AQA A-Level Chemistry: Redox & Electrochemistry

The standard electrode potential of a metal/metal-ion system cannot be measured in isolation; it is always measured relative to a reference half-cell under a fixed set of conditions.

Describe and explain how the standard electrode potential, $E^{\circ}$E∘, of a metal/metal-ion system (for example $\text{Zn}^{2+}(aq)\,/\,\text{Zn}(s)$Zn2+(aq)/Zn(s)) is measured. In your answer you should refer to:

• the construction and operation of the standard hydrogen electrode, and why its $E^{\circ}$E∘ is taken to be $0.00 \text{ V}$0.00 V;
• the standard conditions used;
• the roles of the salt bridge and the high-resistance voltmeter;
• what the sign and magnitude of the measured $E^{\circ}$E∘ tell you.

(6 marks)

An electrochemical cell is set up by combining the two half-cells below under standard conditions, joined by a salt bridge and a high-resistance voltmeter.

Half-equation (written as a reduction)	$E^{\circ}$E∘ / V
$\text{Ag}^{+}(aq) + e^{-} \rightleftharpoons \text{Ag}(s)$Ag+(aq)+e−⇌Ag(s)	$+0.80$+0.80
$\text{Ni}^{2+}(aq) + 2e^{-} \rightleftharpoons \text{Ni}(s)$Ni2+(aq)+2e−⇌Ni(s)	$-0.25$−0.25

(a) Identify the positive electrode and the negative electrode, giving a reason. (2 marks)

(b) Calculate the standard cell EMF, $E^{\circ}_{\text{cell}}$Ecell∘​. (1 mark)

(c) Write the overall cell equation for the spontaneous reaction. (2 marks)

(d) State whether the cell reaction is feasible under standard conditions, and justify your answer. (1 mark)

A single iron supplement tablet, of mass $1.50 \text{ g}$1.50 g, contains iron as iron(II). The whole tablet was dissolved in excess dilute sulfuric acid and the solution made up to $250 \text{ cm}^3$250 cm3 in a volumetric flask. $25.0 \text{ cm}^3$25.0 cm3 portions of this solution were titrated against $0.0200 \text{ mol dm}^{-3}$0.0200 mol dm−3 potassium manganate(VII).

The ionic equation for the reaction is:

$\text{MnO}_4^{-} + 8\text{H}^{+} + 5\text{Fe}^{2+} \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} + 5\text{Fe}^{3+}$MnO4−​+8H++5Fe2+→Mn2++4H2​O+5Fe3+

Titration	Mean titre of manganate(VII) / cm³
Concordant mean	24.0

$(A_r:\ \text{Fe} = 55.8)$(Ar​: Fe=55.8)

(a) Calculate the amount, in moles, of $\text{MnO}_4^{-}$MnO4−​ used in one titration. (1 mark)

(b) Calculate the mass of iron in the whole tablet. (3 marks)

(c) Calculate the percentage by mass of iron in the tablet. (1 mark)

In an industrial electroplating process, copper is deposited onto a steel component from a solution of copper(II) sulfate. The half-equation at the cathode is:

$\text{Cu}^{2+}(aq) + 2e^{-} \rightarrow \text{Cu}(s)$Cu2+(aq)+2e−→Cu(s)

A steady current of $1.50 \text{ A}$1.50 A is passed for $45.0$45.0 minutes.

$(\text{Faraday constant } F = 96500 \text{ C mol}^{-1};\ A_r:\ \text{Cu} = 63.5)$(Faraday constant F=96500 C mol−1; Ar​: Cu=63.5)

Calculate the mass of copper deposited on the component. Show your working clearly. (5 marks)

A student is given the following standard electrode potentials.

Half-equation (written as a reduction)	$E^{\circ}$E∘ / V
$\text{Fe}^{3+}(aq) + e^{-} \rightleftharpoons \text{Fe}^{2+}(aq)$Fe3+(aq)+e−⇌Fe2+(aq)	$+0.77$+0.77
$\frac{1}{2}\text{I}_2(aq) + e^{-} \rightleftharpoons \text{I}^{-}(aq)$21​I2​(aq)+e−⇌I−(aq)	$+0.54$+0.54
$\frac{1}{2}\text{Br}_2(aq) + e^{-} \rightleftharpoons \text{Br}^{-}(aq)$21​Br2​(aq)+e−⇌Br−(aq)	$+1.07$+1.07

(a) Use the $E^{\circ}$E∘ values to predict whether $\text{Fe}^{3+}(aq)$Fe3+(aq) will oxidise iodide ions, $\text{I}^{-}(aq)$I−(aq). Support your answer with a value of $E^{\circ}_{\text{cell}}$Ecell∘​. (2 marks)

(b) Predict whether $\text{Fe}^{3+}(aq)$Fe3+(aq) will oxidise bromide ions, $\text{Br}^{-}(aq)$Br−(aq), supporting your answer with a value of $E^{\circ}_{\text{cell}}$Ecell∘​. (1 mark)

(c) State one limitation of using standard electrode potentials to predict whether a redox reaction will actually occur. (1 mark)

In acidic solution, manganate(VII) ions react with ethanedioate (oxalate) ions according to the equation:

$2\text{MnO}_4^{-} + 5\text{C}_2\text{O}_4^{2-} + 16\text{H}^{+} \rightarrow 2\text{Mn}^{2+} + 10\text{CO}_2 + 8\text{H}_2\text{O}$2MnO4−​+5C2​O42−​+16H+→2Mn2++10CO2​+8H2​O

(a) Deduce the oxidation state of manganese in $\text{MnO}_4^{-}$MnO4−​ and the oxidation state of carbon in $\text{C}_2\text{O}_4^{2-}$C2​O42−​. (2 marks)

(b) Using oxidation states, identify the species that is oxidised and the species that is reduced in this reaction. (1 mark)