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Using the standard enthalpy of combustion values given below, what is the standard enthalpy of formation of liquid propan-2-ol, \(\text{CH}_3\text{CH(OH)CH}_3(l)\)?

\(\Delta H_c^\theta[\text{C}(s)] = -393.5 \text{ kJ mol}^{-1}\)
\(\Delta H_c^\theta[\text{H}_2(g)] = -285.8 \text{ kJ mol}^{-1}\)
\(\Delta H_c^\theta[\text{CH}_3\text{CH(OH)CH}_3(l)] = -2006.0 \text{ kJ mol}^{-1}\)

An element, X, is in Period 3 of the Periodic Table. The first six successive ionisation energies (in \(\text{kJ mol}^{-1}\)) of X are shown in the table below:

1st: 1010
2nd: 1900
3rd: 2910
4th: 4960
5th: 6270
6th: 21300

Which formula represents the chloride formed when element X is reacted with excess chlorine gas?

A sample of 0.180 g of a volatile organic compound of molecular formula \(\text{C}_n\text{H}_{2n+2}\text{O}\) is completely vaporised at a temperature of \(127^\circ\text{C}\) and a pressure of \(1.00 \times 10^5 \text{ Pa}\). The volume of vapor produced is \(99.7 \text{ cm}^3\).

What is the molecular formula of the compound?
[\(R = 8.31 \text{ J K}^{-1}\text{ mol}^{-1}\)]

Four different halogenoalkanes are added separately to mixtures of ethanol and aqueous silver nitrate. The temperature of each mixture is maintained at \(50^\circ\text{C}\).

Which halogenoalkane produces a precipitate in the shortest time?

An increase in temperature from 300 K to 310 K often approximately doubles the rate of a chemical reaction.

Which statement correctly explains this observation?

How many structural isomers are there with the molecular formula \(\text{C}_3\text{H}_6\text{Cl}_2\)?

A solid mixture contains 0.010 mol of anhydrous magnesium nitrate, \(\text{Mg(NO}_3\text{)}_2\), and 0.010 mol of anhydrous barium nitrate, \(\text{Ba(NO}_3\text{)}_2\). The mixture is heated strongly in an open crucible until there is no further change in mass.

What is the total mass of solid residue remaining in the crucible?
[\(A_r\) values: \(\text{Mg} = 24.3\), \(\text{Ba} = 137.3\), \(\text{O} = 16.0\)]

In a closed container, 2.00 mol of \(\text{SO}_2(g)\) and 2.00 mol of \(\text{O}_2(g)\) are mixed at a constant temperature. The equilibrium mixture is represented by the equation:

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

At equilibrium, 1.60 mol of \(\text{SO}_3(g)\) is present. If the total volume of the container is \(2.00 \text{ dm}^3\), what is the value of the equilibrium constant, \(K_c\), at this temperature?

A student carries out an experiment to determine the enthalpy change of combustion of pentan-1-ol (\(M_r = 88.0\)). A spirit burner containing pentan-1-ol is burned beneath a beaker containing \(150\text{ g}\) of water. The temperature of the water rises from \(20.5\text{ }^{\circ}\text{C}\) to \(44.5\text{ }^{\circ}\text{C}\). The mass of the spirit burner decreases by \(0.440\text{ g}\). The specific heat capacity of water is \(4.18\text{ J g}^{-1}\text{ K}^{-1}\). What is the enthalpy change of combustion of pentan-1-ol under these conditions?

For a particular gas-phase reaction, when the temperature is increased from \(300\text{ K}\) to \(310\text{ K}\), the rate of reaction approximately doubles. Which of the following statements correctly explains this observation?

Which of the following species has the same number of unpaired d-electrons as the \(\text{Co}^{2+}\) ion in its ground state?

A student aims to synthesize propanoic acid from ethene in a three-step pathway:

\(\text{ethene} \xrightarrow{\text{Step 1}} \text{Compound X} \xrightarrow{\text{Step 2}} \text{Compound Y} \xrightarrow{\text{Step 3}} \text{propanoic acid}\)

Which set of reagents and conditions is most suitable for this synthesis?

Consider the following heterogeneous equilibrium at a constant temperature \(T\):

\(\text{C(s)} + \text{CO}_2\text{(g)} \rightleftharpoons 2\text{CO(g)} \quad \Delta H = +172.5\text{ kJ mol}^{-1}\)

The volume of the container is suddenly halved while maintaining the temperature at \(T\). Once the new equilibrium is established, how do the value of \(K_c\) and the mass of \(\text{C(s)}\) compare to their initial values?

