Homework 1: Electrochemistry
- Page ID
- 8768
There are select solutions to these problems here.
Q1.1
Determine the oxidation states of the underlined elements in
- Cl2
- NaH
- H2CO
- S2O3 2-
- KMnO4
- FeCl3
- N2
- H2SO4
- HClO2
- CuSO4
Q1.2
Balance the following equations in both acidic and basic environments:
- \(H_2(g) + O_2(g) \rightarrow H_2O(l)\)
- \(Cr_2O_7^{2-}(aq) + C_2H_5OH (l) \rightarrow Cr^{3+}(aq) + CO_2(g)\)
- \(Fe^{2+}(aq) + MnO_4^-(aq) \rightarrow Fe^{3+}(aq) + Mn^{2+}(aq)\)
- \(Zn (s) + NO_3^-(aq) \rightarrow Zn^{2+}(aq) + NO(g)\)
- \(Al (s) + H_2O (l) + O_2 (g) \rightarrow [Al(OH)_4]^-(aq)\)
Q1.3
Using the table of Standard Reduction Potential and the following clues, suggest two possible elements for the metal \(\ce{M}\).
- \(\ce{M}\) will be oxidized by \(\ce{F2(g)}\), but will be reduced by \(\ce{Fe^2+(aq)}\)
- \(\ce{M}\) 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.
- \(\mathrm{Al(s)| Al^{3+}|| Zn^{2+}(aq)| Zn(s)}\)
- \(\mathrm{Pt (s)| Fe^{2+}(aq), Fe^{3+}(aq)|| Cu^{2+}(aq)| Cu(s)}\)
- \(\mathrm{Pt (s)| Cr^{3+}(aq), Cr_2O_7^{2-}(aq)|| Ag^+(aq)| Ag(s)}\)
- \(\mathrm{O_2^-(aq)| O_2(g)|| H^+(aq)|H_2(g)| C(s)}\)
Q1.5
Use the table of Standard Reduction Potential to predict whether these reactions will happen spontaneously.
- \(\mathrm{2Ag^+(aq) + Cu(s) \rightarrow 2Ag(s) + Cu^{2+}(aq)}\)
- \(\mathrm{Fe^{3+}(aq) + Na(s) \rightarrow Fe^{2+}(aq) + Na^+(aq)}\)
- \(\mathrm{Zn^{2+}(aq) + 2I^-(aq) \rightarrow Zn(s) + I_2(l)}\)
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.
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)}\)
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || Fe^{2+}(aq),\,Fe^{3+}(aq) | Pt(s)}\)
- \(\mathrm{PbO_2(s) | Pb^{2+}(aq) || Mn^{3+}(aq) | MnO_2(s)}\)
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || O_2(g) | H_2O_2(aq) | Pt(s)}\)
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.
- \(\mathrm{Cu^{2+}(aq) + Sn^{2+}(aq) \rightarrow Cu(s) + Sn^{4+}(aq)}\)
- \(\ce{Mg(s)}\) displaces \(\ce{Sn^2+(aq)}\) from solution
- \(\ce{Zn(s)}\) donates e- to acidic solution to form \(\ce{H2(g)}\)
- \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) + Sn^{2+}(aq)}\)
Q1.8
Calculate the oxidation number for nitrogen in the following substances:
- NH3
- N2
- NO2
- NO3-
Q1.9
In the oxidation-reduction reaction,
\[\ce{Br2(l) + 2 I- (aq) \rightarrow 2 Br- (aq) + I2(s)}\]
- which substance(s) is being reduced?
- which element(s) is increasing in oxidation number?
- which element(s) is gaining electrons?
- which substance(s) is the oxidizing agent?
- 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:
- \(\ce{NH3}\)
- \(\ce{N2}\)
- \(\ce{NO2}\)
- \(\ce{NO3-}\)
Q1M.2
In the oxidation-reduction reaction,
\[\ce{Br2(l) + 2 I- (aq) \rightarrow 2 Br- (aq) + I2(s)}\]
- which substance(s) is being reduced?
- which element(s) is increasing in oxidation number?
- which element(s) is gaining electrons?
- which substance(s) is the oxidizing agent?
- 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.
- \(\ce{Ag(s) + Cu^2+(aq) \rightarrow Ag+(aq) + Cu(s)}\)
- \(\ce{Ni(s) + MnO_4- (aq) \rightarrow Ni^2+(aq) + Mn^2+(aq)}\)
- \(\ce{Mn^2+(aq) + NO3- (aq) \rightarrow MnO2(s) + NO(g)}\)
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.
- the reaction between iron and iron(III) ions to give iron(II) ions.
- 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
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || Ag^+(aq) | Ag(s)}\)
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || Fe^{2+}(aq),\,Fe^{3+}(aq) | Pt(s)}\)
- \(\mathrm{PbO_2(s) | Pb^{2+}(aq) || Mn^{2+}(aq) | MnO_2(s)}\)
- \(\mathrm{Cu(s) | Cu^{2+}(aq) || O_2(g) | H_2O_2(aq) | Pt(s)}\)
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.
- \(\mathrm{Cu^{2+}(aq) + Sn^{2+}(aq) \rightarrow Cu(s) + Sn^{4+}(aq)}\)
- \(\ce{Mg(s)}\) displaces \(\ce{Sn^2+(aq)}\) from solution
- \(\ce{Zn(s)}\) donates e- to acidic solution to form \(\ce{H2(g)}\)
- \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) + Sn^{2+}(aq)}\)
Q1M.8
Determine \(\mathrm{\Delta G^\circ}\) for the following voltaic cell reactions:
- \(\mathrm{Pb^{2+}(aq) +Sn(s) \rightarrow Pb(s) +Sn^{2+}(aq)}\)
- \(\mathrm{O_2(g) +2H^+(aq) + 2F^-(aq) \rightarrow H_2O_2(aq) + F_2(g)}\)
- \(\mathrm{Br_2(l) +2Fe^{2+}(aq) \rightarrow 2Br^-(aq) + 2Fe^{3+}(aq)}\)
Q1M.9
Which of the following ions will oxidize \(\ce{Br-}\) ion to \(\ce{Br2}\)?
- \(\ce{Pb^2+}\)
- \(\ce{H+}\)
- \(\ce{Au^3+}\)
- \(\ce{MnO4-}\)
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.
- Which has a greater tendency to be oxidized, \(\ce{Al}\) or \(\ce{Pb}\)? Write a balanced equation for the spontaneous reaction.
- Draw a diagram of the voltaic cell, including
- the anode and cathode
- the direction of flow of electrons, positive ions and negative ions
- Which electrode will increase in mass?