科目：化学 Chemistry (9701)
年份：2023 夏季
试卷：Paper 42
总页数：22
章节：19 Lattice energy, 20 Electrochemistry, 21 Further aspects of equilibria, 22 Reaction kinetics, 24 Transition elements, 25 Benzene and its compounds, 26 Carboxylic acids and their derivatives, 29 Organic synthesis, 30 Analytical chemistry
1. (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.
(ii) Describe the variation in the thermal stability of Group 2 carbonates. Explain your answer.
[3]
(b) (i) Define lattice energy.
[2]
(ii) The lattice energy of the Group 2 carbonates, $\Delta H_{\text{latt}}^{\ominus}(\text{MCO}_3)$, becomes less exothermic down the group.
The lattice energy of the Group 2 oxides, $\Delta H_{\text{latt }}^{\ominus}(\text{MO})$, also becomes less exothermic down the group.
$\Delta H_{\text{latt }}^{\circ}(\text{MCO}_3)$ and $\Delta H_{\text{latt }}^\circ(\text{MO})$ change by different amounts going down the group.
Suggest how the(c) Potassium sulfite, K$_2$SO$_3$, 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 $250\text{cm}^3$ solution contains 3.40 g of impure K$_2$SO$_3$.
• $25.0\text{cm}^3$ of this solution requires 22.40 cm$^3$ of $0.0250\text{mol dm}^{-3}$ 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$_2$SO$_3$-are oxidised.
The reaction occurs as shown by the two half-equations.
$$\text{H}_2\text{O} + \text{SO}_3^{2-} \rightarrow \text{SO}_4^{2-} + 2\text{H}^+ + 2\text{e}^-$$
$$\text{MnO}_4^{-} + 8\text{H}^+ + 5\text{e}^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}$$
(i) Give the ionic equation for the2 (a) State two typical chemical properties of a transition element.
1 ... [1]
(b) Aqueous solutions of cobalt(II) salts contain the complex ion $[\text{Co}(\text{H}_2\text{O})_6]^{2+}$.
(i) Define complex ion.
[1]
(ii) Samples of $[\text{Co}(\text{H}_{2}\text{O})_{6}]^{2+}$ 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.
Complete Table 2.1 about the reactions shown by $[\text{Co}(\text{H}_{\text{2}}\text{O})_{6}]^{2 +}$.
Table 2.1
| reagent added to $[\text{Co}(\text{H_{2}O})_{6}]^{2+}$ (aq) | formula of cobalt species formed | colour and state of cobalt species formed | type of reaction |
| :--- | :---: | :---: | :---: |
| an excess of NH$_3$ (aq)(c) The ethanedioate ion, $C_2O_4^{2-}$, can act as a bidentate ligand.
(i) Explain what is meant by a bidentate ligand.
[2]
(ii) The complex $[\text{Co(H}_2\text{O})_2(\text{C}_2\text{O}_4)\text{BrCl}]^-$ exists as stereoisomers.
Complete the three-dimensional diagrams in Fig. 2.2 to show four stereoisomers of $[\text{Co(H}_2\text {O})_2(\text{C}_ {2}\text{O}_4)\text{BrCl}]$.
The $C_2O_4^{2−}$ ligand is represented using $\textcircled{1}$$\textcircled{2}$$\textcircled{3}$$\textcircled{4}$.
Fig. 2.2
[3]
(iii) State the oxidation state of cobalt in this complex and a type of stereoisomerism shown.
oxidation state of cobalt .....
type of stereoisomerism ..... [1]
[Total: 12]3 (a) Complete Table 3.1 by placing one tick ($\checkmark$) in each row to indicate the sign of each type of energy change under standard conditions.
Table 3.1
| energy change | always positive | always negative | can be either negative or positive |
| :---: | :---: | :---: | :---
| bond energy | | | |
| enthalpy change of atomisation | | | |
| enthalpy changeof formation | | | |
[1]
(b) Define standard enthalpy change of atomisation.
[2]
(c) Table 3.2 shows some energy changes.
Table 3.2
| energy change | value/kJ mol$^{-1}$ |
| :---: | :---: |
| standard enthalpy change of atomisation of silver | +285 |
| first ionisation energy of silver | +731 |
| second ionisation energy of silver | +2074 |
| bond energy of O=O | +496 |
| bond energy of O–O | +150 |
| first electron affinity oCalculate the lattice energy, $\Delta H_{\text{latt}}^{\ominus}$ of Ag$_2$O(s) using relevant data from Table 3.2.
It may be helpful to draw a labelled energy cycle.
Show your working.
$$\Delta H_{\text{latt}}^{\circ} \text{ of } \text{Ag}_2\text{O}(s) = \dots\dots\dots\dots\quad \text{kJ mol}^{-1}$$
(d) Suggest the trend in the magnitude of the lattice energies of the silver compounds Ag$_2$S, Ag$_2$O and Ag$_2$Se.
