(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 ]

[Maximum number: 3]

Iodine is found naturally in compounds in many different oxidation states.

(a)

The concentration of $\mathrm{Cu}^{2+}(\mathrm{aq})$ in a solution can be determined by the reaction of $\mathrm{Cu}^{2+}$ ions with $\mathrm{I}^{-}$ ions.

reaction $12 \mathrm{Cu}^{2+}+4 \mathrm{I}^{-} \rightarrow 2 \mathrm{CuI}+\mathrm{I}_{2}$

The $\mathrm{I}_{2}$ produced in reaction 1 is titrated against a solution containing thiosulfate ions, $\mathrm{S}_{2} \mathrm{O}_{3}{ }^{2-}$ , using a suitable indicator.

\text { reaction } 2 \quad 2 \mathrm{~S}_{2} \mathrm{O}_{3}^{2-}+\mathrm{I}_{2} \rightarrow \mathrm{~S}_{4} \mathrm{O}_{6}^{2-}+2 \mathrm{I}^{-}

[ 1 ]

(i)

A $25.0 \mathrm{~cm}^{3}$ portion of a $\mathrm{Cu}^{2+}(\mathrm{aq})$ solution reacts with an excess of $\mathrm{I}^{-}(\mathrm{aq})$ .

The end-point of the titration occurs when $22.30 \mathrm{~cm}^{3}$ of $0.150 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{~S}_{2} \mathrm{O}_{3}{ }^{2-}(\mathrm{aq})$ is added.
Calculate the concentration of $\mathrm{Cu}^{2+}(\mathrm{aq})$ in the original solution.

(ii)

Identify a suitable indicator for the titration.

[ 1 ]

(b)

The reaction of $\mathrm{I}^{-}$ ions with persulfate ions, $\mathrm{S}_{2} \mathrm{O}_{8}{ }^{2-}$ , can be catalysed by $\mathrm{Fe}^{3+}$ ions.

2 \mathrm{I}^{-}+\mathrm{S}_{2} \mathrm{O}_{8}^{2-} \rightarrow \mathrm{I}_{2}+2 \mathrm{SO}_{4}^{2-}

Write equations to show how $\mathrm{Fe}^{3+}$ catalyses this reaction.

[ 2 ]

(a)

Define the term standard electrode potential, $E^{\ominus}$ .

[ 1 ]

(b)

(i)

Complete and label the diagram to show how the standard electrode potential, $E^{\ominus}$ , of $\mathrm{Ag}^{+}(\mathrm{aq}) / \mathrm{Ag}(\mathrm{s})$ could be measured under standard conditions.

[ 4 ]

(ii)

Use the Data Booklet to label the diagram in (e)(i) to show
- which is the positive electrode,
- the direction of electron flow in the external circuit when a current flows.

[ 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 ]

(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.
- 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)

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 ]

(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)

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)

$\mathrm{Co}^{2+}$ 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 ]

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