Electrochemistry: Chemistry 9701 (Cambridge International AS & A Level)

A few drops of potassium iodide solution are added to a concentrated solution of hydrogen peroxide, where the iodide ion acts as a catalyst. The mixture rapidly decomposes with vigorous fizzing, releasing a stream of oxygen gas:

$2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g)$

What is the oxidation number of oxygen in $\text{H}_2\text{O}_2(aq)$, and what type of redox process does oxygen as a whole undergo in this reaction?

Sodium chlorate(V), $\text{NaClO}_3$, is used industrially as a bleaching agent and, in dilute acidified solution, as a laboratory oxidising agent. When acidified sodium chlorate(V) solution is mixed with a solution of sodium sulfite, $\text{Na}_2\text{SO}_3$, the sulfite ions are oxidised to sulfate ions while the chlorate(V) ions are reduced to chloride ions.

(a) Determine the oxidation number of chlorine in $\text{ClO}_3^-$ and in $\text{Cl}^-$, and of sulfur in $\text{SO}_3^{2-}$ and in $\text{SO}_4^{2-}$. [2]

(b) Using oxidation numbers, construct two separate balanced ionic half-equations, including $\text{H}^+(aq)$ and $\text{H}_2\text{O}(l)$ as appropriate: one for the reduction of $\text{ClO}_3^-(aq)$ to $\text{Cl}^-(aq)$, and one for the oxidation of $\text{SO}_3^{2-}(aq)$ to $\text{SO}_4^{2-}(aq)$, both in acidic solution. [3]

(c) Combine your two half-equations from (b) to give the overall balanced ionic equation for the reaction, showing that both the atoms and the charge balance. [2]

(d) Identify, with a reason, the oxidising agent and the reducing agent in this reaction. [1]

The standard electrode potentials of two half-cells are given below:

$\text{Sn}^{4+}(aq) + 2e^- \rightleftharpoons \text{Sn}^{2+}(aq) \qquad E^{\ominus} = +0.15\ \text{V}$

$\text{Br}_2(aq) + 2e^- \rightleftharpoons 2\text{Br}^-(aq) \qquad E^{\ominus} = +1.09\ \text{V}$

Both half-cells use inert platinum electrodes, since neither redox couple includes a solid metal.

(a) Construct the cell diagram (cell notation) for the electrochemical cell formed from these two half-cells under standard conditions. [2]

(b) Calculate the standard cell potential, $E_{cell}^{\ominus}$, for this cell. [1]

(c) State, with a reason based on the $E^{\ominus}$ values given above, whether aqueous bromine is able to oxidise $\text{Sn}^{2+}(aq)$ to $\text{Sn}^{4+}(aq)$ under standard conditions. [2]

(d) Write the overall ionic equation for the feasible reaction identified in (c). [2]

A steel spoon is electroplated with silver. The spoon is made the cathode in a cell containing aqueous silver nitrate, with a pure silver rod as the anode, and a constant current of $0.850\ \text{A}$ is passed for $15.0$ minutes.

(a) Write the half-equation for the reaction occurring at the cathode. [1]

(b) Calculate the quantity of electric charge, $Q$, passed during the electroplating process. [1]

(c) Calculate the amount, in mol, of electrons transferred, and hence the mass of silver deposited on the spoon. ($F = 96500\ \text{C mol}^{-1}$; $A_r(\text{Ag}) = 108$) [3]

(d) The same total charge calculated in (b) is then passed through a second electrolytic cell, connected in series with the first, containing molten aluminium oxide dissolved in cryolite (as in the industrial extraction of aluminium). Calculate the mass of aluminium that would be deposited at the cathode of this second cell by the same quantity of charge. ($A_r(\text{Al}) = 27$) [3]

A $\text{Cu}^{2+}(aq)/\text{Cu}(s)$ half-cell is set up under standard conditions, with $[\text{Cu}^{2+}(aq)] = 1.00\ \text{mol dm}^{-3}$, giving a standard electrode potential of $+0.34\ \text{V}$:

$\text{Cu}^{2+}(aq) + 2e^- \rightleftharpoons \text{Cu}(s)$

Distilled water is then added to the half-cell, reducing the concentration of $\text{Cu}^{2+}(aq)$ to well below $1.00\ \text{mol dm}^{-3}$, with the temperature and the copper electrode unchanged.