In the redox reaction between acidified potassium dichromate(VI) and tin(II) ions, the unbalanced equation is shown below:

\(\text{Cr}_2\text{O}_7^{2-} + a\text{H}^+ + b\text{Sn}^{2+} \rightarrow 2\text{Cr}^{3+} + c\text{H}_2\text{O} + d\text{Sn}^{4+}\)

What is the value of the coefficient \(b\) when the equation is balanced with the simplest whole-number coefficients?

The standard enthalpy changes of combustion, \(\Delta H_c^{\ominus}\), for carbon, hydrogen, and propene are given below:

- \(\Delta H_c^{\ominus}[\text{C(graphite)}] = -393.5\text{ kJ mol}^{-1}\)
- \(\Delta H_c^{\ominus}[\text{H}_2\text{(g)}] = -285.8\text{ kJ mol}^{-1}\)
- \(\Delta H_c^{\ominus}[\text{C}_3\text{H}_6\text{(g)}] = -2058.0\text{ kJ mol}^{-1}\)

What is the standard enthalpy change of formation, \(\Delta H_f^{\ominus}\), of propene gas?

The reaction between peroxodisulfate ions, \(\text{S}_2\text{O}_8^{2-}\), and iodide ions, \(\text{I}^-\), is catalyzed by aqueous iron(II) ions, \(\text{Fe}^{2+}\)(aq):

\(\text{S}_2\text{O}_8^{2-}(aq) + 2\text{I}^-(aq) \xrightarrow{\text{Fe}^{2+}(aq)} 2\text{SO}_4^{2-}(aq) + \text{I}_2(aq)\)

Which statement best explains why \(\text{Fe}^{2+}\)(aq) acts as an effective homogeneous catalyst in this reaction?

A student burns 0.32 g of methanol (\(CH_3OH\), \(M_r = 32.0\)) to heat 150 cm³ of water in a copper beaker. The temperature of the water rises by 12.0 °C. The specific heat capacity of water is 4.18 J g⁻¹ K⁻¹. What is the experimental enthalpy change of combustion of methanol, \(\Delta H_c^\theta\), in kJ mol⁻¹?

Which statement about the Maxwell–Boltzmann energy distribution curve for a gaseous reaction is correct when the temperature is increased?

Nitrogen dioxide dissociates according to the equation: \(2\text{NO}_2(\text{g}) \rightleftharpoons 2\text{NO}(\text{g}) + \text{O}_2(\text{g})\). At a certain temperature, 2.00 mol of \(\text{NO}_2\) is placed in a closed \(2.00 \text{ dm}^3\) vessel. At equilibrium, 0.60 mol of \(\text{O}_2\) is present. What is the value of \(K_c\) at this temperature?

In the redox reaction between sodium thiosulfate and iodine, sodium tetrathionate (\(\text{Na}_2\text{S}_4\text{O}_6\)) and sodium iodide are formed: \(2\text{Na}_2\text{S}_2\text{O}_3 + \text{I}_2 \rightarrow \text{Na}_2\text{S}_4\text{O}_6 + 2\text{NaI}\). What is the change in the oxidation number of sulfur in this reaction?

An organic compound X has the molecular formula \(\text{C}_4\text{H}_{10}\text{O}\). When X is warmed with acidified potassium dichromate(VI), the solution turns green. The organic product Y does not react with Fehling's solution. Which of the following is the IUPAC name of X?

Bromoethane can be converted into propanoic acid in a two-step synthesis: \(\text{CH}_3\text{CH}_2\text{Br} \xrightarrow{\text{Step 1}} \text{Compound X} \xrightarrow{\text{Step 2}} \text{CH}_3\text{CH}_2\text{CO}_2\text{H}\). Which reagents and conditions are correct for Step 1 and Step 2?

Two halogenoalkanes, 2-chloro-2-methylpropane and 1-chlorobutane, are hydrolyzed separately with aqueous sodium hydroxide. Which row correctly describes the predominant mechanism for each reaction?

Which statement about the trends in the properties of Group 2 elements and their compounds from magnesium to barium is correct?

An element X in Period 3 has the following successive ionisation energies in \(\text{kJ mol}^{-1}\):

\(1^{\text{st}} = 578\)

\(2^{\text{nd}} = 1817\)

\(3^{\text{rd}} = 2745\)

\(4^{\text{th}} = 11578\)

\(5^{\text{th}} = 14831\)

What is the formula of the oxide of X?

A student uses a simple calorimeter to determine the enthalpy change of combustion of methanol (\(\text{CH}_3\text{OH}\), \(M_{\text{r}} = 32.0\)).

0.80 g of methanol is burned to heat 150 g of water from \(20.0^\circ\text{C}\) to \(41.5^\circ\text{C}\).

The specific heat capacity of water is \(4.18\text{ J g}^{-1}\text{ K}^{-1}\).

