2026 Cambridge IAL Chemistry (9701) Practice Paper with Answers

A sample of an unknown gaseous hydrocarbon has a mass of \(0.210\text{ g}\) and occupies a volume of \(120\text{ cm}^3\) at a temperature of \(298\text{ K}\) and a pressure of \(1.01 \times 10^5\text{ Pa}\).

What is the molecular formula of this hydrocarbon?

[\(R = 8.31\text{ J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1}\)]

The first eight successive ionisation energies of an element \(\text{X}\), in \(\text{kJ}\cdot\text{mol}^{-1}\), are shown:

\(738,\ 1450,\ 7730,\ 10540,\ 13630,\ 18000,\ 21700,\ 25660\)

Which formula represents the chloride of element \(\text{X}\)?

Which molecule has a trigonal pyramidal shape and a bond angle of approximately \(107^\circ\)?

The standard enthalpy changes of combustion, \(\Delta H^\circ_\text{c}\), are given:

- \(\Delta H^\circ_\text{c}(\text{C(s)}) = -394\text{ kJ}\cdot\text{mol}^{-1}\)
- \(\Delta H^\circ_\text{c}(\text{H}_2\text{(g)}) = -286\text{ kJ}\cdot\text{mol}^{-1}\)
- \(\Delta H^\circ_\text{c}(\text{C}_2\text{H}_5\text{OH(l)}) = -1367\text{ kJ}\cdot\text{mol}^{-1}\)

What is the standard enthalpy change of formation, \(\Delta H^\circ_\text{f}\), of ethanol, \(\text{C}_2\text{H}_5\text{OH(l)}\)?

Solid sodium iodide is treated with concentrated sulfuric acid at room temperature.

Which combination of sulfur-containing reduction products is formed?

The reversible reaction between dinitrogen tetraoxide and nitrogen dioxide is represented by the equation:

\[\text{N}_2\text{O}_4\text{(g)} \rightleftharpoons 2\text{NO}_2\text{(g)} \quad \Delta H = +57\text{ kJ}\cdot\text{mol}^{-1}\]

\(\text{N}_2\text{O}_4\text{(g)}\) is colourless and \(\text{NO}_2\text{(g)}\) is dark brown.

The temperature of an equilibrium mixture held in a sealed vessel of fixed volume is increased.

Which statement about the resulting equilibrium mixture is correct?

An organic compound \(\text{Y}\) with the molecular formula \(\text{C}_4\text{H}_{10}\text{O}\) forms a yellow precipitate when warmed with alkaline aqueous iodine.

When compound \(\text{Y}\) is heated under reflux with acidified potassium dichromate(VI), an organic compound \(\text{Z}\) is produced.

What is the structural formula of compound \(\text{Z}\)?

An unlabelled liquid sample is known to be one of the following compounds:

- propanoic acid
- propan-1-ol
- propanal
- ethyl ethanoate

The infrared spectrum of the liquid displays a very broad absorption band in the region \(2500\text{--}3000\text{ cm}^{-1}\) and a strong absorption peak at \(1715\text{ cm}^{-1}\).

Which compound is the liquid?

A \(0.120\text{ g}\) sample of an unknown volatile liquid \(X\) is vaporised in a gas syringe at \(373\text{ K}\) and \(1.01\times 10^5\text{ Pa}\).

The vapour occupies a volume of \(52.0\text{ cm}^3\).

What is the relative molecular mass, \(M_\text{r}\), of \(X\)?

Which species has the smallest bond angle?

Standard enthalpy changes of combustion, \(\Delta H_\text{c}^\ominus\), are given.

- \(\Delta H_\text{c}^\ominus [\text{C(graphite)}] = -394\text{ kJ mol}^{-1}\)
- \(\Delta H_\text{c}^\ominus [\text{H}_2\text{(g)}] = -286\text{ kJ mol}^{-1}\)
- \(\Delta H_\text{c}^\ominus [\text{CH}_3\text{COOH(l)}] = -874\text{ kJ mol}^{-1}\)

What is the standard enthalpy change of formation, \(\Delta H_\text{f}^\ominus\), of ethanoic acid, \(\text{CH}_3\text{COOH(l)}\)?

Which statement describing the properties of the Period 3 chlorides, from sodium chloride to phosphorus(V) chloride, is correct?

An excess of aqueous silver nitrate is added to an aqueous solution containing a single sodium halide, \(Y\). A cream precipitate is formed.

Concentrated aqueous ammonia is then added to this precipitate.

Which row correctly identifies halide \(Y\) and describes the observation when concentrated aqueous ammonia is added?

Compound \(W\) has the molecular formula \(\text{C}_4\text{H}_8\text{O}\).

- \(W\) forms an orange precipitate with 2,4-dinitrophenylhydrazine.
- \(W\) gives a yellow precipitate with alkaline aqueous iodine.
- \(W\) does not react with Tollens' reagent.

What is the identity of compound \(W\)?

How many structural isomers with the molecular formula \(\text{C}_5\text{H}_{12}\text{O}\) are secondary alcohols?

An organic compound \(Z\) contains carbon, hydrogen, and oxygen only.

- The mass spectrum of \(Z\) shows a molecular ion peak, \(\text{M}^+\), at \(m/e = 88\) with a relative abundance of \(24.0\).
- The \([\text{M}+1]^+\) peak has a relative abundance of \(1.06\).
- The infrared spectrum of \(Z\) shows a strong absorption peak at \(1740\text{ cm}^{-1}\) and no broad absorption between \(3200\text{ cm}^{-1}\) and \(3650\text{ cm}^{-1}\).

Which compound is \(Z\)?

An organic compound contains carbon, hydrogen, and oxygen only. Complete combustion of \(0.116\text{ g}\) of this compound in excess oxygen produces \(0.264\text{ g}\) of carbon dioxide and \(0.108\text{ g}\) of water.

What is the empirical formula of the compound?

A hydrocarbon molecule has the molecular formula \(\text{C}_4\text{H}_6\). It contains no rings and no triple bonds.

How many \(\sigma\) (sigma) bonds and how many \(\pi\) (pi) bonds are present in one molecule of this hydrocarbon?

Standard enthalpy changes of combustion, \(\Delta H^\circ_\text{c}\), are given:

\(\text{C}(\text{s}) + \text{O}_2(\text{g}) \rightarrow \text{CO}_2(\text{g})\quad \Delta H^\circ_\text{c} = -394\text{ kJ}\cdot\text{mol}^{-1}\)
\(\text{H}_2(\text{g}) + \frac{1}{2}\text{O}_2(\text{g}) \rightarrow \text{H}_2\text{O}(\text{l})\quad \Delta H^\circ_\text{c} = -286\text{ kJ}\cdot\text{mol}^{-1}\)
\(\text{C}_2\text{H}_5\text{OH}(\text{l}) + 3\text{O}_2(\text{g}) \rightarrow 2\text{CO}_2(\text{g}) + 3\text{H}_2\text{O}(\text{l})\quad \Delta H^\circ_\text{c} = -1368\text{ kJ}\cdot\text{mol}^{-1}\)

What is the standard enthalpy change of formation, \(\Delta H^\circ_\text{f}\), of liquid ethanol, \(\text{C}_2\text{H}_5\text{OH}(\text{l})\)?

