(a)

$\mathrm{Co}^{2+}$ and $\mathrm{Co}^{3+}$ both form complexes with edta4-.

Use the data in the table to predict what happens, if anything, when separate aqueous solutions of $\mathrm{Co}^{3+}$ and $[\mathrm{Co}(\text { edta })]^{-}$ are left to stand in the air.
aqueous solution of $\mathrm{Co}^{3+}$
aqueous solution of $[\mathrm{Co}(\mathrm{edta})]^{-}$

[ 3 ]

(a)

Both $\mathrm{KMnO}_{4}$ and $\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}$ are used as oxidising agents, usually in acidic solution.

[ 4 ]

(i)

Use information from the Data Booklet to explain why their oxidising power increases as the $\left[\mathrm{H}^{+}(\mathrm{aq})\right]$ in the solution increases.

(ii)

What colour changes would you observe when each of these oxidising agents is completely reduced?
- $\mathrm{KMnO}_{4}$
from to
- $\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}$
from. to

[ 4 ]

(b)

Manganese(IV) oxide, $\mathrm{MnO}_{2}$ , is a dark brown solid, insoluble in water and dilute acids. Passing a stream of $\mathrm{SO}_{2}(\mathrm{~g})$ through a suspension of $\mathrm{MnO}_{2}$ in water does, however, cause it to dissolve, to give a colourless solution.

[ 4 ]

(i)

Use the Data Booklet to suggest an equation for this reaction, and explain what happens to the oxidation states of manganese and of sulfur during the reaction.

(ii)

The pH of the suspension of $\mathrm{MnO}_{2}$ is reduced. Explain what effect, if any, this would have on the extent of this reaction.

[ 4 ]

(c)

The main ore of manganese, pyrolusite, is mainly $\mathrm{MnO}_{2}$ . A solution of $\mathrm{SnCl}_{2}$ can be used to estimate the percentage of $\mathrm{MnO}_{2}$ in a sample of pyrolusite, using the following method.
- A known mass of pyrolusite is warmed with an acidified solution containing a known amount of $\mathrm{SnCl}_{2}$ .
- The excess $\mathrm{Sn}^{2+}(\mathrm{aq})$ ions are titrated with a standard solution of $\mathrm{KMnO}_{4}$ .

In one such experiment, 0.100 g of pyrolusite was warmed with an acidified solution containing $2.00 \times 10^{-3} \mathrm{molSn}^{2+}$ . After the reaction was complete, the mixture was titrated with $0.0200 \mathrm{moldm}^{-3} \mathrm{KMnO}_{4}$ , and required $18.1 \mathrm{~cm}^{3}$ of this solution to reach the end point.

The equation for the reaction between $\mathrm{Sn}^{2+}(\mathrm{aq})$ and $\mathrm{MnO}_{4}^{-}(\mathrm{aq})$ is as follows.

2 \mathrm{MnO}_{4}^{-}+5 \mathrm{Sn}^{2+}+16 \mathrm{H}^{+} \rightarrow 2 \mathrm{Mn}^{2+}+5 \mathrm{Sn}^{4+}+8 \mathrm{H}_{2} \mathrm{O}

(i)

Use the Data Booklet to construct an equation for the reaction between $\mathrm{MnO}_{2}$ and $\mathrm{Sn}^{2+}$ ions in acidic solution.

[Maximum number: 9]

$1 \mathrm{EDTA}^{4-}$ , is a polydentate ligand.

(a)

Cadmium ions form complexes with methylamine, $\mathrm{CH}_{3} \mathrm{NH}_{2}$ , and with 1,2-diaminoethane, $\mathrm{H}_{2} \mathrm{NCH}_{2} \mathrm{CH}_{2} \mathrm{NH}_{2}$ , as shown in equilibriums 2 and 3. 1,2-diaminoethane is shown as en.
equilibrium $2\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+4 \mathrm{CH}_{3} \mathrm{NH}_{2} \rightleftharpoons\left[\mathrm{Cd}\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad K_{\text {stab2 }}=3.60 \times 10^{6}$
equilibrium $3\left[\mathrm{Cd}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{2+}+2 \mathrm{en} \rightleftharpoons\left[\mathrm{Cd}(\mathrm{en})_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad \mathrm{K}_{\text {stab3 }}=4.20 \times 10^{10}$