Assuming that only \(60\%\) of the heat released by the combustion of methanol is transferred to the water, what is the calculated enthalpy of combustion of methanol?

For a gas-phase reaction, the temperature of the system is increased from \(T_1\) to \(T_2\) in the absence of a catalyst.

How do the most probable energy of the molecules and the activation energy, \(E_{\text{a}}\), of the reaction change?

The following dynamic equilibrium is established in a closed container of fixed volume:

\[2\text{SO}_2(\text{g}) + \text{O}_2(\text{g}) \rightleftharpoons 2\text{SO}_3(\text{g}) \quad \Delta H < 0\]

A small amount of helium gas is injected into the container at constant volume.

How do the position of the equilibrium and the value of the equilibrium constant, \(K_{\text{c}}\), change?

An organic compound Y is a secondary alcohol with the molecular formula \(\text{C}_5\text{H}_{12}\text{O}\). When Y is heated with concentrated sulfuric acid, it undergoes dehydration to produce a mixture containing exactly three isomeric alkenes (including stereoisomers).

What is the IUPAC name of Y?

Equal amounts of 1-chlorobutane, 1-bromobutane, and 1-iodobutane are heated separately with aqueous silver nitrate in ethanol at \(50^\circ\text{C}\).

Which statement correctly explains the relative rate of reaction?

Equal masses of anhydrous calcium nitrate, \(\text{Ca(NO}_3)_2\), and anhydrous barium nitrate, \(\text{Ba(NO}_3)_2\), are heated strongly in separate test tubes until they completely decompose.

Which statement about these experiments is correct?

Which molecule has a non-zero permanent dipole moment?

In a calorimetry experiment, a sample of 2-methylpropan-1-ol (\(C_4H_{10}O\)) was burned completely in air. The heat released was used to heat water in a copper calorimeter. The experimental data collected is shown below:

* Mass of 2-methylpropan-1-ol burned = \(1.48\text{ g}\)
* Mass of water in the calorimeter = \(200.0\text{ g}\)
* Temperature rise of the water = \(45.0^\circ\text{C}\)
* Specific heat capacity of water = \(4.18\text{ J g}^{-1}\text{ K}^{-1}\)

What is the experimental enthalpy change of combustion, \(\Delta H_c\), of 2-methylpropan-1-ol under these conditions?

When the temperature of a gaseous reaction mixture is increased from \(298\text{ K}\) to \(308\text{ K}\), the rate of reaction approximately doubles.

What is the main reason for this significant increase in reaction rate?

An anhydrous Group 2 nitrate, \(X(NO_3)_2\), is heated strongly in a test-tube. A brown gas is evolved and a white solid residue remains.

When \(2.97\text{ g}\) of this anhydrous nitrate is decomposed completely, \(1.20\text{ dm}^3\) of gas (measured at room temperature and pressure, r.t.p.) is produced.

What is the identity of the metal \(X\)?

[Molar volume of a gas at r.t.p. = \(24.0\text{ dm}^3\text{ mol}^{-1}\)]

Four different halogenoalkanes are treated separately with aqueous silver nitrate in ethanol at \(50^\circ\text{C}\). The time taken for a precipitate to appear is recorded.

Which halogenoalkane produces a precipitate in the shortest time?

An organic compound, \(Y\), reacts with 2,4-dinitrophenylhydrazine reagent to form an orange precipitate.

However, \(Y\) does not produce a silver mirror with Tollens' reagent, and does not form a yellow precipitate when warmed with alkaline aqueous iodine.

Which compound could be \(Y\)?

The table shows the first five successive ionisation energies, in \(\text{kJ mol}^{-1}\), of an element \(Z\) in Period 3 of the Periodic Table.

| Ionisation energy | Value / \(\text{kJ mol}^{-1}\) |
| :--- | :--- |
| 1st | 578 |
| 2nd | 1817 |
| 3rd | 2745 |
| 4th | 11578 |
| 5th | 14831 |

Which oxide is formed by element \(Z\)?

Phosphorus pentachloride decomposes reversibly according to the following equation:

\[ PCl_5(g) \rightleftharpoons PCl_3(g) + Cl_2(g) \quad \Delta H = +88\text{ kJ mol}^{-1} \]

An equilibrium mixture is established in a sealed container of volume \(V\) at temperature \(T\).

Which change will increase the value of the equilibrium constant, \(K_c\)?

Which molecule has a permanent, non-zero net dipole moment?

Standard enthalpy changes are fundamental to understanding chemical thermodynamics.