The following reversible reaction reaches dynamic equilibrium in a closed vessel:

\[2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g}) \quad \Delta H = -196\text{ kJ}\cdot\text{mol}^{-1}\]

Which single change will increase the numerical value of the equilibrium constant, \(K_\text{c}\)?

Solid sodium halides react when treated separately with concentrated sulfuric acid.

Which row correctly describes the observations and the role of concentrated sulfuric acid in these reactions?

The temperature of a gaseous reaction mixture is increased from \(300\text{ K}\) to \(320\text{ K}\).

Which statement best explains the significant increase in the rate of reaction?

An alkene, \(\text{Z}\), has the molecular formula \(\text{C}_6\text{H}_{12}\) and exists as a pair of cis-trans (\(E\)-\(Z\)) stereoisomers.

When \(\text{Z}\) is reacted with hot concentrated acidified \(\text{KMnO}_4\), propanoic acid is obtained as the only organic product.

What is the identity of alkene \(\text{Z}\)?

An organic compound, \(\text{Q}\), gives an orange precipitate with 2,4-dinitrophenylhydrazine reagent.

\(\text{Q}\) does not produce a silver mirror when warmed with Tollens' reagent.

When \(\text{Q}\) is warmed with alkaline aqueous iodine, a yellow precipitate is formed.

Which compound is \(\text{Q}\)?

A \(10.0\text{ cm}^3\) sample of a gaseous hydrocarbon, \(\text{C}_x\text{H}_y\), is completely combusted in \(70.0\text{ cm}^3\) of oxygen (an excess).

After cooling to room temperature, the total volume of the remaining gaseous mixture is \(55.0\text{ cm}^3\).
When this remaining mixture is shaken with excess aqueous sodium hydroxide, the volume decreases to \(25.0\text{ cm}^3\).

All gas volumes are measured at the same temperature and pressure.

What is the molecular formula of the hydrocarbon?

The first six successive ionisation energies, in \(\text{kJ}\cdot\text{mol}^{-1}\), of an element \(X\) are given:

578, 1817, 2745, 11578, 14842, 18379

Which statement about element \(X\) is correct?

Which species has a bond angle of approximately \(107^\circ\) and possesses exactly one lone pair of electrons on its central atom?

The standard enthalpy changes of combustion, \(\Delta H^\circ_\text{c}\), of carbon(graphite), hydrogen(g), and propan-1-ol(l) are given:

\(\Delta H^\circ_\text{c}[\text{C(graphite)}] = -394\text{ kJ}\cdot\text{mol}^{-1}\)
\(\Delta H^\circ_\text{c}[\text{H}_2\text{(g)}] = -286\text{ kJ}\cdot\text{mol}^{-1}\)
\(\Delta H^\circ_\text{c}[\text{CH}_3\text{CH}_2\text{CH}_2\text{OH(l)}] = -2021\text{ kJ}\cdot\text{mol}^{-1}\)

What is the standard enthalpy change of formation, \(\Delta H^\circ_\text{f}\), of propan-1-ol(l)?

Two Period 3 elements, \(M\) and \(N\), each react with chlorine to form a chloride.

- The chloride of \(M\) is a white solid that dissolves in water to form a neutral solution with \(\text{pH} \approx 7\).
- The chloride of \(N\) is a colourless liquid at room temperature that reacts vigorously with water to form white misty fumes and an acidic solution with \(\text{pH} \approx 1-2\).

Which row correctly identifies elements \(M\) and \(N\)?

In an experiment, \(2.00\text{ mol}\) of \(\text{SO}_2\text{(g)}\) and \(2.00\text{ mol}\) of \(\text{O}_2\text{(g)}\) are mixed in a sealed vessel of volume \(2.00\text{ dm}^3\) at a constant temperature \(T\).

The equilibrium reaction is established:
\[2\text{SO}_2\text{(g)} + \text{O}_2\text{(g)} \rightleftharpoons 2\text{SO}_3\text{(g)}\]

At equilibrium, the amount of \(\text{SO}_3\text{(g)}\) present is \(1.60\text{ mol}\).

What is the numerical value of \(K_\text{c}\) under these conditions?

An unbranched alkene, \(W\), has the molecular formula \(\text{C}_6\text{H}_{12}\).

When \(W\) is heated under reflux with hot concentrated acidified potassium manganate(VII), the only organic product obtained is propanoic acid, \(\text{CH}_3\text{CH}_2\text{COOH}\).

How many stereoisomers exist for alkene \(W\)?

An organic compound \(Y\) with the molecular formula \(\text{C}_4\text{H}_{10}\text{O}\) exhibits the following properties:

- It turns acidified potassium dichromate(VI) from orange to green when heated.
- It forms a pale yellow precipitate when warmed with alkaline aqueous iodine.

What is the identity of compound \(Y\)?

A sample of \(0.295\text{ g}\) of an anhydrous metal carbonate, \(M\text{CO}_3\), was completely dissolved in \(50.0\text{ cm}^3\) of \(0.200\text{ mol dm}^{-3}\) hydrochloric acid.

The unreacted hydrochloric acid required \(15.0\text{ cm}^3\) of \(0.200\text{ mol dm}^{-3}\) sodium hydroxide for complete neutralisation.

What is the relative atomic mass, \(A_\text{r}\), of metal \(M\)?

Which sequence places the given species in order of increasing bond angle?

The standard enthalpy changes of combustion, \(\Delta H^\circ_\text{c}\), are given in the table.

$$\begin{array}{|c|c|}\hline\text{Substance} & \Delta H^\circ_\text{c} / \text{kJ mol}^{-1} \\hline\text{C(graphite)} & -394 \\text{H}_2\text{(g)} & -286 \\text{CH}_3\text{CH(OH)CH}_3\text{(l)} & -2006 \\hline\end{array}$$

What is the standard enthalpy change of formation, \(\Delta H^\circ_\text{f}\), of propan-2-ol, \(\text{CH}_3\text{CH(OH)CH}_3\text{(l)}\)?

In an experiment, \(2.00\text{ mol}\) of \(\text{SO}_2\text{(g)}\) and \(1.00\text{ mol}\) of \(\text{O}_2\text{(g)}\) are sealed in a vessel of fixed volume \(V\text{ dm}^3\) at a constant temperature.

The reaction reaches dynamic equilibrium according to the equation:

$$2\text{SO}_2\text{(g)} + \text{O}_2\text{(g)} \rightleftharpoons 2\text{SO}_3\text{(g)}$$

At equilibrium, \(1.60\text{ mol}\) of \(\text{SO}_3\text{(g)}\) is present in the mixture.

What is the value of the equilibrium constant, \(K_\text{c}\), in terms of \(V\)?

Solid samples of sodium halides are added separately to concentrated sulfuric acid.

Which statement is correct?

Hydrocarbon \(\text{W}\) has the molecular formula \(\text{C}_6\text{H}_{12}\) and exhibits geometrical (\(\text{cis}\)-\(\text{trans}\)) isomerism.

When \(\text{W}\) is heated under reflux with hot concentrated acidified potassium manganate(VII), propanoic acid, \(\text{CH}_3\text{CH}_2\text{COOH}\), is the only organic product obtained.