Explain your answer.
| least exothermic | most exothermic |
| :--- | :--- |
| | |
[2](e) Silver sulfite, Ag$_2$SO$_3$(s), is sparingly soluble in water.
(i) Give an expression for the solubility product, $K_{\text{sp}}$, of Ag$_2$SO$_3$.
$$K_{\text{sp}} =$$
[1]
(ii) Calculate the equilibrium concentration of Ag$^+$ in a saturated solution of Ag$_2$SO$_3$ at 298 K.
$$[K_{\text{sp}}: \text{Ag}_2\text{SO}_3, 1.50 \times 10^{-14} \text{mol}^3\text{dm}^{-9}\text{at } 298\text{K}]$$
$$[\text{Ag}^+] = \dots\dots\dots\dots\mathrm{mol}\,\mathrm{dm}^{-3}$$
(f) The standard enthalpy change of solution, $\Delta H^{\ominus}_{\text{sol}}$, of AgNO$_3$(s) in water is +22.6 kJ mol$^{-1}$.
Suggest how the feasibility of dissolving AgNO$_3$(s) in water changes with temperature.
Explain your answer.
[Total: 12]4 (a) In aqueous solution, iron(III) ions react with iodide ions, as shown.
$$2\text{Fe}^{3+} + 2\text{I}^- \rightarrow 2\text{Fe}^{2+} + \text{I}_2$$
A series of experiments is carried out using different concentrations of Fe$^{3+}$ and I$^{-}$, as shown in Table 4.1.
Table 4.1
| experiment | $[\text{Fe}^{3+}]/\text{mol dm}^{-3}$ | $[\text{I}^-]/\text{mol dm}^{-3}$$\text{s}^{-1}$ |
| :--- | :---: | :---: |
| 1 | 0.0400 | 0.0200 | $2.64 \times 10^{-4}$ |
| 2 | 0.1200 | 0.0209 | $7.92 \times 10^{-4}$ |
tr>
| 3 | 0.0800 | 0.0400 | $2.11 \times 10^{-3}$ |
(i) Explain what is meant by overall order of reaction.
[1]
(ii) Use the data in Table 4.1 to deduce the order of reaction with respect to Fe$^{3+}$ and with respect to I$^{-}$.
Explain your reasoning.
[2]
(iii) Use your answer to (b) In aqueous solution, iodide ions react with acidified hydrogen peroxide, as shown. $$2I^- + H_2O_2 + 2H^+ \rightarrow I^-, first order with respect to H_2O_2 and zero order with respect to H^+.$$ Fig. 4.1 shows a possible four-step mechanism for this reaction. step 1 $H_2O_2 + I^- \rightarrow IO^- + H_2O$ step 2 $H^+ + IO^- \rightarrow HIO$ step 3 $HIO + I^- \rightarrow I_2 + OH^-$ step 4 $OH^- + H^+ \rightarrow H_2O$ 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. [1] (ii) Identify a step in Fig. 4.1 that involves a redox reaction. Explain your answer in terms of oxidation numbers. [1] (iii) Suggest the role of HIO in this mechanism. Explain your reasoning. [Total: 10]
5 (a) Methylbenzene can undergo different reactions, as shown in Fig. 5.1. Fig. 5.1 (i) Draw structures in Fig. 5.1 for the possible organic products of the three reactions shown. (ii) Complete Table 5.1. Table 5.1 | type of reaction | mechanism | | :--- | :--- | | reaction 1 | | | reaction 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. [1]
(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.
$$\text{SOBr}_2 + \text{FeBr}_3 \rightarrow \text{SOBr}^+ + \text{FeBr}_4^-$$
(i) The reaction of methylbenzenewith 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.
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.
$$\text{SOBr}_2+\text{FeBr}_3\rightarrow\text{SOBr}^++\text{FeBr}_4^-$$(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.
[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$_2$.
[1]
(ii) Suggest the relative ease of hydrolysis of acyl bromides, RCOBr, acyl chlorides, RCOCl,$\text{and}$ alkyl chlorides, RC$\text{l}$.
Explain your answer.
easiest to hydrolyse > hardest to hydrolyse
[3]
[Total: 14]6 (a) Perindopril is a drug used to treat heart disease. Fig. 6.1 (i) State the number of chiral carbon atoms present in one molecule of perindopril. [1] (ii) Suggest **one** benefit and **one** disadvantage of producing a drug such as perindopril as a single pure optical isomer. benefit ... disadvantage ... [2] (b) (i) Name all the functional groups in perindopril. [2]
(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.
7 (a) Explain why phenol is brominated much more easily than benzene is brominated.
[3]
(b) Iodine monobromide, I–Br, reacts with benzene in the presence of an $AlBr_3$ catalyst.
Predict whether the organic product will be bromobenzene or iodobenzene.
Explain your answer.