(a) Define the term *standard enthalpy change of reaction*. [2]

(b) An experiment was carried out to determine the enthalpy change for the thermal decomposition of sodium hydrogencarbonate:
\( 2\text{NaHCO}_3(\text{s}) \rightarrow \text{Na}_2\text{CO}_3(\text{s}) + \text{CO}_2(\text{g}) + \text{H}_2\text{O}(\text{l}) \)
This direct enthalpy change cannot be easily measured. Instead, an indirect method was used by reacting both \( \text{NaHCO}_3(\text{s}) \) and anhydrous \( \text{Na}_2\text{CO}_3(\text{s}) \) with excess dilute hydrochloric acid.

(i) In Experiment 1, a student added 3.36 g of \( \text{NaHCO}_3(\text{s}) \) to \( 50.0 \text{ cm}^3 \) of \( 2.00 \text{ mol dm}^{-3} \) \( \text{HCl} \) (an excess) in a polystyrene cup. The temperature of the mixture decreased by \( 4.2 \text{ }^{\circ}\text{C} \).
Calculate the enthalpy change, \( \Delta H_1 \), in \( \text{kJ mol}^{-1} \), for the reaction of 1 mole of \( \text{NaHCO}_3(\text{s}) \) with \( \text{HCl}(\text{aq}) \).
Assume the specific heat capacity of the solution is \( 4.18 \text{ J g}^{-1}\text{ K}^{-1} \) and its density is \( 1.00 \text{ g cm}^{-3} \). Assume the mass of the solution is equal to the mass of the acid solution. Show your working. [4]

(ii) In Experiment 2, the student added 2.12 g of anhydrous \( \text{Na}_2\text{CO}_3(\text{s}) \) to \( 50.0 \text{ cm}^3 \) of \( 2.00 \text{ mol dm}^{-3} \) \( \text{HCl} \) (an excess) in a polystyrene cup. The temperature of the mixture increased by \( 3.8 \text{ }^{\circ}\text{C} \).
Calculate the enthalpy change, \( \Delta H_2 \), in \( \text{kJ mol}^{-1} \), for the reaction of 1 mole of \( \text{Na}_2\text{CO}_3(\text{s}) \) with \( \text{HCl}(\text{aq}) \).
Use the same assumptions as in (b)(i). Show your working. [3]

(c) Construct a Hess's Law cycle relating the decomposition of \( \text{NaHCO}_3(\text{s}) \) to its reactions with hydrochloric acid. Use your cycle and the values from (b)(i) and (b)(ii) to calculate the standard enthalpy change of reaction, \( \Delta H_{\text{r}} \), in \( \text{kJ mol}^{-1} \), for the decomposition of 2 moles of \( \text{NaHCO}_3(\text{s}) \):
\( 2\text{NaHCO}_3(\text{s}) \rightarrow \text{Na}_2\text{CO}_3(\text{s}) + \text{CO}_2(\text{g}) + \text{H}_2\text{O}(\text{l}) \)
Show your working. [4]

(d) State one reason, other than heat loss to the surroundings, why the experimental value of \( \Delta H_{\text{r}} \) determined via this indirect method might differ slightly from the literature value. [1]

(e) Predict the effect on the temperature change in Experiment 2 if a glass beaker were used instead of a polystyrene cup. Explain your answer. [1]

The rate of chemical reactions depends on factors such as temperature, concentration, and the presence of a catalyst.

(a) Define the term *activation energy*. [1]

(b) The Maxwell-Boltzmann distribution shows the distribution of molecular energies in a gas at a constant temperature.

(i) Sketch a Maxwell-Boltzmann distribution curve for a gas at a temperature \( T_1 \). On your sketch, label the axes and mark the activation energy, \( E_{\text{a}} \), of a reaction.
On the same diagram, sketch the curve you would expect for the same gas at a higher temperature, \( T_2 \). [3]

(ii) Use your diagram to explain, in terms of the collision theory, why a relatively small increase in temperature leads to a very large increase in the rate of reaction. [3]

(c) Many industrial processes utilize catalysts to increase reaction rates.

(i) State what is meant by the term *homogeneous catalyst*. [1]

(ii) Explain, by referencing a reaction pathway diagram, how a catalyst increases the rate of reaction. [3]

(d) During a reaction at a constant temperature, the rate of reaction decreases as time progresses.
Explain this observation in terms of collision theory. [2]

(e) Some reactions, such as the reaction between nitrogen and oxygen to form nitrogen monoxide, have extremely high activation energies.

(i) Explain why these reactions are exceptionally slow at room temperature. [1]

(ii) Suggest one source of energy in nature that provides the activation energy for this reaction. [1]

This question is about Period 3 elements and their compounds.

(a) Describe and explain the trend in electronegativity across Period 3 from sodium to chlorine. [3]

(b) Elements in Period 3 react with oxygen to form oxides.