What is the structure of hydrocarbon \(\text{W}\)?

An organic compound \(\text{X}\) has the molecular formula \(\text{C}_4\text{H}_8\text{O}\).

- Compound \(\text{X}\) reacts with 2,4-dinitrophenylhydrazine to form an orange precipitate.
- Compound \(\text{X}\) does not produce a silver mirror when warmed with Tollens' reagent.
- Compound \(\text{X}\) forms a pale yellow precipitate when warmed with alkaline aqueous iodine.

What is the identity of \(\text{X}\)?

The mass spectrum of a halogenoalkane \(\text{Y}\) containing only carbon, hydrogen, and chlorine displays the following peak data:

- a molecular ion peak, \(\text{M}^+\), at \(m/e = 78\) with a relative abundance of \(60.0\%\)
- an \([\text{M}+1]^+\) peak at \(m/e = 79\) with a relative abundance of \(2.0\%\)
- an \([\text{M}+2]^+\) peak at \(m/e = 80\) with a relative abundance of \(19.8\%\)

How many carbon atoms and how many chlorine atoms are present in one molecule of \(\text{Y}\)?

(a) (i) Define relative atomic mass, \(A_\text{r}\).

(ii) A sample of magnesium consists of three isotopes: \(^{24}\text{Mg}\) (relative isotopic mass 23.99, abundance 78.99%), \(^{25}\text{Mg}\) (relative isotopic mass 24.99, abundance 10.00%), and \(^{26}\text{Mg}\) (relative isotopic mass 25.98, abundance 11.01%).

Calculate the relative atomic mass, \(A_\text{r}\), of this sample of magnesium. Give your answer to two decimal places.

(b) Compound \(\mathbf{X}\) is a volatile organic liquid containing only carbon, hydrogen, and oxygen.

Complete combustion of a \(1.480\text{ g}\) sample of \(\mathbf{X}\) produces \(3.520\text{ g}\) of carbon dioxide, \(\text{CO}_2\), and \(1.800\text{ g}\) of water, \(\text{H}_2\text{O}\).

(i) Show by calculation that the empirical formula of \(\mathbf{X}\) is \(\text{C}_4\text{H}_{10}\text{O}\).

(ii) In a separate experiment, a \(0.295\text{ g}\) sample of \(\mathbf{X}\) was vaporised in a gas syringe at \(373\text{ K}\) and \(1.00 \times 10^5\text{ Pa}\). The vapour occupied a volume of \(124\text{ cm}^3\).

Calculate the relative molecular mass, \(M_\text{r}\), of compound \(\mathbf{X}\) using the ideal gas equation. Give your answer to three significant figures. [\(R = 8.31\text{ J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1}\)]

(iii) Deduce the molecular formula of compound \(\mathbf{X}\).

(c) State two assumptions of the kinetic theory of ideal gases that do not hold true for real gases at high pressure and low temperature.

(a) (i) Define standard enthalpy change of formation, \(\Delta H_\text{f}^\circ\).

(ii) Hydrazine, \(\text{N}_2\text{H}_4(\text{l})\), reacts with dinitrogen tetroxide, \(\text{N}_2\text{O}_4(\text{g})\), to form nitrogen and steam according to the equation:

\[ 2\text{N}_2\text{H}_4(\text{l}) + \text{N}_2\text{O}_4(\text{g}) \rightarrow 3\text{N}_2(\text{g}) + 4\text{H}_2\text{O}(\text{g}) \]

Standard enthalpy changes of formation are given below:
- \(\Delta H_\text{f}^\circ [\text{N}_2\text{H}_4(\text{l})] = +50.6\text{ kJ}\cdot\text{mol}^{-1}\)
- \(\Delta H_\text{f}^\circ [\text{N}_2\text{O}_4(\text{g})] = +9.2\text{ kJ}\cdot\text{mol}^{-1}\)
- \(\Delta H_\text{f}^\circ [\text{H}_2\text{O}(\text{g})] = -241.8\text{ kJ}\cdot\text{mol}^{-1}\)

Calculate the standard enthalpy change of this reaction, \(\Delta H_\text{r}^\circ\).

(b) (i) State what is meant by the term dynamic equilibrium.

(ii) Nitrogen monoxide and oxygen react in a sealed container to establish the following equilibrium:

\[ 2\text{NO}(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{NO}_2(\text{g}) \quad \Delta H = -114\text{ kJ}\cdot\text{mol}^{-1} \]

Write the expression for the equilibrium constant, \(K_\text{c}\), for this reaction and deduce its units.

(iii) Predict and explain the effect of increasing the temperature on the position of equilibrium and the yield of \(\text{NO}_2(\text{g})\).

(iv) Explain the effect of increasing the total pressure on the initial rate of the forward reaction.

(a) Across Period 3 of the Periodic Table, elements display regular trends in electronic configuration, bonding, and physical properties.

(i) Explain why the first ionisation energy of sulfur is lower than that of phosphorus.

(ii) Describe and explain the trend in melting point from sodium to aluminium.

(b) Phosphorus and sulfur react with chlorine to form covalent chlorides.

(i) Solid phosphorus(V) chloride, \(\text{PCl}_5\), exists as a covalent molecule in the gas phase.
Draw the 3D shape of a gaseous \(\text{PCl}_5\) molecule, state the bond angles present, and give the name of its molecular geometry.

(ii) Write a balanced chemical equation for the reaction of solid phosphorus(V) chloride with water, and state the approximate pH of the resulting solution.

(c) Chlorine is used extensively in water treatment.

(i) Define the term disproportionation.

(ii) Write an ionic equation for the reaction of chlorine gas with cold, dilute aqueous sodium hydroxide.

(a) Halogenoalkanes are important synthetic intermediates undergoing substitution and elimination reactions.

(i) 2-bromo-2-methylpropane, \((\text{CH}_3)_3\text{CBr}\), reacts with aqueous sodium hydroxide to form 2-methylpropan-2-ol via an \(\text{S}_\text{N}1\) mechanism.

Outline the complete two-step mechanism for this reaction.
Include all relevant dipoles, curly arrows, charges, and lone pairs.

(ii) Explain why tertiary halogenoalkanes react predominantly via an \(\text{S}_\text{N}1\) mechanism rather than an \(\text{S}_\text{N}2\) mechanism.

(b) 1-bromobutane, \(\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_2\text{Br}\), reacts with hot ethanolic potassium hydroxide to produce but-1-ene.

(i) State the role of the hydroxide ion, \(\text{OH}^-\), in this elimination reaction.

(ii) When but-1-ene is reacted with hydrogen bromide, \(\text{HBr}\), two isomeric bromoalkanes, \(\mathbf{A}\) and \(\mathbf{B}\), are formed.

Identify the major organic product and explain why it is formed in greater yield by referring to the relative stabilities of the carbocation intermediates.

(c) Describe a chemical test to distinguish between 1-chlorobutane and 1-iodobutane using aqueous silver nitrate dissolved in ethanol. State the observations for each compound and explain the difference in the rate of precipitate formation.

(a) Compound \(\mathbf{W}\) has the molecular formula \(\text{C}_4\text{H}_8\text{O}\).