[1]
(c) Fig. 7.1 shows some reactions of phenol.
$$\text{Na(s)} \xrightarrow{\text{OH}} \text{excess Br}_2(\text{aq}) \quad R$$
reaction 1
reaction 2
Fig. 7.1
(i) Give an equation for the reaction of phenol with Na(s).
[1](ii) Draw the structure of the organic product, R, formed when phenol reacts with an excess of Br$_2$(aq). [1] (iii) State the reagents and conditions for reaction 1 and reaction 2 in Fig. 7.1. reaction 1 ..... [2] reaction 2 ..... [Total: 8]
8 (a) Describe the difference in reactivity between chloroethane and chlorobenzene with OH⁻(aq). Explain your answer. [2] (b) Compound T, C₅H₉O₂Cl, is a useful synthetic intermediate. Fig. 8.1 shows some reactions of T. Fig. 8.1
(i) Give the systematic name for T. (ii) Draw the structures of W, X, Y and Z in Fig. 8.1. (iii) State the reagents and conditions for steps 1 and 2 in Fig. 8.1. step 1 ... step 2 ...
(c) The proton (1H) NMR spectrum of compound T, C$_5$H$_9$O$_2$Cl, in CDCl$_3$ is shown in Fig. 8.2. Fig. 8.2 Table 8.1 | environment of proton | example | chemical shift range $\delta$/ppm | | :--- | :--- | :--- | | alkane | $-CH_3$, $-CH_2-$, $>CH-$ | 0.9–1.7 | | alkyl next to C=O | CH$_3$-C=O, $-CH_2$-C=O, $>CH-C=O$ | 2.2–3.0 | | alkyl next to aromatic ring | CH$_3$-Ar, $-CH_2$-Ar, $>CH-Ar$ | 2.3–3.0 | | alkyl next到 electronegative atom | CH$_3$-O, $-CH_2-O$, $-CH_2-Cl$ | 3.2–4.0 | | attached to alkene | =CHR | 4.5–6.0 | | attached to aromatic ring | H–Ar | 6.0–9.0 | | aldehyde | HCOR | 9.3–10.5 | | alcohol | ROH | 0.5–6.0 | | phenol | Ar–OH | 4.5–7.0 | | carboxylic acid | RCOOH | 9.0–13.0 | | alkyl amine | R-NH- | 1.0–5.0 | | aryl amine | Ar-NH$_2$ | 3.0–6.0 | | amide | RCONHR | 5.0
(i) Suggest why CDCl$_3$ is used as a solvent instead of CHCl$_3$ for the proton ($^1\text{H}$) NMR spectrum.
(ii) Complete Table 8.2 for the proton ($^1\text{ H}$) NMR spectrum of T.
Table 8.2
| chemical shift $\delta$/ppm | environment of proton | splitting pattern | number of $^1\text{H}$ atoms responsible for the peak |
| :--- | :--- | :--- | :--- |
| 1.2 | | | |
| 2.8 | | | |
| 3.7 | | | |
| 3.9 | | |
(iii) Explain the splitting pattern of the peak at $\delta$ 3.9 ppm.
[Total: 15]9 (a) Define standard cell potential, $E_{\text{cell}}^{\ominus}$.
[1]
(b) An electrochemical cell is set up to measure $E_{\text{cell}}^{\ominuss}$ of a cell consisting of an Fe$^{3+}$/Fe$^{2+}$ half-cell and a Cl$_2$/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.
$$\text{Cl}_2 + 2\text{Fe}^{2+} \rightarrow 2\text{Fe}^{3+} + 2\text{Cl}^-$$
$$E_{\text{cell}}^{\ominuss} = +0.59\text{V}$$
Calculate $\Delta G^{\ominus}$, in kJ mol$^{-1}$, for this cell reaction.
$\Delta G^{\ominuss} = \dots\dots\dots\dots\kappa\text{J mol}^{-1}$ [2](d) Another experiment is set up using the same electrochemical cell.
In this experiment the Fe$^{2+}$ concentration is 0.15 mol dm$^{-3}$. All other concentrations remain at their standard values.
The Nernst equation is shown.
$$E = E^{\ominus} + (0.059/z) \log \frac{[\text{oxidised species}]}{[\text{reduced species}]}$$
(i) Use the Nernst equation to calculate the electrode potential, $E$, for the Fe$^{3+}/\text{Fe}^{2+}$ half-cell in this experiment.
$$[E^{\ominus}: \text{Fe}^{3+}/\text{Fe}^{2+\!=\!+0.77\text{ V}]}$$
$$E = \dots\dots\dots\dots\quad \text{V [1]}$$
(ii) Use your answer to (d)(i) to calculate $E_{\text{cell}}$ for this electrochemical cell.
$$E_{\text{cell}} = \dots\dots\dotted\dots\quad \text{V {[1]}}$$
[Total: 8]
Important values, constants and standards
| mol