(i) Write equations for the reactions of sodium and phosphorus with oxygen to form sodium oxide, \( \text{Na}_2\text{O} \), and phosphorus(V) oxide, \( \text{P}_4\text{O}_{10} \), respectively. [2]

(ii) State the structure and bonding of silicon dioxide, \( \text{SiO}_2 \), and sulfur trioxide, \( \text{SO}_3 \). [2]

(iii) Write an equation for the reaction of sulfur trioxide with water and state the pH of the resulting solution. [2]

(iv) Aluminium oxide, \( \text{Al}_2\text{O}_3 \), is described as amphoteric. Describe its reactions with both hydrochloric acid and aqueous sodium hydroxide, writing a balanced chemical equation for each reaction. [4]

(c) Silicon tetrachloride, \( \text{SiCl}_4 \), reacts rapidly with water, whereas carbon tetrachloride, \( \text{CCl}_4 \), does not react with water. Explain this difference in reactivity. [2]

Alcohols, carbonyl compounds, and halogenoalkanes are key intermediates in organic synthesis.

(a) Compound A is propan-1-ol, \( \text{CH}_3\text{CH}_2\text{CH}_2\text{OH} \).

(i) State the class of alcohol to which propan-1-ol belongs. [1]

(ii) Propan-1-ol can be dehydrated to form an alkene, Compound D. State a suitable reagent and condition for this reaction, and name the alkene D formed. [2]

(b) Propan-1-ol can be oxidized to form different products depending on the reaction conditions.

(i) State the reagents and conditions required to obtain a good yield of propanal from propan-1-ol. [2]

(ii) State the reagents and conditions required to obtain propanoic acid from propan-1-ol. [2]

(iii) State the color change of the oxidizing agent observed during these oxidation reactions. [1]

(iv) Suggest a chemical test that could distinguish between propanal and propanoic acid. State the reagent and the observations for both compounds. [2]

(c) Propan-1-ol can be converted into 1-chloropropane, Compound E, by reaction with phosphorus trichloride, \( \text{PCl}_3 \).

(i) Write a balanced chemical equation for this reaction. [1]

(ii) Compound E reacts with hot aqueous sodium hydroxide to reform propan-1-ol via an \( \text{S}_{\text{N}}2 \) mechanism.
Draw the mechanism for this reaction. Your diagram must show relevant dipoles, lone pairs, curly arrows, the structure of the transition state, and any charges. [4]

FA 1 is an aqueous solution containing \(8.40\text{ g dm}^{-3}\) of an impure sample of hydrated ethanedioic acid, \(\text{H}_2\text{C}_2\text{O}_4 \cdot 2\text{H}_2\text{O}\). FA 2 is \(0.0200\text{ mol dm}^{-3}\) potassium manganate(VII), \(\text{KMnO}_4\). Dilute sulfuric acid is also provided.

(a) You are to determine the concentration of ethanedioic acid by titration. Pipette \(25.0\text{ cm}^3\) of FA 1 into a conical flask. Use a measuring cylinder to add \(20\text{ cm}^3\) of dilute sulfuric acid. Heat the mixture to approximately \(60^\circ\text{C}\), then titrate with FA 2 until a permanent pale pink color is reached. Repeat the titration until concordant results are achieved. Record your titration results in a standard table. (For calculations, assume your average concordant titre of FA 2 is \(20.00\text{ cm}^3\).)

(b) Calculate the number of moles of \(\text{KMnO}_4\) present in the average titre of FA 2.

(c) Using the equation: \(2\text{MnO}_4^- + 5\text{H}_2\text{C}_2\text{O}_4 + 6\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 10\text{CO}_2 + 8\text{H}_2\text{O}\), calculate the concentration, in \(\text{mol dm}^{-3}\), of ethanedioic acid in FA 1.

(d) Calculate the percentage purity of the \(\text{H}_2\text{C}_2\text{O}_4 \cdot 2\text{H}_2\text{O}\) sample (\(M_r = 126.0\)).

(e) Explain why dilute hydrochloric acid cannot be used to acidify this titration mixture instead of dilute sulfuric acid.

In this experiment, you will determine the enthalpy change of decomposition of calcium carbonate indirectly using Hess's Law. You are provided with:
FA 3: \(3.00\text{ g}\) of calcium carbonate, \(\text{CaCO}_3\) (\(M_r = 100.1\))
FA 4: \(1.40\text{ g}\) of calcium oxide, \(\text{CaO}\) (\(M_r = 56.1\))
Dilute hydrochloric acid, \(\text{HCl}\), of concentration \(2.0\text{ mol dm}^{-3}\) (excess)

(a) Reaction 1: Pour \(50.0\text{ cm}^3\) of the acid into a polystyrene cup. Record the initial temperature \(T_1\). Add all the FA 3, stir continuously, and record the maximum temperature achieved \(T_2\).
(Assume: \(T_1 = 22.0^\circ\text{C}\), \(T_2 = 24.3^\circ\text{C}\).)
Calculate the heat energy released, \(q_1\), in kJ, and the enthalpy change, \(\Delta H_1\), in \(\text{kJ mol}^{-1}\), for the reaction of \(1\text{ mol}\) of \(\text{CaCO}_3\) with \(\text{HCl}\). (Assume density of solution is \(1.00\text{ g cm}^{-3}\) and specific heat capacity \(c = 4.18\text{ J g}^{-1}\text{ K}^{-1}\)).