- Compound \(\mathbf{W}\) reacts with 2,4-dinitrophenylhydrazine (2,4-DNPH) to give an orange precipitate.
- Compound \(\mathbf{W}\) shows no reaction when warmed with Tollens' reagent.

(i) Deduce the structural formula of \(\mathbf{W}\) and state its IUPAC systematic name.

(ii) State the characteristic absorption range in the infrared spectrum that confirms the presence of the carbonyl functional group in \(\mathbf{W}\).

(iii) In the mass spectrum of \(\mathbf{W}\), a major fragmentation peak appears at \(m/e = 43\). Give the formula of the fragment ion responsible for this peak.

(b) Compound \(\mathbf{W}\) can be synthesised by the oxidation of a four-carbon alcohol, \(\mathbf{V}\).

(i) Draw the displayed formula of alcohol \(\mathbf{V}\), and state the reagents and conditions required to convert \(\mathbf{V}\) into \(\mathbf{W}\).

(ii) State the type of reaction occurring in (b)(i).

(c) Propanal, \(\text{CH}_3\text{CH}_2\text{CHO}\), reacts with hydrogen cyanide, \(\text{HCN}\), in the presence of a catalytic amount of \(\text{KCN}\) to form a hydroxynitrile.

(i) Complete the mechanism for this reaction.
Include all relevant lone pairs, dipoles, charges, and curly arrows, as well as the structures of the intermediate and the organic product.

(ii) Explain the role of \(\text{KCN}\) in this mechanism.

In this experiment, you are to determine the percentage purity by mass of hydrated iron(II) sulfate crystals, \(\text{FeSO}_4 \cdot 7\text{H}_2\text{O}\) (molar mass \(277.9\text{ g}\cdot\text{mol}^{-1}\)), in an impure commercial sample labelled FA 1.

You are provided with:
* FA 1: An aqueous solution prepared by dissolving \(14.20\text{ g}\) of impure hydrated iron(II) sulfate crystals in dilute sulfuric acid and making up to \(500.0\text{ cm}^3\) in a volumetric flask.
* FA 2: \(0.0200\text{ mol}\cdot\text{dm}^{-3}\) aqueous potassium manganate(VII), \(\text{KMnO}_4\).

(a) A candidate pipettes \(25.0\text{ cm}^3\) of FA 1 into a conical flask and titrates it with FA 2 from a burette until the first permanent pale pink colour appears. The recorded titration data are shown below:

$$\begin{array}{|l|c|c|c|}
\hline
\text{Titration number} & 1 & 2 & 3 \\
\hline
\text{Final burette reading } / \text{ cm}^3 & 24.35 & 48.40 & 24.10 \\
\hline
\text{Initial burette reading } / \text{ cm}^3 & 0.00 & 24.35 & 0.05 \\
\hline
\text{Titre } / \text{ cm}^3 & 24.35 & 24.05 & 24.05 \\
\hline
\end{array}$$

From the table above, select suitable titres and calculate the mean titre of FA 2. Show which titres you used in your calculation.

(b) Calculate the amount, in moles, of \(\text{MnO}_4^-\text{(aq)}\) present in your mean titre.

(c) The ionic equation for the redox titration is:
$$\text{MnO}_4^-(\text{aq}) + 5\text{Fe}^{2+}(\text{aq}) + 8\text{H}^+(\text{aq}) \rightarrow \text{Mn}^{2+}(\text{aq}) + 5\text{Fe}^{3+}(\text{aq}) + 4\text{H}_2\text{O}(\text{l})$$

(i) Calculate the amount, in moles, of \(\text{Fe}^{2+}\) present in the \(25.0\text{ cm}^3\) sample of FA 1.
(ii) Calculate the total amount, in moles, of \(\text{Fe}^{2+}\) present in the \(500.0\text{ cm}^3\) volumetric flask.

(d) Using your answer to (c)(ii), calculate the mass of pure \(\text{FeSO}_4 \cdot 7\text{H}_2\text{O}\) in the original sample and deduce the percentage purity by mass of the commercial sample. Give your answer to 3 significant figures.

(e) Each individual reading of the burette has an uncertainty of \(\pm 0.05\text{ cm}^3\). Calculate the maximum percentage apparatus error in the titre volume obtained in Titration 2.

In this experiment, you are to determine the value of \(x\) in hydrated zinc sulfate, \(\text{ZnSO}_4 \cdot x\text{H}_2\text{O}\), by measuring the mass loss upon thermal decomposition.

$$\text{ZnSO}_4 \cdot x\text{H}_2\text{O}(\text{s}) \xrightarrow{\Delta} \text{ZnSO}_4(\text{s}) + x\text{H}_2\text{O}(\text{g})$$

(a) A sample of hydrated zinc sulfate was placed in a dry, pre-weighed porcelain crucible with a lid and heated using a Bunsen burner. The table of results obtained is shown below:

$$\begin{array}{|l|c|}
\hline
\text{Measurement} & \text{Mass } / \text{ g} \\
\hline
\text{Mass of crucible + lid} & 18.42 \\
\hline
\text{Mass of crucible + lid + } \text{ZnSO}_4 \cdot x\text{H}_2\text{O} & 22.73 \\
\hline
\text{Mass of crucible + lid + residue after 1st heating} & 20.91 \\
\hline
\text{Mass of crucible + lid + residue after 2nd heating} & 20.84 \\
\hline
\text{Mass of crucible + lid + residue after 3rd heating} & 20.84 \\
\hline
\end{array}$$

Using the data from the table:
(i) Calculate the initial mass of \(\text{ZnSO}_4 \cdot x\text{H}_2\text{O}\) heated.
(ii) Calculate the mass of anhydrous \(\text{ZnSO}_4\) residue formed.
(iii) Calculate the mass of water, \(\text{H}_2\text{O}\), driven off.

(b) Use your results from (a) to:
(i) Calculate the amount, in moles, of anhydrous \(\text{ZnSO}_4\) formed. [\(M_r(\text{ZnSO}_4) = 161.5\)]
(ii) Calculate the amount, in moles, of \(\text{H}_2\text{O}\) lost. [\(M_r(\text{H}_2\text{O}) = 18.0\)]
(iii) Determine the empirical formula and the integer value of \(x\) in \(\text{ZnSO}_4 \cdot x\text{H}_2\text{O}\).

(c) State why the crucible and residue were heated a second and third time.

(d) A student noticed that some solid 'spat' out of the crucible during vigorous initial heating before the lid was placed on.
State and explain the effect this loss of solid would have on the calculated value of \(x\).

You are provided with two aqueous solutions, FA 3 and FA 4, each containing one cation and one anion from the Qualitative Analysis Notes.

