s) | Pb^{2+}(aq) || Mn^{3+}(aq) | MnO_2(s)}\)
    4. \(\mathrm{Cu(s) | Cu^{2+}(aq) || O_2(g) | H_2O_2(aq) | Pt(s)}\)

    Q1.7

    Sketch a voltaic cell for each of the following conditions, labeling the anode, cathode, and electron flow. Balance the redox equation if necessary. Calculate \(\mathrm{E^\circ_{cell}}\)​ for each cell.

    1. \(\mathrm{Cu^{2+}(aq) + Sn^{2+}(aq) \rightarrow Cu(s) + Sn^{4+}(aq)}\)
    2. \(\ce{Mg(s)}\)​ displaces \(\ce{Sn^2+(aq)}\)​ from solution
    3. \(\ce{Zn(s)}\)​ donates e- to acidic solution to form \(\ce{H2(g)}\)​
    4. \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) + Sn^{2+}(aq)}\)

    Q1.8

    Calculate the oxidation number for nitrogen in the following substances:

    1. NH3
    2. N2
    3. NO2
    4. NO3-

    Q1.9

    In the oxidation-reduction reaction,

    \[\ce{Br2(l) + 2 I- (aq) \rightarrow 2 Br- (aq) + I2(s)}\]

    1. which substance(s) is being reduced?
    2. which element(s) is increasing in oxidation number?
    3. which element(s) is gaining electrons?
    4. which substance(s) is the oxidizing agent?
    5. which substance(s) is the reducing agent?

    Q1.10

    Classify each of the following substances as an oxidizing agent, reducing agent or both. List the oxidizing agents in order of decreasing strength; list the reducing agents in order of decreasing strength (use SRP Table):

    Ni(s), H+(aq), Au(s), Cl2(g), Sn2+(aq), Mg(s), Fe2+(aq)

    More Question (without solutions)

    Q1M.1

    Calculate the oxidation number for nitrogen in the following substances:

    1. \(\ce{NH3}\)
    2. \(\ce{N2}\)
    3. \(\ce{NO2}\)
    4. \(\ce{NO3-}\)

    Q1M.2

    In the oxidation-reduction reaction,

    \[\ce{Br2(l) + 2 I- (aq) \rightarrow 2 Br- (aq) + I2(s)}\]

    1. which substance(s) is being reduced?
    2. which element(s) is increasing in oxidation number?
    3. which element(s) is gaining electrons?
    4. which substance(s) is the oxidizing agent?
    5. which substance(s) is the reducing agent?

    Q1M.3

    Classify each of the following substances as an oxidizing agent, reducing agent or both. List the oxidizing agents in order of decreasing strength; list the reducing agents in order of decreasing strength (use SRP Table):

    \(\ce{Ni(s)}\), \(\ce{H+ (aq)}\), \(\ce{Au(s)}\), \(\ce{Cl2(g)}\), \(\ce{Sn^2+ (aq)}\), \(\ce{Mg(s)}\), \(\ce{Fe^2+ (aq)}\)

    Q1M.4

    For each of the following reactions: 1) identify the oxidation and reduction half-equations, and 2) balance the equation (adding \(\ce{H+}\) and \(\ce{H2O}\) as needed), 3) find \(\mathrm{E^\circ}\) (in volts) and 4) determine whether the reaction is spontaneous under standard conditions.

    1. \(\ce{Ag(s) + Cu^2+(aq) \rightarrow Ag+(aq) + Cu(s)}\)
    2. \(\ce{Ni(s) + MnO_4- (aq) \rightarrow Ni^2+(aq) + Mn^2+(aq)}\)
    3. \(\ce{Mn^2+(aq) + NO3- (aq) \rightarrow MnO2(s) + NO(g)}\)

    Q1M.5

    For each of the following reactions: 1) find \(\mathrm{E^\circ}\) (in volts) and 2) determine whether the reaction is spontaneous under standard conditions.

    1. the reaction between iron and iron(III) ions to give iron(II) ions.
    2. the following cell: \(\mathrm{ I^- \, |\, I_2 \,||\, Zn^{2+} \,|\, Zn}\)

    Q1M.6

    For the electrochemical cells shown below, write the cell reactions and find \(\mathrm{E^\circ_{cell}}\) using the Standard Reduction Potential table

    1. \(\mathrm{Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)}\)
    2. \(\mathrm{Cu(s) | Cu^{2+}(aq) || Fe^{2+}(aq),\,Fe^{3+}(aq) | Pt(s)}\)
    3. \(\mathrm{PbO_2(s) | Pb^{2+}(aq) || Mn^{2+}(aq) | MnO_2(s)}\)
    4. \(\mathrm{Cu(s) | Cu^{2+}(aq) || O_2(g) | H_2O_2(aq) | Pt(s)}\)

    Q1M.7

    Sketch a voltaic cell for each of the following conditions, labeling the anode, cathode, and electron flow. Balance the redox equation if necessary. Calculate \(\mathrm{E^\circ_{cell}}\)​ for each cell.

    1. \(\mathrm{Cu^{2+}(aq) + Sn^{2+}(aq) \rightarrow Cu(s) + Sn^{4+}(aq)}\)
    2. \(\ce{Mg(s)}\)​ displaces \(\ce{Sn^2+(aq)}\)​ from solution
    3. \(\ce{Zn(s)}\)​ donates e- to acidic solution to form \(\ce{H2(g)}\)​
    4. \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) + Sn^{2+}(aq)}\)

    Q1M.8

    Determine \(\mathrm{\Delta G^\circ}\)​ for the following voltaic cell reactions:

    1. \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) +Sn^{2+}(aq)}\)
    2. \(\mathrm{O_2(g) +2H^+(aq) + 2F^-(aq) \rightarrow H_2O_2(aq) + F_2(g)}\)
    3. \(\mathrm{Br_2(l) +2Fe^{2+}(aq) \rightarrow 2Br^-(aq) + 2Fe^{3+}(aq)}\)

    Q1M.9

    Which of the following ions will oxidize \(\ce{Br-}\) ion to \(\ce{Br2}\)?

    1. \(\ce{Pb^2+}\)
    2. \(\ce{H+}\)
    3. \(\ce{Au^3+}\)
    4. \(\ce{MnO4-}\)

    Q1M.10

    A voltaic cell has an aluminum electrode in \(\ce{Al2(SO4)3}\) solution in one compartment and the other compartment has a lead electrode in \(\ce{PbSO4}\) solution.

    1. Which has a greater tendency to be oxidized, \(\ce{Al}\) or \(\ce{Pb}\)? Write a balanced equation for the spontaneous reaction.
    2. Draw a diagram of the voltaic cell, including
      1. the anode and cathode
      2. the direction of flow of electrons, positive ions and negative ions
    3. Which electrode will increase in mass?