(b) Reaction 2: Repeat the procedure using \(50.0\text{ cm}^3\) of acid and all the FA 4. Record the initial temperature \(T_3\) and maximum temperature \(T_4\).
(Assume: \(T_3 = 21.5^\circ\text{C}\), \(T_4 = 43.8^\circ\text{C}\).)
Calculate the heat energy released, \(q_2\), in kJ, and the enthalpy change, \(\Delta H_2\), in \(\text{kJ mol}^{-1}\), for the reaction of \(1\text{ mol}\) of \(\text{CaO}\) with \(\text{HCl}\).

(c) Use your calculated values of \(\Delta H_1\) and \(\Delta H_2\) to determine the enthalpy change, \(\Delta H_r\), for the decomposition:
\(\text{CaCO}_3(s) \rightarrow \text{CaO}(s) + \text{CO}_2(g)\)

(d) State one significant source of thermal energy loss in this experiment and suggest a specific practical improvement to reduce it.

You are provided with three solutions: FA 5, FA 6, and FA 7. Each solution contains one cation and one anion from the standard Qualitative Analysis Notes. Carry out the following tests and record your observations.

(a) Test 1: Add aqueous sodium hydroxide, \(\text{NaOH}(aq)\), dropwise then in excess to separate portions of FA 5, FA 6, and FA 7.
Test 2: Add aqueous ammonia, \(\text{NH}_3(aq)\), dropwise then in excess to separate portions of FA 5, FA 6, and FA 7.

(b) Test 3: Add acidified barium nitrate, \(\text{Ba(NO}_3)_2(aq)\), to FA 5. Add acidified silver nitrate, \(\text{AgNO}_3(aq)\), to FA 6. Add sodium hydroxide and a piece of aluminium foil, then heat gently, to FA 7, testing any gas evolved with damp red litmus paper.

(c) Identify the cations and anions in FA 5, FA 6, and FA 7 based on your tests.

(d) Write the ionic equation (including state symbols) for the reaction of FA 6 with acidified silver nitrate.

This question is about Born-Haber cycles and the lattice energy of barium sulfide, BaS. The following standard enthalpy change values are given: Enthalpy change of formation of BaS(s): \(-460\text{ kJ mol}^{-1}\), Enthalpy change of atomisation of Ba(s): \(+180\text{ kJ mol}^{-1}\), First ionisation energy of Ba(g): \(+503\text{ kJ mol}^{-1}\), Second ionisation energy of Ba(g): \(+965\text{ kJ mol}^{-1}\), Enthalpy change of atomisation of S(s): \(+279\text{ kJ mol}^{-1}\), First electron affinity of S(g): \(-200\text{ kJ mol}^{-1}\), Second electron affinity of S(g): \(+640\text{ kJ mol}^{-1}\). (a) Define the term lattice energy. [2] (b) Use the given data to calculate the lattice energy of BaS(s). Show your working. [3] (c) Suggest how the lattice energy of MgS(s) compares with that of BaS(s). Explain your answer. [3] (d) Explain why the second electron affinity of sulfur is endothermic, whereas the first electron affinity is exothermic. [3]

An electrochemical cell is set up consisting of a standard \(\text{Fe}^{3+}(\text{aq})/\text{Fe}^{2+}(\text{aq})\) half-cell connected to a non-standard \(\text{Ag}^{+}(\text{aq})/\text{Ag}(\text{s})\) half-cell. The standard electrode potentials are: \(\text{Fe}^{3+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Fe}^{2+}(\text{aq})\) \(E^{\ominus} = +0.77\text{ V}\); \(\text{Ag}^{+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Ag}(\text{s})\) \(E^{\ominus} = +0.80\text{ V}\). In the non-standard silver half-cell, \([\text{Ag}^{+}] = 0.025\text{ mol dm}^{-3}\) at \(298\text{ K}\). (a) Draw a fully labelled diagram of this electrochemical cell. Include details of the electrodes, solutions, and salt bridge. [4] (b) State the function of the salt bridge and suggest a suitable electrolyte to prepare it. [2] (c) Calculate the electrode potential, \(E\), of the silver half-cell under these non-standard conditions using the Nernst equation. [2] (d) Calculate the overall cell potential, \(E_{\text{cell}}\), under these conditions and write the overall ionic equation for the cell reaction. [3]