(a) Carry out the tests described in the table below and deduce the observations. Write your observations in full, specifying the colour and state (precipitate/solution) of any products formed.

$$\begin{array}{|l|l|l|}
\hline
\textbf{Test} & \textbf{Observations for FA 3} & \textbf{Observations for FA 4} \\
\hline
\text{(i) Add aqueous sodium hydroxide} & & \\
\quad \text{dropwise, then in excess.} & & \\
\hline
\text{(ii) To the mixture from (i),} & & \\
\quad \text{warm gently with a piece of} & & \\
\quad \text{aluminium foil and test any} & & \\
\quad \text{gas evolved with damp red} & & \\
\quad \text{litmus paper.} & & \\
\hline
\text{(iii) Add aqueous ammonia} & & \\
\quad \text{dropwise, then in excess.} & & \\
\hline
\text{(iv) Add dilute nitric acid} & & \\
\quad \text{followed by aqueous barium nitrate.} & & \\
\hline
\text{(v) Add dilute nitric acid} & & \\
\quad \text{followed by aqueous silver nitrate,} & & \\
\quad \text{then add excess aqueous ammonia.} & & \\
\hline
\end{array}$$

Given observations:
* For FA 3:
- In (i): Green precipitate formed, insoluble in excess \(\text{NaOH}\).
- In (ii): No gas evolved; damp red litmus paper remains red.
- In (iii): Green precipitate formed, insoluble in excess \(\text{NH}_3\).
- In (iv): No precipitate formed.
- In (v): Cream/off-white precipitate formed, partially soluble/dissolves only in concentrated ammonia.
* For FA 4:
- In (i): No precipitate; on gentle warming without aluminium, no gas.
- In (ii): Effervescence; gas turns damp red litmus paper blue.
- In (iii): No precipitate.
- In (iv): Dense white precipitate formed, insoluble in excess dilute nitric acid.
- In (v): No precipitate.

(b) Identify the ions present in FA 3 and FA 4:
* Cation in FA 3:
* Anion in FA 3:
* Cation in FA 4:
* Anion in FA 4:

(c) Write an ionic equation, including state symbols, for the reaction between the cation in FA 3 and aqueous sodium hydroxide in test (a)(i).

(d) Write an ionic equation, including state symbols, for the precipitation reaction observed for FA 4 in test (a)(iv).

Calcium fluoride, \(\text{CaF}_2\), is a sparingly soluble ionic salt found naturally as the mineral fluorite.

(a) Define the term standard enthalpy change of atomisation.

(b) Relevant thermodynamic data for calcium and fluorine are given in the table below:

$$\begin{array}{|l|c|}
\hline
\text{Thermodynamic quantity} & \text{Value / } \text{kJ}\cdot\text{mol}^{-1} \\
\hline
\text{Standard enthalpy change of formation of } \text{CaF}_2(\text{s}) & -1220 \\
\text{Standard enthalpy change of atomisation of } \text{Ca}(\text{s}) & +178 \\
\text{First ionisation energy of calcium} & +590 \\
\text{Second ionisation energy of calcium} & +1145 \\
\text{Bond enthalpy of } \text{F}_2(\text{g}) & +158 \\
\text{First electron affinity of fluorine} & -328 \\
\hline
\end{array}$$

(i) Calculate the lattice energy, \(\Delta H_{\text{latt}}^\theta\), of calcium fluoride, \(\text{CaF}_2(\text{s})\).

(ii) The standard enthalpy changes of hydration of the gaseous ions are given below:
$$\Delta H_{\text{hyd}}^\theta(\text{Ca}^{2+}(\text{g})) = -1650 \text{ kJ}\cdot\text{mol}^{-1}$$
$$\Delta H_{\text{hyd}}^\theta(\text{F}^-(\text{g})) = -506 \text{ kJ}\cdot\text{mol}^{-1}$$

Calculate the standard enthalpy change of solution, \(\Delta H_{\text{sol}}^\theta\), of \(\text{CaF}_2(\text{s})\).

(c) The decomposition of solid calcium carbonate is an endothermic process:
$$\text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g}) \quad \Delta H^\theta = +178 \text{ kJ}\cdot\text{mol}^{-1}$$

The standard entropies, \(S^\theta\), of the species involved are:
$$\begin{array}{|c|c|}
\hline
\text{Species} & S^\theta / \text{J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1} \\
\hline
\text{CaCO}_3(\text{s}) & 92.9 \\
\text{CaO}(\text{s}) & 39.7 \\
\text{CO}_2(\text{g}) & 213.6 \\
\hline
\end{array}$$

(i) Calculate the standard entropy change, \(\Delta S^\theta\), for this reaction in \(\text{J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1}\).

(ii) Calculate the minimum temperature, in \(\text{K}\), at which this decomposition becomes spontaneous under standard pressure. State any assumption made.

Propanoic acid, \(\text{CH}_3\text{CH}_2\text{COOH}\), is a weak monobasic carboxylic acid.

(a) (i) Write the expression for the acid dissociation constant, \(K_{\text{a}}\), of propanoic acid.

(ii) At \(298\text{ K}\), a \(0.150\text{ mol}\cdot\text{dm}^{-3}\) solution of propanoic acid has a \(\text{pH}\) of \(2.85\). Calculate the value of \(K_{\text{a}}\) for propanoic acid at this temperature. State the units of \(K_{\text{a}}\).

(b) A buffer solution was prepared by adding \(50.0\text{ cm}^3\) of \(0.200\text{ mol}\cdot\text{dm}^{-3}\text{ NaOH}(\text{aq})\) to \(100.0\text{ cm}^3\) of \(0.300\text{ mol}\cdot\text{dm}^{-3}\text{ CH}_3\text{CH}_2\text{COOH}(\text{aq})\).

Assume \(K_{\text{a}} = 1.34 \times 10^{-5}\text{ mol}\cdot\text{dm}^{-3}\) for propanoic acid.

Calculate the \(\text{pH}\) of this buffer solution at \(298\text{ K}\).

(c) Propanoic acid partitions between ethoxyethane and water. In an experiment, \(100.0\text{ cm}^3\) of an aqueous solution containing \(0.0450\text{ mol}\) of propanoic acid was shaken with \(40.0\text{ cm}^3\) of ethoxyethane until equilibrium was reached at \(298\text{ K}\).

The partition coefficient is defined as:
$$K_{\text{partition}} = \frac{[\text{CH}_3\text{CH}_2\text{COOH}(\text{ethoxyethane})]}{[\text{CH}_3\text{CH}_2\text{COOH}(\text{aq})]} = 3.50$$

Calculate the mass, in \(\text{g}\), of propanoic acid that remains extracted into the organic layer (ethoxyethane).

Peroxodisulfate(\(\text{VI}\)) ions, \(\text{S}_2\text{O}_8^{2-}\), oxidise iodide ions to iodine in aqueous solution according to the equation:
$$\text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow 2\text{SO}_4^{2-}(\text{aq}) + \text{I}_2(\text{aq})$$

(a) The rate of this reaction was investigated at \(298\text{ K}\) by measuring initial rates with different initial concentrations of the reactants. The results are shown:

$$\begin{array}{|c|c|c|c|}
\hline
\text{Experiment} & [\text{S}_2\text{O}_8^{2-}] / \text{mol}\cdot\text{dm}^{-3} & [\text{I}^-] / \text{mol}\cdot\text{dm}^{-3} & \text{Initial rate} / \text{mol}\cdot\text{dm}^{-3}\cdot\text{s}^{-1} \\
\hline
1 & 0.0400 & 0.0300 & 1.44 \times 10^{-4} \\
2 & 0.0800 & 0.0300 & 2.88 \times 10^{-4} \\
3 & 0.0800 & 0.0900 & 8.64 \times 10^{-4} \\
\hline
\end{array}$$

(i) Deduce the order of reaction with respect to \(\text{S}_2\text{O}_8^{2-}\) and with respect to \(\text{I}^-\). Explain your reasoning.

