AS & A Level Chemistry - 9701 Paper 4 2023 Summer Zone 2

(a) Group 2 carbonates decompose when heated to form the metal oxide and carbon dioxide.
(i) Suggest a mechanism for the decomposition of the carbonate ion by adding two curly arrows in Fig. 1.1.
!(CO2 + O2-)
Fig. 1.1 [1]
(ii) Describe the variation in the thermal stability of Group 2 carbonates. Explain your answer.
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(b) (i) Define lattice energy.
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(ii) The lattice energy of the Group 2 carbonates, $\Delta H_{\text{latt}}^{\circ}(MCO_3)$, becomes less exothermic down the group. The lattice energy of the Group 2 oxides, $\Delta H_{\text{latt}}^{\circ}(MO)$, also becomes less exothermic down the group. $\Delta H_{\text{latt}}^{\circ}(MCO_3)$ and $\Delta H_{\text{latt}}^{\circ}(MO)$ change by different amounts going down the group.
Suggest how the standard enthalpy change of the decomposition reaction for Group 2 carbonates changes down the group.
Explain your reasoning in terms of the relative sizes of the anions and the relative changes in lattice energy down the group.
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(c) Potassium sulfite, $K_2SO_3$, is used as a food additive.
The concentration of sulfite ions, $SO_3^{2-}$, can be determined by titration using aqueous acidified manganate(VII) ions, $MnO_4^-$.

• A 250 cm³ solution contains 3.40 g of impure $K_2SO_3$.
• 25.0 cm³ of this solution requires 22.40 cm³ of 0.0250 mol dm⁻³ acidified $MnO_4^-$ to reach the end-point. All the $SO_3^{2-}$ ions are oxidised. None of the other species in the impure $K_2SO_3$ are oxidised.

(a) State two typical chemical properties of a transition element.
1 ................................................................................................................................................................
2 ................................................................................................................................................................ [1]

(b) Aqueous solutions of cobalt(II) salts contain the complex ion $\text{[Co(H}_2\text{O)}_6\text{]}^{2+}$.

(i) Define complex ion.
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(ii) Samples of $\text{[Co(H}_2\text{O)}_6\text{]}^{2+}$ are reacted separately with an excess of aqueous ammonia, with an excess of concentrated HCl and with an excess of aqueous sodium hydroxide, as shown in Fig. 2.1.

$$\text{}$$

Complete Table 2.1 about the reactions shown by $\text{[Co(H}_2\text{O)}_6\text{]}^{2+}$.

$$\text{[Table_1]}$$

reagent added to $\text{[Co(H}_2\text{O)}_6\text{]}^{2+}$ (aq) | formula of cobalt species formed | colour and state of cobalt species formed | type of reaction

an excess of NH$_3$(aq): A = ...........................................................................................................................................
an excess of concentrated HCl: B = ..........................................................................................................................
an excess of NaOH(aq): C = ...................................................................................................................................... [4]

(c) The ethanedioate ion, $\text{C}_2\text{O}_4^{2-}$, can act as a bidentate ligand.

(i) Explain what is meant by a bidentate ligand.
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(ii) The complex $\text{[Co(H}_2\text{O)}_2\text{(C}_2\text{O}_4\text{)}\text{BrCl]}^{-}$ exists as stereoisomers. Complete the three-dimensional diagrams in Fig. 2.2 to show four stereoisomers of $\text{[Co(H}_2\text{O)}_2\text{(C}_2\text{O}_4\text{)}\text{BrCl]}^{-}$.
The $\text{C}_2\text{O}_4$ ligand is represented using $$\text{O ◌O}$$.
$$\text{}$$ [3]

(iii) State the oxidation state of cobalt in this complex and a type of stereoisomerism shown.
oxidation state of cobalt ....................................................................................................................................
type of stereoisomerism ............................................................................................................................... [1]

(a) Complete Table 3.1 by placing one tick (✓) in each row to indicate the sign of each type of
energy change under standard conditions.

Table 3.1
[Table_1]

(b) Define standard enthalpy change of atomisation.
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(c) Table 3.2 shows some energy changes.

Table 3.2
[Table_2]

Calculate the lattice energy, $\Delta H_{\text{latt}}^{\circ}$, of $\text{Ag}_2\text{O(s)}$ using relevant data from Table 3.2.
It may be helpful to draw a labelled energy cycle.
Show your working.

$\Delta H_{\text{latt}}^{\circ}$ of $\text{Ag}_2\text{O(s)}$ = ........................................ kJ mol\(^{-1}\)

(d) Suggest the trend in the magnitude of the lattice energies of the silver compounds \(\text{Ag}_2\text{S}, \\\,\text{Ag}_2\text{O and Ag}_2\text{Se}\).
Explain your answer.
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least exothermic most exothermic
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(e) Silver sulfite, $\text{Ag}_2\text{SO}_3\text{(s)}$, is sparingly soluble in water.