Nitric oxide, \(\text{NO}\), reacts with hydrogen gas according to the equation: \(2\text{NO}(\text{g}) + 2\text{H}_2(\text{g}) \rightarrow \text{N}_2(\text{g}) + 2\text{H}_2\text{O}(\text{g})\). The following initial rate data were obtained at \(800\text{ K}\): Experiment 1: \([\text{NO}] = 0.010\text{ mol dm}^{-3}\), \([\text{H}_2] = 0.010\text{ mol dm}^{-3}\), Initial Rate = \(1.2 \times 10^{-5}\text{ mol dm}^{-3}\text{ s}^{-1}\); Experiment 2: \([\text{NO}] = 0.020\text{ mol dm}^{-3}\), \([\text{H}_2] = 0.010\text{ mol dm}^{-3}\), Initial Rate = \(4.8 \times 10^{-5}\text{ mol dm}^{-3}\text{ s}^{-1}\); Experiment 3: \([\text{NO}] = 0.010\text{ mol dm}^{-3}\), \([\text{H}_2] = 0.020\text{ mol dm}^{-3}\), Initial Rate = \(2.4 \times 10^{-5}\text{ mol dm}^{-3}\text{ s}^{-1}\). (a) Deduce the order of reaction with respect to both \(\text{NO}\) and \(\text{H}_2\). Show your reasoning. [4] (b) Write the rate equation for the reaction, and calculate the value of the rate constant, \(k\), including its units. [3] (c) Suggest a multi-step reaction mechanism consistent with this rate equation, indicating which step is the rate-determining step. [4]

Aqueous copper(II) ions exist as \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\). When excess aqueous ammonia is added, ligand substitution occurs: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{NH}_3(\text{aq}) \rightleftharpoons [\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}(\text{aq}) + 4\text{H}_2\text{O}(\text{l})\) for which \(K_{\text{stab}} = 2.1 \times 10^{13}\text{ dm}^{12}\text{ mol}^{-4}\). When 1,2-diaminoethane (en) is added, it forms a tris-chelate complex: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 3\text{en}(\text{aq}) \rightleftharpoons [\text{Cu}(\text{en})_3]^{2+}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\) for which \(K_{\text{stab}} = 2.0 \times 10^{18}\text{ dm}^9\text{ mol}^{-3}\). (a) Write the expression for the stability constant, \(K_{\text{stab}}\), for the formation of \([\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}\). [2] (b) Explain why the tris-chelate complex is much more stable than the ammonia complex, referencing entropy changes, \(\Delta S^{\ominus}\). [4] (c) Describe the color change when excess ammonia is added to aqueous copper(II) sulfate, and state the geometry of the resulting complex. [3] (d) Draw the structure of one of the optical isomers of \([\text{Cu}(\text{en})_3]^{2+}\), showing its three-dimensional shape. [2]

An organic acid, X, is soluble in both water and ethoxyethane (ether). The partition coefficient, \(K_{\text{pc}}\), of X between ether and water is defined as: \(K_{\text{pc}} = \frac{[\text{X}]_{\text{ether}}}{[\text{X}]_{\text{water}}} = 4.50\). A solution is prepared by dissolving \(10.0\text{ g}\) of X in \(100\text{ cm}^3\) of water. (a) State the definition of partition coefficient. [2] (b) Calculate the mass of X extracted from this aqueous solution if it is shaken with: (i) One single portion of \(60\text{ cm}^3\) of ether. [3] (ii) Two successive portions of \(30\text{ cm}^3\) of ether. [4] (c) Compare the efficiency of the extraction in (b)(i) and (b)(ii) and state a general rule for solvent extraction. [2]

A synthesis of 4-aminobenzoic acid (PABA) starting from methylbenzene is outlined below: Methylbenzene \(\xrightarrow{\text{Step 1}} \) 4-nitromethylbenzene \(\xrightarrow{\text{Step 2}} \) 4-nitrobenzoic acid \(\xrightarrow{\text{Step 3}} \) 4-aminobenzoic acid. (a) Give the reagents and conditions for Step 1. State the role of sulfuric acid in this step and write an equation for the formation of the electrophile. [4] (b) Draw the structures of 4-nitromethylbenzene and 4-nitrobenzoic acid. [2] (c) State the reagents and conditions required for: (i) Step 2 [2] (ii) Step 3 [1] (d) State how methylbenzene and 4-aminobenzoic acid can be distinguished using a simple chemical test, describing the observations. [2]