An equilibrium is set up between these two complexes as shown in equilibrium 4.
equilibrium $4\left[\mathrm{Cd}\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+2 \mathrm{en} \rightleftharpoons\left[\mathrm{Cd}(\mathrm{en})_{2}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]^{2+}+4 \mathrm{CH}_{3} \mathrm{NH}_{2}$

\begin{aligned} & \Delta H^{\ominus}=+0.840 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ & \Delta S^{\ominus}=+80.9 \mathrm{JK}^{-1} \mathrm{~mol}^{-1} \end{aligned}

(i)

Calculate the value of the standard Gibbs free energy change, $\Delta G^{\ominus}$ , for equilibrium 4 at 298 K .
$\Delta G^{\ominus}=$ $\mathrm{kJ} \mathrm{mol}^{-1}$

(a)

Iron(III) chloride, $\mathrm{FeCl}_{3}$ , is used to dissolve unwanted copper from printed circuit boards (PCBs) by the following reaction.

2 \mathrm{FeCl}_{3}(\mathrm{aq})+\mathrm{Cu}(\mathrm{~s}) \rightarrow 2 \mathrm{FeCl}_{2}(\mathrm{aq})+\mathrm{CuCl}_{2}(\mathrm{aq})

A solution in which $\left[\mathrm{Fe}^{3+}(\mathrm{aq})\right]$ was originally equal to $1.50 \mathrm{~mol} \mathrm{dm}^{-3}$ was re-used several times to dissolve copper from the PCBs, and was then titrated as follows.

A $2.50 \mathrm{~cm}^{3}$ sample of the partially-used-up solution was acidified and titrated with $0.0200 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{KMnO}_{4}$ .
This oxidised any $\mathrm{FeCl}_{2}$ in the solution back to $\mathrm{FeCl}_{3}$ .
It was found that $15.0 \mathrm{~cm}^{3}$ of $\mathrm{KMnO}_{4}(\mathrm{aq})$ was required to reach the end point.

[ 6 ]

(i)

Construct an ionic equation for the reaction between $\mathrm{Fe}^{2+}$ and $\mathrm{MnO}_{4}^{-}$ in acid solution.

(ii)

State here the $\mathrm{Fe}^{2+}: \mathrm{MnO}_{4}^{-}$ ratio from your equation in (i).

(iii)

Calculate the mass of copper that could still be dissolved by $100 \mathrm{~cm}^{3}$ of the partially-used-up solution.
mass of copper =

[ 6 ]

(a)

The $\left[\mathrm{Cr}_{2}\left(\mathrm{O}_{2} \mathrm{CCH}_{3}\right)_{4}\left(\mathrm{H}_{2} \mathrm{O}\right)_{2}\right]$ complex reacts with aqueous acid to form $\mathrm{Cr}^{2+}(\mathrm{aq})$ ions. $\mathrm{Cr}^{2+}(\mathrm{aq})$ ions react with $\mathrm{O}_{2}(\mathrm{aq})$ under acidic conditions. $\mathrm{Cr}^{3+}(\mathrm{aq})$ ions are formed.

Use the Data Booklet to answer the following questions.

[ 3 ]

(i)

Construct an ionic equation for the reaction of $\mathrm{Cr}^{2+}(\mathrm{aq})$ with $\mathrm{O}_{2}(\mathrm{aq})$ under acidic conditions.

[ 2 ]

(ii)

Calculate $E_{\text {cell }}^{\ominus}$ for the reaction in (e)(i).

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

V

[ 1 ]

(a)

When aqueous solutions of KI and $\mathrm{K}_{2} \mathrm{~S}_{2} \mathrm{O}_{8}$ are mixed almost no reaction occurs, but when a few drops of $\mathrm{Fe}^{2+}(\mathrm{aq})$ or $\mathrm{Fe}^{3+}(\mathrm{aq})$ are added, iodine, $\mathrm{I}_{2}(\mathrm{aq})$ , is produced at a steady rate.