(ii) Write the rate equation for this reaction.

(iii) Calculate the rate constant, \(k\), for this reaction at \(298\text{ K}\). State its units.

(b) In a separate experiment carried out with a large excess of iodide ions, \([\text{I}^-]\) remains effectively constant at \(0.500\text{ mol}\cdot\text{dm}^{-3}\). The half-life, \(t_{1/2}\), for the decay of \([\text{S}_2\text{O}_8^{2-}]\) is observed to be constant.

Calculate the half-life, in seconds, under these pseudo-first-order conditions.

(c) The uncatalysed reaction between \(\text{S}_2\text{O}_8^{2-}\) and \(\text{I}^-\) is slow because it has a high activation energy. \(\text{Fe}^{2+}(\text{aq})\) acts as a homogeneous catalyst.

(i) State why two negatively charged ions have a high activation energy for reaction.

(ii) Write two ionic equations to show how \(\text{Fe}^{2+}(\text{aq})\) catalyses this reaction.

Cobalt is a transition element that forms complex ions in different oxidation states.

(a) State the electronic configuration of the cobalt(\(\text{II}\)) ion, \(\text{Co}^{2+}\).

(b) When concentrated hydrochloric acid is added to a pink aqueous solution containing \([\text{Co}(\text{H}_2\text{O})_6]^{2+}\), the solution turns blue as the tetrahedral complex \([\text{CoCl}_4]^{2-}\) is formed.

(i) Write an equation for this equilibrium reaction.

(ii) State the coordination number and geometry of cobalt in \([\text{Co}(\text{H}_2\text{O})_6]^{2+}\) and in \([\text{CoCl}_4]^{2-}\).

(iii) Explain why transition metal complexes are coloured, making reference to d-orbitals, ligands, and visible light absorption.

(c) The stability constant, \(K_{\text{stab}}\), for the complex ion formed between \(\text{Co}^{2+}\) and 1,2-diaminoethane (abbreviated as \(\text{en}\), \(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2\)) is \(1.0 \times 10^{14}\text{ dm}^9\cdot\text{mol}^{-3}\) at \(298\text{ K}\).

$$[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 3\text{en} \rightleftharpoons [\text{Co}(\text{en})_3]^{2+} + 6\text{H}_2\text{O}$$

(i) Write the expression for \(K_{\text{stab}}\) for this equilibrium.

(ii) In terms of entropy change, explain why \([\text{Co}(\text{en})_3]^{2+}\) is significantly more stable than \([\text{Co}(\text{NH}_3)_6]^{2+}\) (the chelate effect).

Standard electrode potentials can be used to predict the feasibility of redox reactions.

(a) Standard reduction potentials at \(298\text{ K}\) are given below:
$$\begin{array}{rcll}
\text{Fe}^{3+}(\text{aq}) + \text{e}^- & \rightleftharpoons & \text{Fe}^{2+}(\text{aq}) & E^\theta = +0.77\text{ V} \\
\text{Ag}^+(\text{aq}) + \text{e}^- & \rightleftharpoons & \text{Ag}(\text{s}) & E^\theta = +0.80\text{ V} \\
\text{Cu}^{2+}(\text{aq}) + 2\text{e}^- & \rightleftharpoons & \text{Cu}(\text{s}) & E^\theta = +0.34\text{ V}
\end{array}$$

(i) Define standard electrode potential.

(ii) An electrochemical cell is constructed under standard conditions using the \(\text{Fe}^{3+}/\text{Fe}^{2+}\) half-cell and the \(\text{Ag}^+/\text{Ag}\) half-cell. Calculate the standard cell potential, \(E_{\text{cell}}^\theta\), and state which electrode is the positive terminal.

(iii) Write the overall ionic equation for the cell reaction that occurs spontaneously under standard conditions.

(b) The electrode potential of a non-standard half-cell can be calculated using the Nernst equation:
$$E = E^\theta + \frac{0.059}{z}\log_{10}\left(\frac{[\text{oxidised form}]}{[\text{reduced form}]}\right)$$

Calculate the electrode potential, \(E\), of a copper half-cell when \([\text{Cu}^{2+}] = 0.00200\text{ mol}\cdot\text{dm}^{-3}\) at \(298\text{ K}\).

(c) In an industrial electroplating experiment, a metal object is coated with copper by electrolysing aqueous copper(II) sulfate. A steady current of \(3.50\text{ A}\) was passed through the cell for \(45.0\text{ minutes}\).

Calculate the mass of copper, in \(\text{g}\), deposited on the electrode. \([F = 9.65 \times 10^4\text{ C}\cdot\text{mol}^{-1}; A_{\text{r}}(\text{Cu}) = 63.5]\)

Benzene and its derivatives undergo characteristic electrophilic substitution reactions.

(a) Benzene reacts with a mixture of concentrated nitric acid and concentrated sulfuric acid at \(50-55^\circ\text{C}\) to form nitrobenzene.

(i) Write an equation for the generation of the electrophile, \(\text{NO}_2^+\), from the two concentrated acids.

(ii) Describe the mechanism for the nitration of benzene. In your answer, include curly arrows to show the movement of electron pairs, the intermediate ion structure, and the regenerator of the catalyst.

(b) Nitrobenzene can be converted into phenylamine, \(\text{C}_6\text{H}_5\text{NH}_2\).

(i) State the reagents and conditions needed to reduce nitrobenzene to phenylamine in the laboratory.

(ii) Explain why phenylamine is a weaker base than ethylamine, \(\text{CH}_3\text{CH}_2\text{NH}_2\).

(c) Phenylamine is used to synthesise azo dyes. In the first step, phenylamine is treated with nitrous acid, \(\text{HNO}_2\), at \(5^\circ\text{C}\) to form a diazonium salt.

(i) State how nitrous acid is prepared in situ.

(ii) Draw the structural formula of the azo dye formed when benzenediazonium chloride reacts with alkaline phenol, \(\text{C}_6\text{H}_5\text{O}^-\).

Acyl chlorides are versatile synthetic intermediates in organic chemistry.

(a) Ethanoyl chloride, \(\text{CH}_3\text{COCl}\), reacts vigorously with water at room temperature.

(i) Write the equation for the reaction of ethanoyl chloride with water.

(ii) Compare and explain the relative rates of hydrolysis of ethanoyl chloride, chloroethane (\(\text{CH}_3\text{CH}_2\text{Cl}\)), and chlorobenzene (\(\text{C}_6\text{H}_5\text{Cl}\)).

(b) Ethanoyl chloride reacts with methylamine, \(\text{CH}_3\text{NH}_2\), to form an \(N\)-substituted amide.

(i) Give the structural formula and systematic name of the organic product formed.

(ii) Complete the mechanism for the reaction between \(\text{CH}_3\text{COCl}\) and \(\text{CH}_3\text{NH}_2\). Include dipoles, lone pairs, and curly arrows showing addition and subsequent elimination steps.