(i) Give an expression for the solubility product, $K_{sp}$, of $\text{Ag}_2\text{SO}_3$.
$K_{sp}$ = .................................................................................................................................
[1]

(ii) Calculate the equilibrium concentration of \(\text{Ag}^+\) in a saturated solution of $\text{Ag}_2\text{SO}_3$ at 298 K.
[\(K_{sp}\): $\text{Ag}_2\text{SO}_3$, $1.50 \times 10^{-14} \text{mol}^3 \text{dm}^{-9}$ at 298 K]
[Ag\(^+\)] = ........................................... mol dm\(^{-3}\)

(f) The standard enthalpy change of solution, $\Delta H_{\text{sol}}^{\circ}$, of $\text{AgNO}_3\text{(s)}$ in water is +22.6 kJ mol\(^{-1}\).
Suggest how the feasibility of dissolving $\text{AgNO}_3\text{(s)}$ in water changes with temperature.
Explain your answer.
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(a) In aqueous solution, iron(III) ions react with iodide ions, as shown.

$$2\text{Fe}^{3+} + 2\text{I}^{-} \rightarrow 2\text{Fe}^{2+} + \text{I}_2$$
A series of experiments is carried out using different concentrations of $\text{Fe}^{3+}$ and $\text{I}^{-}$, as shown in Table 4.1.

$$\begin{array}{|c|c|c|c|} \hline \text{experiment} & [\text{Fe}^{3+}]/\text{mol dm}^{-3} & [\text{I}^{-}]/\text{mol dm}^{-3} & \text{initial rate}/\text{mol dm}^{-3}\text{s}^{-1} \\ \hline 1 & 0.0400 & 0.0200 & 2.64 \times 10^{-4} \\ 2 & 0.1200 & 0.0200 & 7.92 \times 10^{-4} \\ 3 & 0.0800 & 0.0400 & 2.11 \times 10^{-3} \\ \hline \end{array}$$

(i) Explain what is meant by overall order of reaction.
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(ii) Use the data in Table 4.1 to deduce the order of reaction with respect to $\text{Fe}^{3+}$ and with respect to $\text{I}^{-}$.
Explain your reasoning.
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(iii) Use your answer to (a)(ii) to construct the rate equation for this reaction.
rate = ..................................................................................................................................................... [1]

(iv) Use your answer to (a)(iii) and the data from experiment 1 to calculate the rate constant, $k$, for this reaction. Include the units of $k$.
$$k = \text{...............................} \text{ units ...............................}$$ [2]

(v) Describe qualitatively the effect of an increase in temperature on the rate constant and on the rate of this reaction.
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(b) In aqueous solution, iodide ions react with acidified hydrogen peroxide, as shown.

$$2\text{I}^{-} + \text{H}_2\text{O}_2 + 2\text{H}^{+} \rightarrow \text{I}_2 + 2\text{H}_2\text{O}$$
The initial rate of reaction is found to be first order with respect to $\text{I}^{-}$, first order with respect to $\text{H}_2\text{O}_2$ and zero order with respect to $\text{H}^{+}$.

Fig. 4.1 shows a possible four-step mechanism for this reaction.

$$\text{step 1} \quad \text{H}_2\text{O}_2 + \text{I}^{-} \rightarrow \text{IO}^{-} + \text{H}_2\text{O}$$
$$\text{step 2} \quad \text{H}^{+} + \text{IO}^{-} \rightarrow \text{HIO}$$
$$\text{step 3} \quad \text{HIO} + \text{I}^{-} \rightarrow \text{I}_2 + \text{OH}^{-}$$
$$\text{step 4} \quad \text{OH}^{-} + \text{H}^{+} \rightarrow \text{H}_2\text{O}$$

Fig. 4.1

(i) Suggest which of the steps, 1, 2, 3 or 4, in this mechanism is the rate-determining step.
Explain your answer.
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(ii) Identify a step in Fig. 4.1 that involves a redox reaction.
Explain your answer in terms of oxidation numbers.
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(iii) Suggest the role of HIO in this mechanism.
Explain your reasoning.
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(a) Methylbenzene can undergo different reactions, as shown in Fig. 5.1.

(i) Draw structures in Fig. 5.1 for the possible organic products of the three reactions shown. [3]
(ii) Complete Table 5.1.
[Table_1]
[2]
(b) When methylbenzene reacts with an electrophile, a substitution reaction occurs. No addition reaction takes place under these conditions.
Explain why no addition reaction takes place.
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(c) The reaction of methylbenzene with thionyl bromide, SOBr$_2$, in the presence of an iron(III) bromide catalyst, FeBr$_3$, is shown in Fig. 5.2.

The mechanism of this reaction is similar to that of the bromination of benzene.
The first step of the mechanism generates the SOBr$^+$ electrophile, as shown.
$ ext{SOBr}_2 + ext{FeBr}_3 ightarrow ext{SOBr}^+ + ext{FeBr}_4^-$
(i) The reaction of methylbenzene with SOBr$^+$ ions is shown in Fig. 5.3. Complete the mechanism in Fig. 5.3.
Include all relevant curly arrows and charges. Draw the structure of the organic intermediate.

