CAIE A-Level Chemistry 24.2 Electrode and Cell Potentials

Practise standard electrodes, cell potentials, redox feasibility, concentration effects and electrochemical ΔG.

Exam points

draw and label half-cells with a salt bridge, voltmeter and inert electrode where required
calculate E°cell from reduction potentials and use its sign to select feasible cell direction
link concentration changes qualitatively to E and calculate ΔG° from −nE°cellF

Covered syllabus points

24.2.1—Terms 24.2.2—Standard hydrogen electrode 24.2.3—Measuring standard electrode potentials 24.2.4—A standard cell potential by combining two 24.2.5—Standard cell potentials 24.2.6—From E 24.2.7—Redox equations using the relevant 24.2.8—How the value of an electrode potential, E 24.2.9—Nernst equation 24.2.10—Gibbs free energy and cell potential

[Maximum number: 11]

(a)

Define the term standard cell potential.

[ 2 ]

the potential difference between two half-cells/two electrodes (in a cell)

under standard conditions of $1 \mathrm{~atm} ., 298 \mathrm{~K}$ , (all) solutions being $1 \mathrm{~mol} \mathrm{dm}^{-3}$

(b)

Draw a fully labelled diagram of the experimental set-up you could use to measure the standard electrode potential of the $\mathrm{Pb}^{2+}(\mathrm{aq}) / \mathrm{Pb}(\mathrm{s})$ electrode. Include the necessary chemicals.

[ 4 ]

Figure for Answer to Question (b) — CAIE A-Level Chemistry A2

8 marking points, any 2 points for each mark

$\mathrm{H}_{2}$ / hydrogen

correct delivery system for $\mathrm{H}_{2}$ $\mathrm{Pb}^{2+}$ (aq)

Pb electrode

Pt electrode

$\mathrm{H}^{+}(\mathrm{aq})$ solution

salt bridge

voltmeter/V labelled

(c)

The $E^{\ominus}$ for a $\mathrm{Pb}^{2+}(\mathrm{aq}) / \mathrm{Pb}(\mathrm{s})$ electrode is -0.13 V .

Suggest how the E for this electrode would differ from its $E^{\ominus}$ value if the concentration of $\mathrm{Pb}^{2+}(\mathrm{aq})$ ions is reduced. Indicate this by placing a tick $(\checkmark)$ in the appropriate box in the table.

Table for Question (c) — CAIE A-Level Chemistry A2

Explain your answer.

[ 2 ]

more negative

shifts $\mathrm{Pb}^{2+}\left(+2 \mathrm{e}^{-}\right) \leftrightharpoons \mathrm{Pb}$ equilibrium / reaction to the left

(d)

Car batteries are made up of rechargeable lead-acid cells. Each cell consists of a negative electrode made of Pb metal and a positive electrode made of $\mathrm{PbO}_{2}$ . The electrolyte is $\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq})$ .

When a lead-acid cell is in use, $\mathrm{Pb}^{2+}$ ions are precipitated out as $\mathrm{PbSO}_{4}(\mathrm{~s})$ at the negative electrode.

\mathrm{Pb}(\mathrm{~s})+\mathrm{SO}_{4}{ }^{2-}(\mathrm{aq}) \rightarrow \mathrm{PbSO}_{4}(\mathrm{~s})+2 \mathrm{e}^{-}

[ 1 ]

(i)

Complete the half-equation for the reaction taking place at the positive electrode.

\mathrm{PbO}_{2}(\mathrm{~s})+\mathrm{SO}_{4}{ }^{2-}(\mathrm{aq})+\ldots \ldots \ldots \ldots . .+\ldots \ldots \ldots \ldots . \rightarrow \mathrm{PbSO}_{4}(\mathrm{~s})+\ldots \ldots \ldots \ldots .

[ 1 ]

$\mathrm{PbO}_{2}(\mathrm{~s})+\mathrm{SO}_{4}{ }^{2-}(\mathrm{aq})+\mathbf{4 H}^{+}+\mathbf{2 e}^{-} \rightarrow \mathrm{PbSO}_{4}(\mathrm{~s})+\mathbf{2 H}_{2} \mathbf{O}$

(e)

The diagrams show how the voltage across two different cells changes with time when each cell is used to provide an electric current.

Figure for Question (e) — CAIE A-Level Chemistry A2

Suggest a reason why
- the voltage of the lead-acid cell changes after several hours,
- the voltage of the fuel cell remains constant.