(c) Polyamides and polyesters are condensation polymers.

A section of a biodegradable polymer chain is shown below:
$$-\left[\text{O}-\text{CH}_2\text{CH}_2-\text{O}-\text{CO}-\text{C}_6\text{H}_4-\text{CO}\right]_n-$$

(i) Draw the skeletal or structural formulas of the two monomers required to synthesise this polymer.

(ii) Explain why condensation polymers containing ester linkages are biodegradable, whereas addition polymers like poly(ethene) are non-biodegradable.

Compound X is an aromatic compound with the molecular formula \(\text{C}_9\text{H}_{10}\text{O}_2\).

(a) Complete combustion of \(0.300\text{ g}\) of X gives \(0.792\text{ g}\) of \(\text{CO}_2\) and \(0.180\text{ g}\) of \(\text{H}_2\text{O}\).

Show that this combustion data is consistent with the empirical formula \(\text{C}_9\text{H}_{10}\text{O}_2\).

(b) The spectroscopic data for compound X are summarised below:
- Infrared spectrum: shows a strong sharp absorption at \(1715\text{ cm}^{-1}\) and a broad absorption between \(2500-3300\text{ cm}^{-1}\).
- **\(^{13}\text{C}\) NMR spectrum**: displays \(7\) distinct peaks.
- **\(^1\text{H}\) NMR spectrum** (in \(\text{CDCl}_3\)):
$$\begin{array}{|c|c|c|c|}
\hline
\text{Chemical shift, } \delta / \text{ppm} & \text{Splitting pattern} & \text{Relative peak area} & \text{Assignment} \\
\hline
1.20 & \text{doublet} & 6 & \text{Protons on } \text{C}\mathbf{H}_3 \\
2.95 & \text{septet} & 1 & \text{Proton on } \text{C}\mathbf{H} \\
7.20-7.35 & \text{multiplet} & 2 & \text{Aromatic ring protons} \\
7.95-8.05 & \text{multiplet} & 2 & \text{Aromatic ring protons} \\
11.80 & \text{singlet} & 1 & \text{Exchangeable with } \text{D}_2\text{O} \\
\hline
\end{array}$$

(i) Identify the functional group responsible for the IR absorptions at \(1715\text{ cm}^{-1}\) and \(2500-3300\text{ cm}^{-1}\).

(ii) Identify the alkyl substituent attached to the aromatic ring indicated by the signals at \(\delta = 1.20\text{ ppm}\) and \(\delta = 2.95\text{ ppm}\). Explain the observed splitting patterns.

(iii) Deduce the substitution pattern on the aromatic ring from the \(^1\text{H}\) and \(^{13}\text{C}\) NMR spectra.

(iv) Draw the displayed or structural formula of compound X.

Group 2 carbonates and nitrates exhibit distinct trends in their thermal stability down the group.

(a) State the trend in the thermal stability of Group 2 carbonates from \(\text{MgCO}_3\) to \(\text{BaCO}_3\). Explain this trend in terms of ionic radius and polarising power of the cation.

(b) When a \(3.47\text{ g}\) sample of anhydrous Group 2 nitrate, \(\text{M}(\text{NO}_3)_2\), is heated strongly until no further reaction occurs, brown nitrogen dioxide gas and oxygen gas are evolved according to the equation:
$$2\text{M}(\text{NO}_3)_2(\text{s}) \rightarrow 2\text{MO}(\text{s}) + 4\text{NO}_2(\text{g}) + \text{O}_2(\text{g})$$

The total volume of gas collected at room conditions (\(298\text{ K}\) and \(101\text{ kPa}\)) was \(1.20\text{ dm}^3\).

[Molar volume of gas at room conditions = \(24.0\text{ dm}^3\cdot\text{mol}^{-1}\)]

(i) Calculate the total amount, in moles, of gas collected.

(ii) Calculate the amount, in moles, of \(\text{M}(\text{NO}_3)_2\) that reacted.

(iii) Calculate the relative formula mass, \(M_{\text{r}}\), of \(\text{M}(\text{NO}_3)_2\) and hence identify metal \(\text{M}\).

(c) The solid residue, \(\text{MO}\), reacts with dilute sulfuric acid to form a white precipitate.

(i) State the trend in the solubility of Group 2 sulfates down the group.

(ii) Write an ionic equation, including state symbols, for the formation of the white precipitate when the aqueous solution of metal \(\text{M}\) reacts with aqueous sulfate ions.

Nickel is a transition element that forms a variety of octahedral and tetrahedral complexes in the +2 oxidation state. (a) Define what is meant by the stability constant, \(K_{\text{stab}}\), of a complex ion. [1] (b) Aqueous nickel(II) ions exist as the octahedral complex \([\text{Ni}(\text{H}_2\text{O})_6]^{2+}\). When 1,2-diaminoethane, \(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2\) (abbreviated as \(\text{en}\)), is added to an aqueous solution of nickel(II) sulfate, complete ligand exchange occurs to form \([\text{Ni}(\text{en})_3]^{2+}\). (i) Write an ionic equation for this ligand exchange reaction. [1] (ii) Explain, in terms of entropy and Gibbs free energy, why \([\text{Ni}(\text{en})_3]^{2+}\) is significantly more stable than \([\text{Ni}(\text{NH}_3)_6]^{2+}\), despite both complexes having six nickel-nitrogen coordinate bonds. [2] (iii) State the type of stereoisomerism shown by \([\text{Ni}(\text{en})_3]^{2+}\) and draw three-dimensional representations of the two stereoisomers. [2] (c) When concentrated hydrochloric acid is added dropwise to \([\text{Ni}(\text{H}_2\text{O})_6]^{2+}(\text{aq})\), the green solution changes to yellow-green due to the formation of the complex ion \([\text{NiCl}_4]^{2-}\). (i) State the coordination number and geometry of \([\text{NiCl}_4]^{2-}\). [1] (ii) Describe and explain why aqueous solutions containing \([\text{Ni}(\text{H}_2\text{O})_6]^{2+}\) are coloured. [3]

Phenylamine, \(\text{C}_6\text{H}_5\text{NH}_2\), is a primary aromatic amine widely used as a precursor in the pharmaceutical and dye industries. (a) (i) State the reagents and conditions necessary to synthesise phenylamine from nitrobenzene in the laboratory. [1] (ii) Explain why phenylamine is a significantly weaker base than ethylamine, \(\text{CH}_3\text{CH}_2\text{NH}_2\). [2] (b) When phenylamine is reacted with bromine water, \(\text{Br}_2(\text{aq})\), at room temperature, an immediate reaction occurs without needing a halogen carrier catalyst. (i) State the observations made during this reaction. [1] (ii) Draw the structural formula of the organic product formed and state its systematic IUPAC name. [1] (c) Compound X is an aromatic amide with the molecular formula \(\text{C}_8\text{H}_9\text{NO}\). When heated under reflux with aqueous sodium hydroxide, Compound X is hydrolysed into phenylamine and sodium ethanoate. (i) Deduce the structural formula of Compound X and give its systematic name. [2] (ii) Deduce the number of peaks present in the \(^{13}\text{C}\text{ NMR}\) spectrum of Compound X. [1] (iii) Predict the splitting pattern and relative peak area of the signal corresponding to the methyl group protons in the \(^1\text{H}\text{ NMR}\) spectrum of Compound X. [2]

A student plans an experiment to determine the partition coefficient, \(K_{\text{partition}}\), of ethanoic acid between ethoxyethane (an organic solvent) and water at \(298\text{ K}\).