What is the effect on the electrode potential of this half-cell, and why?

Acidified potassium manganate(VII) solution is a strong oxidising agent. During its reaction with a reducing agent, the manganate(VII) ion, $\text{MnO}_4^-(aq)$, is reduced to $\text{Mn}^{2+}(aq)$ in acidic solution.

Which of the following is the correctly balanced ionic half-equation for this reduction?

Copper reacts with dilute nitric acid to form copper(II) nitrate solution, nitrogen monoxide gas and water:

$\text{Cu}(s) + \text{HNO}_3(aq) \rightarrow \text{Cu(NO}_3)_2(aq) + \text{NO}(g) + \text{H}_2\text{O}(l) \quad \text{(unbalanced)}$

(a) State the oxidation number of copper in $\text{Cu}(s)$, of nitrogen in $\text{HNO}_3(aq)$, and of nitrogen in $\text{NO}(g)$. [2]

(b) Using the change in oxidation number of copper and of the nitrogen atoms that are reduced, deduce the simplest whole-number ratio of moles of copper oxidised to moles of nitrogen atoms reduced, such that the total number of electrons lost equals the total number of electrons gained. [2]

(c) Hence deduce the balancing numbers for the full equation above, and explain how the number of $\text{HNO}_3(aq)$ formula units is obtained even though not all of the nitrogen atoms are reduced. [3]

(d) Verify that your balanced equation from (c) conserves both atoms and charge (noting that this is a reaction between neutral molecules, so the overall charge on each side should be zero). [2]

A student wants to measure the standard electrode potential of the $\text{Ni}^{2+}(aq)/\text{Ni}(s)$ half-cell. They set up a nickel electrode dipping into $1.00\ \text{mol dm}^{-3}\ \text{Ni}^{2+}(aq)$ solution, connected by a salt bridge to a standard hydrogen electrode, with a high-resistance voltmeter completing the external circuit. At $298\ \text{K}$, the voltmeter reads $0.25\ \text{V}$, and electrons are found to flow through the external wire from the nickel electrode to the platinum/hydrogen electrode.

(a) State the conditions required for the platinum/hydrogen electrode to function as a standard hydrogen electrode. [2]

(b) State which electrode is negative, and write the half-equation occurring at each electrode. [3]

(c) Given that $E^{\ominus}$ of the standard hydrogen electrode is defined as $0.00\ \text{V}$, calculate the standard electrode potential, $E^{\ominus}$, of the $\text{Ni}^{2+}(aq)/\text{Ni}(s)$ half-cell. [2]

(d) Write the cell notation for this electrochemical cell. [2]

The table below gives the standard electrode potentials of four metal/metal-ion half-cells:

Based only on these standard electrode potentials, which of the following reactions is feasible under standard conditions?

Dilute sulfuric acid is electrolysed using inert platinum electrodes. A constant current of $0.500\ \text{A}$ is passed for $40.0$ minutes. ($F = 96500\ \text{C mol}^{-1}$; 1 mole of gas occupies $24000\ \text{cm}^3$ at room temperature and pressure, rtp.)

(a) Write half-equations for the reactions occurring at the cathode and at the anode. [2]

(b) Calculate the quantity of electric charge, $Q$, passed during electrolysis. [1]

(c) Calculate the amount, in mol, of electrons transferred, and hence the volume of gas produced at the cathode, measured at rtp. [3]

(d) Calculate the volume of gas produced at the anode over the same time period, measured at rtp, and state the simplest whole-number ratio of the volume of gas at the cathode to the volume of gas at the anode. [3]