[ 2 ]

reagents/ $\mathrm{PbO}_{2} / \mathrm{H}_{2} \mathrm{SO}_{4}$ and used up/concentration decreases

as fuel/hydrogen is being continuously supplied/fuel has not run out

[Maximum number: 14]

Some electrode potentials are shown in Table 3.1.

Table 3.1

(a)

Complete the diagram to show a standard hydrogen electrode.

Label your diagram. Identify all substances. You do not need to state standard conditions.

Figure for Question (a) — CAIE A-Level Chemistry A2

[ 1 ]

hydrogen, delivery system, $\mathrm{H}^{+}$ , platinum, [1]

(b)

An electrochemical cell is set up using an $\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}$ electrode and a standard hydrogen electrode.

Identify the positive electrode in the electrochemical cell and the direction of electron flow in the external circuit.
positive electrode
Electrons flow from the electrode to the electrode.

[ 1 ]

(c)

The vanadium-containing species in the electrode reactions given in Table 3.1 are $\mathrm{V}, \mathrm{V}^{2+}$ , $\mathrm{V}^{3+}, \mathrm{VO}^{2+}$ and $\mathrm{VO}_{2}{ }^{+}$ .

[ 3 ]

(i)

Identify one vanadium-containing species that does not react with $\mathrm{Fe}^{2+}$ ions under standard conditions.

Use data from Table 3.1 to explain your answer.

[ 1 ]

(for specified $\mathrm{V}^{2+}, \mathrm{V}^{3+}$ or $\mathrm{VO}^{2+}$ ) $\mathrm{E}^{\ominus}$ is more positive than / above -0.44 AND more negative than / below 0.77 V [1]

(ii)

Identify all the vanadium-containing species that will react with $\mathrm{Fe}^{2+}$ ions under standard conditions.

[ 1 ]

V and $\mathrm{VO}_{2}{ }^{+}$ [1]

(iii)

Write an equation for one of the possible reactions identified in (ii).

[ 1 ]

$\begin{aligned} \mathrm{V}+\mathrm{Fe}^{2+} \rightarrow \mathrm{V}^{2+}+\mathrm{Fe} \mathrm{VO}_{2}{ }^{+}+2 \mathrm{H}^{+}+\mathrm{Fe}^{2+} \rightarrow \mathrm{VO}^{2+}+\mathrm{H}_{2} \mathrm{O}+\mathrm{Fe}^{3+}[1] \end{aligned}$

(d)

Another electrochemical cell is set up using an $\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}$ electrode and an alkaline $\mathrm{ClO}^{-} / \mathrm{Cl}^{-}$ electrode.

The concentration of $\mathrm{Fe}^{3+}$ is 1000 times greater than the concentration of $\mathrm{Fe}^{2+}$ in the $\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}$ electrode. All other conditions are standard.

[ 3 ]

(i)

Use the Nernst equation to calculate the E value of the $\mathrm{Fe}^{3+} / \mathrm{Fe}^{2+}$ electrode.

Show your working.

[ 2 ]

Nernst: $\mathrm{E}=0.77+(0.059 / \mathrm{z}) \log [\mathrm{ox}] /[\mathrm{red}][1]$
0.947 [1]

(ii)

Write an equation for the reaction that occurs in the cell, under these conditions.

[ 1 ]

$2 \mathrm{Fe}^{3+}+\mathrm{Cl}^{-}+2 \mathrm{OH}^{-} \rightarrow 2 \mathrm{Fe}^{2+}+\mathrm{ClO}^{-}+\mathrm{H}_{2} \mathrm{O}[1]$

(e)

Another electrochemical cell is set up using an $\mathrm{Fe}^{2+} / \mathrm{Fe}$ electrode and an alkaline $\mathrm{ClO}^{-} / \mathrm{Cl}^{-}$ electrode under standard conditions.

Calculate the value of $\Delta G^{\ominus}$ for the cell.
$\Delta G^{\ominus}=$ $\mathrm{kJ} \mathrm{mol}^{-1}$

[ 3 ]

$\begin{aligned} \mathrm{E}^{\ominus}{ }_{\text {cell }}=1.33 \mathrm{~V}[1] \Delta \mathrm{G}^{\ominus}=-\mathrm{nE}^{\ominus}{ }_{\text {cell }} F[1] -257[1] \end{aligned}$

(f)

Iron(II) chloride, $\mathrm{FeCl}_{2}$ , is oxidised by chlorine to form iron(III) chloride, $\mathrm{FeCl}_{3}$ , under standard conditions.