\(K_{\text{partition}}\) is defined by the expression:
\[K_{\text{partition}} = \frac{[\text{CH}_3\text{COOH}_{(\text{ethoxyethane})}]}{[\text{CH}_3\text{COOH}_{(\text{aq})}]}\]

Ethoxyethane and water are immiscible liquids. Ethoxyethane has a density of \(0.713\text{ g}\cdot\text{cm}^{-3}\) and is highly flammable and volatile (boiling point \(34.6\,^\circ\text{C}\)).

You are provided with:
- pure ethanoic acid, \(\text{CH}_3\text{COOH}(\text{l})\)
- pure ethoxyethane, \(\text{CH}_3\text{CH}_2\text{OCH}_2\text{CH}_3(\text{l})\)
- distilled water
- standard \(0.100\text{ mol}\cdot\text{dm}^{-3}\ \text{NaOH}(\text{aq})\)
- phenolphthalein indicator
- standard laboratory glassware and equipment.

(a) Identify the independent variable and the dependent variable in this investigation. [2]

(b) (i) Describe the preparation of \(250.0\text{ cm}^3\) of an aqueous solution containing approximately \(0.500\text{ mol}\cdot\text{dm}^{-3}\) ethanoic acid, using pure liquid ethanoic acid (density = \(1.05\text{ g}\cdot\text{cm}^{-3}\), \(M_{\text{r}} = 60.0\)). Include the calculated mass and volume of ethanoic acid required. [3]

(ii) Describe how the two immiscible layers are brought to equilibrium and separated safely. State the specific apparatus used. [3]

(c) Describe the titration procedure to determine the concentration of ethanoic acid remaining in the aqueous layer after equilibrium has been established. [3]

(d) State one specific safety hazard associated with ethoxyethane and describe an appropriate precaution to minimise the risk during the experiment. [1]

(e) In one experiment, \(50.0\text{ cm}^3\) of the \(0.500\text{ mol}\cdot\text{dm}^{-3}\) aqueous ethanoic acid was shaken with \(50.0\text{ cm}^3\) of ethoxyethane. After equilibrium was established and the layers separated, a \(10.0\text{ cm}^3\) sample of the aqueous layer required \(32.40\text{ cm}^3\) of \(0.100\text{ mol}\cdot\text{dm}^{-3}\ \text{NaOH}(\text{aq})\) for neutralisation.

Calculate:
(i) the concentration of ethanoic acid in the aqueous layer at equilibrium, [1]
(ii) the concentration of ethanoic acid in the ethoxyethane layer at equilibrium, [1]
(iii) the value of \(K_{\text{partition}}\). [1]

A student investigates the kinetics of the reaction between benzenediazonium chloride, \(\text{C}_6\text{H}_5\text{N}_2\text{Cl}(\text{aq})\), and water at different temperatures:
\[\text{C}_6\text{H}_5\text{N}_2\text{Cl}(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightarrow \text{C}_6\text{H}_5\text{OH}(\text{aq}) + \text{HCl}(\text{aq}) + \text{N}_2(\text{g})\]

The reaction is first order with respect to \([\text{C}_6\text{H}_5\text{N}_2\text{Cl}]\). The rate constant, \(k\), is related to temperature, \(T\), by the Arrhenius equation:
\[\ln k = -\frac{E_{\text{a}}}{R}\left(\frac{1}{T}\right) + \ln A\]
where \(R = 8.31\text{ J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1}\) and \(E_{\text{a}}\) is the activation energy in \(\text{J}\cdot\text{mol}^{-1}\).

The student measures the time taken, \(t_{1/2}\), for half of the benzenediazonium chloride to decompose at five different temperatures. The rate constant \(k\) is calculated using:
\[k = \frac{0.693}{t_{1/2}}\]

The experimental results are recorded in the table below.

\[\begin{array}{|c|c|c|c|c|c|} \hline \text{Experiment} & T / \,^\circ\text{C} & T / \text{K} & \frac{1}{T} / 10^{-3}\text{ K}^{-1} & k / 10^{-4}\text{ s}^{-1} & \ln (k / \text{s}^{-1}) \\ \hline 1 & 25.0 & 298.2 & 3.353 & 0.420 & -10.08 \\ 2 & 35.0 & 308.2 & 3.245 & 1.62 & -8.73 \\ 3 & 45.0 & 318.2 & 3.143 & 5.75 & -7.46 \\ 4 & 55.0 & 328.2 & 3.047 & 15.1 & \text{........} \\ 5 & 65.0 & 338.2 & \text{........} & 62.0 & -5.08 \\ \hline \end{array}\]

(a) Complete the table by calculating:
(i) the value of \(\frac{1}{T} \times 10^{-3}\text{ K}^{-1}\) for Experiment 5, [1]
(ii) the value of \(\ln (k / \text{s}^{-1})\) for Experiment 4. [1]

(b) Describe how a graph of \(\ln (k / \text{s}^{-1})\) against \(\frac{1}{T}\) is used to determine the activation energy, \(E_{\text{a}}\). State how the gradient of the line is related to \(E_{\text{a}}\). [2]

(c) The coordinates of two points on the line of best fit drawn by the student are:
- Point 1: \(\left(2.950 \times 10^{-3}\text{ K}^{-1},\ -4.30\right)\)
- Point 2: \(\left(3.350 \times 10^{-3}\text{ K}^{-1},\ -10.02\right)\)

(i) Calculate the gradient of the line of best fit. Include appropriate signs. [2]
(ii) Use your gradient from (c)(i) and \(R = 8.31\text{ J}\cdot\text{K}^{-1}\cdot\text{mol}^{-1}\) to calculate the activation energy, \(E_{\text{a}}\), in \(\text{kJ}\cdot\text{mol}^{-1}\). Give your answer to 3 significant figures. [2]

(d) The volume of \(\text{N}_2(\text{g})\) evolved was collected in a \(100.0\text{ cm}^3\) gas syringe with graduations every \(1.0\text{ cm}^3\). In Experiment 1, the total volume of \(\text{N}_2(\text{g})\) collected at completion was \(48.0\text{ cm}^3\).

(i) Calculate the percentage uncertainty in measuring this volume of gas. [1]
(ii) Suggest one reason why collecting nitrogen gas over water in an inverted burette might lead to an inaccurate volume measurement compared to using a gas syringe. [1]

(e) In Experiment 4, the measured half-life was slightly shorter than expected, causing the calculated value of \(k\) to be an anomalous point above the line of best fit. Suggest one experimental error in temperature control that could cause this anomaly. [1]

(f) Explain why the student must keep the total volume of the reaction mixture constant in all five experiments. [2]

(g) Suggest a suitable piece of apparatus to maintain a constant temperature of \(65.0\,^\circ\text{C}\) throughout Experiment 5. [1]