[3]
(ii) The reaction shown in Fig. 5.2 produces a small amount of a by-product, P, with the molecular formula C$_{14}$H$_{14}$OS.
Suggest a structure for by-product P.
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[1]

(d) Acyl bromides, RCOBr, can be synthesised by the reaction of a carboxylic acid and SOBr$_2$.
This is a similar reaction to the synthesis of acyl chlorides using SOCl$_2$.
(i) Give an equation for the reaction between ethanoic acid and SOBr$_2$.
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[1]
(ii) Suggest the relative ease of hydrolysis of acyl bromides, RCOBr, acyl chlorides, RCOCl, and alkyl chlorides, RCl.
Explain your answer.
.............................................. > .............................................. > ..............................................
easiest to hydrolyse              hardest to hydrolyse
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[3]

(a) Perindopril is a drug used to treat heart disease.

[Image_1: perindopril structure]

(i) State the number of chiral carbon atoms present in one molecule of perindopril.
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(ii) Suggest one benefit and one disadvantage of producing a drug such as perindopril as a single pure optical isomer.
benefit ......................................................................................................................................................
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disadvantage ...............................................................................................................................................
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(b) (i) Name all the functional groups in perindopril.
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(ii) A sample of perindopril is hydrolysed with hot aqueous acid. Draw the structures of the three organic products of the complete acid hydrolysis of perindopril.
[Image_2: Three boxes for drawing structures] [3]

(a) Explain why phenol is brominated much more easily than benzene is brominated.
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(b) Iodine monobromide, $\text{I--Br}$, reacts with benzene in the presence of an $\text{AlBr}_3$ catalyst.
Predict whether the organic product will be bromobenzene or iodobenzene.
Explain your answer.
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(c) Fig. 7.1 shows some reactions of phenol.


(i) Give an equation for the reaction of phenol with $\text{Na(s)}$.
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(ii) Draw the structure of the organic product, $R$, formed when phenol reacts with an excess of $\text{Br}_2(\text{aq})$.
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(iii) State the reagents and conditions for reaction 1 and reaction 2 in Fig. 7.1.

reaction 1 ..........................................................................................................

reaction 2 ..........................................................................................................

(a) Describe the difference in reactivity between chloroethane and chlorobenzene with OH⁻(aq).

Explain your answer.

(b) Compound T, C₅H₉O₂Cl, is a useful synthetic intermediate.
Fig. 8.1 shows some reactions of T.



(i) Give the systematic name for T.
[1]

(ii) Draw the structures of W, X, Y and Z in Fig. 8.1.
[4]

(iii) State the reagents and conditions for steps 1 and 2 in Fig. 8.1.
step 1 ........................................................ [2]
step 2 ........................................................

(c) The proton ($^1$H) NMR spectrum of compound T, C₅H₉O₂Cl, in CDCl₃ is shown in Fig. 8.2.



Table 8.1

[Table_1]

(i) Suggest why CDCl₃ is used as a solvent instead of CHCl₃ for the proton ($^1$H) NMR spectrum.
[1]

(ii) Complete Table 8.2 for the proton ($^1$H) NMR spectrum of T.

Table 8.2

[Table_2]

[4]

(iii) Explain the splitting pattern of the peak at δ 3.9 ppm.
[1]

(a) Define standard cell potential, $E_{\text{cell}}^{\circ}$.
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(b) An electrochemical cell is set up to measure $E_{\text{cell}}^{\circ}$ of a cell consisting of an $\text{Fe}^{3+}/\text{Fe}^{2+}$ half-cell and a $\text{Cl}_2/\text{Cl}^-$ half-cell.

Draw a labelled diagram of this electrochemical cell.
Include all necessary substances. It is not necessary to state conditions used. [3]

(c) The cell reaction for the electrochemical cell in (b) is shown.

$$\text{Cl}_2 + 2\text{Fe}^{2+} \rightarrow 2\text{Fe}^{3+} + 2\text{Cl}^- \qquad E_{\text{cell}}^{\circ} = +0.59\text{ V}$$

Calculate $\Delta G^{\circ}$, in kJ mol$^{-1}$, for this cell reaction.

$\Delta G^{\circ} = \text{.........................} \text{ kJ mol}^{-1}$ [2]

(d) Another experiment is set up using the same electrochemical cell.
In this experiment the $\text{Fe}^{2+}$ concentration is 0.15 mol dm$^{-3}$. All other concentrations remain at their standard values.
The Nernst equation is shown.
$$E = E^{\circ} + (0.059/z) \log \left(\frac{\text{[oxidised species]}}{\text{[reduced species]}}\right)$$
(i) Use the Nernst equation to calculate the electrode potential, $E$, for the $\text{Fe}^{3+}/\text{Fe}^{2+}$ half-cell in this experiment.
$$[E^{\circ}:\ \text{Fe}^{3+}/\text{Fe}^{2+} = +0.77 \text{ V}]$$
$E = \text{.........................} \text{ V}$ [1]
(ii) Use your answer to (d)(i) to calculate $E_{\text{cell}}$ for this electrochemical cell.
$E_{\text{cell}} = \text{.........................} \text{ V}$ [1]
[Total: 8]