(i)

Write an equation for the overall reaction.

(a)

The following diagram shows the apparatus used to measure the standard electrode potential, $E^{\ominus}$ , of a cell composed of a Cu(II) / Cu electrode and an Fe(II) / Fe electrode.

[ 8 ]

(i)

Finish the diagram by adding components to show the complete circuit. Label the components you add.

(ii)

In the spaces below, identify or describe what the four letters A-D represent.

A

B

C

D

(iii)

Use the Data Booklet to calculate the $E^{\ominus}$ for this cell.

(iv)

Predict how the size of the overall cell potential would change, if at all, as the concentration of solution C is increased.
Explain your reasoning.

[ 8 ]

[Maximum number: 6]

An electrochemical cell is constructed using two half-cells.
- an $\mathrm{Sn}^{4+} / \mathrm{Sn}^{2+}$ half-cell
- an $\mathrm{A} l^{3+} / \mathrm{A} l$ half-cell

(a)

State the material used for the electrode in each half-cell.
- $\mathrm{Sn}^{4+} / \mathrm{Sn}^{2+}$ half-cell
- $\mathrm{A} l^{3+} / \mathrm{A} l$ half-cell

[ 1 ]

(b)

The cell is operated at 298 K .

The $\mathrm{A} l^{3+} / \mathrm{A} l$ half-cell has standard concentrations.
The $\mathrm{Sn}^{4+} / \mathrm{Sn}^{2+}$ half-cell has $\left[\mathrm{Sn}^{4+}\right]=0.300 \mathrm{~mol} \mathrm{dm}^{-3}$ and $\left[\mathrm{Sn}^{2+}\right]=0.150 \mathrm{~mol} \mathrm{dm}^{-3}$ .

[ 5 ]

(i)

Use the Nernst equation to calculate the electrode potential, E, of the $\mathrm{Sn}^{4+} / \mathrm{Sn}^{2+}$ half-cell under these conditions.

E =V

[ 2 ]

(ii)

Calculate the $E_{\text {cell }}$ under these conditions.

E_{\text {cell }}=

V

[ 1 ]

(iii)

Write an equation for the overall cell reaction that occurs.

[ 2 ]

(a)

(i)

Use $E^{\ominus}$ values from the Data Booklet to predict the relative oxidising abilities of fluorine and chlorine.

[ 2 ]

(ii)

Predict the type of reaction that would occur between the interhalogen compound chlorine fluoride, ClF, and potassium bromide solution.

[ 1 ]

(iii)

Construct an equation for this reaction.

[ 1 ]

[Maximum number: 6]

An electrochemical cell is constructed using two half-cells.
- a $\mathrm{Br}_{2} / \mathrm{Br}^{-}$ half-cell
- an $\mathrm{Mn}^{3+} / \mathrm{Mn}^{2+}$ half-cell

(a)

State the material used for the electrode in each half-cell.
$\mathrm{Br}_{2} / \mathrm{Br}^{-}$ half-cell
$\mathrm{Mn}^{3+} / \mathrm{Mn}^{2+}$ half-cell

[ 1 ]

(b)

The cell is operated at 298 K .

The $\mathrm{Br}_{2} / \mathrm{Br}^{-}$ half-cell has standard concentrations.
The $\mathrm{Mn}^{3+} / \mathrm{Mn}^{2+}$ half-cell has $\left[\mathrm{Mn}^{3+}\right]=0.500 \mathrm{moldm}{ }^{-3}$ and $\left[\mathrm{Mn}^{2+}\right]=0.100 \mathrm{moldm}{ }^{-3}$ .

[ 5 ]

(i)

Use the Nernst equation to calculate the electrode potential, E, of the $\mathrm{Mn}^{3+} / \mathrm{Mn}^{2+}$ half-cell under these conditions.

E=

[ 2 ]

(ii)

Calculate the $E_{\text {cell }}$ under these conditions.

E_{\text {cell }}=

[ 1 ]

(iii)

Write an equation for the overall cell reaction that occurs.

[ 2 ]