2 \mathrm{FeCl}_{2}(\mathrm{~s})+\mathrm{Cl}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{FeCl}_{3}(\mathrm{~s})

\Delta H^{\ominus}=-128 \mathrm{~kJ} \mathrm{~mol}^{-1}

Table 3.2

[ 3 ]

(i)

Use Table 3.2 and other data to calculate the Gibbs free energy change, $\Delta G^{\ominus}$ , for this reaction.

Show your working.

\Delta G^{\ominus}=

\mathrm{kJ} \mathrm{~mol}^{-1}

[ 3 ]

$\Delta \mathrm{S}=-179[1]$ $\Delta \mathrm{G}=\Delta \mathrm{H}-\mathrm{T} \Delta \mathrm{S}[1]$

-74.7

[Maximum number: 8]

Hypophosphorous acid is an inorganic acid.
The conjugate base of hypophosphorous acid is $\mathrm{H}_{2} \mathrm{PO}_{2}^{-}$ .

(a)

$\mathrm{H}_{2} \mathrm{PO}_{2}^{-}$ is a strong reducing agent. It can be used to reduce metal cations without the need for electrolysis.
equation $1 \quad \mathrm{HPO}_{3}^{2-}+2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{PO}_{2}^{-}+3 \mathrm{OH}^{-}$ $E^{\ominus}=-1.57 \mathrm{~V}$

[ 8 ]

(i)

In an experiment, an alkaline $\mathrm{HPO}_{3}{ }^{2-} / \mathrm{H}_{2} \mathrm{PO}_{2}{ }^{-}$ half-cell is constructed with $\left[\mathrm{H}_{2} \mathrm{PO}_{2}^{-}\right]=0.050 \mathrm{~mol} \mathrm{dm}^{-3}$ .

All other ions are at their standard concentration.

Predict how the value of E of this half-cell differs from its $E^{\ominus}$ value.

Explain your answer.

[ 2 ]

electrode potential E would become more positive / less negative (than $E \ominus$ )

lower $\left[\mathrm{H}_{2} \mathrm{PO}_{2}{ }^{-}\right]$ AND shifts equilibrium to the right-hand side

(ii)

The $\mathrm{Cr}^{3+} / \mathrm{Cr}$ half-cell has a standard electrode potential of -0.74 V .

An electrochemical cell consists of an alkaline $\mathrm{HPO}_{3}{ }^{2-} / \mathrm{H}_{2} \mathrm{PO}_{2}{ }^{-}$ half-cell and a $\mathrm{Cr}^{3+} / \mathrm{Cr}$ half-cell.

Calculate the standard cell potential, $E_{\text {cell }}^{\ominus}$ .

E_{\text {cell }}^{\ominus}=

[ 1 ]

$+1.57-0.74=(+) 0.83(\mathrm{~V})$

(iii)

Complete the diagram in Fig. 2.1 to show how the standard electrode potential of the $\mathrm{Cr}^{3+} / \mathrm{Cr}$ half-cell can be measured relative to that of the standard hydrogen electrode.

Identify the chemicals, conditions and relevant pieces of apparatus.

Fig. 2.1

[ 3 ]

(iv)

Label Fig. 2.1 to show:
- which is the positive electrode
- the direction of electron flow in the external circuit.

[ 1 ]

(v)

$\mathrm{H}_{2} \mathrm{PO}_{2}^{-}$ reduces $\mathrm{Ni}^{2+}$ to Ni in alkaline conditions.

Use equation 1 to construct the ionic equation for this reaction.
equation $1 \quad \mathrm{HPO}_{3}{ }^{2-}+2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{PO}_{2}{ }^{-}+3 \mathrm{OH}^{-}$

[ 1 ]

$\mathrm{H}_{2} \mathrm{PO}_{2}{ }^{-}+3 \mathrm{OH}^{-}+\mathrm{Ni}^{2+} \rightarrow \mathrm{HPO}_{3}{ }^{2-}+2 \mathrm{H}_{2} \mathrm{O}+\mathrm{Ni}$

All question bank results loaded

Covered syllabus points

24.2 Standard electrode potentials E ⦵ , standard cell potentials E ⦵ cell and the Nernst question 1

Question (a)

Question (b)

Question (c)

Question (d)

Question (i)

Question (e)

24.2 Standard electrode potentials E ⦵ , standard cell potentials E ⦵ cell and the Nernst question 2

Question (a)

Question (b)

Question (c)

Question (i)

Question (ii)

Question (iii)

Question (d)

Question (i)

Question (ii)

Question (e)

Question (f)

Question (i)

24.2 Standard electrode potentials E ⦵ , standard cell potentials E ⦵ cell and the Nernst question 3

Question (a)

Question (i)

Question (ii)

Question (iii)

Question (iv)

Question (v)