Calcium carbonate decomposes at high temperatures according to the following equation: \(\text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g})\). The following thermodynamic data at \(298\text{ K}\) are given: \(\Delta H_{\text{f}}^{\ominus}[\text{CaCO}_3(\text{s})] = -1207\text{ kJ mol}^{-1}\), \(\Delta H_{\text{f}}^{\ominus}[\text{CaO}(\text{s})] = -635\text{ kJ mol}^{-1}\), \(\Delta H_{\text{f}}^{\ominus}[\text{CO}_2(\text{g})] = -394\text{ kJ mol}^{-1}\). Standard entropies, \(S^{\ominus}\): \(\text{CaCO}_3(\text{s}) = 92.9\text{ J K}^{-1}\text{ mol}^{-1}\), \(\text{CaO}(\text{s}) = 39.7\text{ J K}^{-1}\text{ mol}^{-1}\), \(\text{CO}_2(\text{g}) = 213.6\text{ J K}^{-1}\text{ mol}^{-1}\). (a) Define the term standard entropy of a substance. [2] (b) Calculate the standard enthalpy change, \(\Delta H^{\ominus}\), for this reaction at \(298\text{ K}\). [2] (c) Calculate the standard entropy change, \(\Delta S^{\ominus}\), for this reaction at \(298\text{ K}\). [2] (d) Calculate the standard Gibbs free energy change, \(\Delta G^{\ominus}\), for this reaction at \(298\text{ K}\). State and explain whether the reaction is spontaneous at this temperature. [3] (e) Determine the temperature (in K) above which this decomposition reaction becomes feasible. [2]

This question is about transition elements and their physical and chemical properties. (a) Explain in detail why transition metal complex ions are typically colored. Your answer should make reference to: d-orbitals and their splitting, the role of ligands, energy absorption and transition of electrons. [5] (b) When concentrated hydrochloric acid is added to a pale blue solution of aqueous copper(II) sulfate, the solution changes color to yellow-green. Write an ionic equation for this reaction and explain the color change in terms of ligand exchange and the size of the ligands. [4] (c) Suggest why solutions of \(\text{Sc}^{3+}\) and \(\text{Zn}^{2+}\) ions are completely colorless. [2]

(a) (i) Define the term *stability constant*, \(K_{\text{stab}}\), of a complex ion.

(ii) Write an expression for the stability constant, \(K_{\text{stab}}\), of the following ligand exchange reaction, and state its units.

\[[\text{Fe}(\text{H}_2\text{O})_6]^{3+}(\text{aq}) + \text{SCN}^-(\text{aq}) \rightleftharpoons [\text{Fe}(\text{SCN})(\text{H}_2\text{O})_5]^{2+}(\text{aq}) + \text{H}_2\text{O}(\text{l})\]

(b) Cobalt(II) ions exist in aqueous solution as the pink hexaaquacobalt(II) ion, \[[\text{Co}(\text{H}_2\text{O})_6]^{2+}\]. When concentrated hydrochloric acid is added, a ligand exchange reaction occurs to form a blue tetrahedral complex:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{Cl}^-(\text{aq}) \rightleftharpoons [\text{CoCl}_4]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\]

The stability constant, \(K_{\text{stab}}\), for this equilibrium is \(3.2 \times 10^1\text{ dm}^{12}\text{ mol}^{-4}\).

At equilibrium, a solution of cobalt(II) chloride in concentrated hydrochloric acid has a total concentration of cobalt species of \(0.10\text{ mol dm}^{-3}\) and a free chloride ion concentration, \[[\text{Cl}^-]\], of \(2.5\text{ mol dm}^{-3}\).

Calculate the equilibrium concentration of the remaining hexaaquacobalt(II) ions, \[[[\text{Co}(\text{H}_2\text{O})_6]^{2+}]\).

(c) The bidentate ligand 1,2-diaminoethane is represented by \(\text{en}\).

The reaction of hexaaquacobalt(II) with \(\text{en}\) has a very large stability constant:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 3\text{en} \rightleftharpoons [\text{Co}(\text{en})_3]^{2+} + 6\text{H}_2\text{O} \quad \log K_{\text{stab}} = 13.8\]

In contrast, the corresponding reaction with the monodentate ligand ammonia, \(\text{NH}_3\), has a much lower stability constant:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 6\text{NH}_3 \rightleftharpoons [\text{Co}(\text{NH}_3)_6]^{2+} + 6\text{H}_2\text{O} \quad \log K_{\text{stab}} = 5.1\]

Explain this difference in the stability of the two complexes in terms of entropy change, \(\Delta S^\theta\), and Gibbs free energy change, \(\Delta G^\theta\).

(a) (i) Define the term *stability constant*, \(K_{\text{stab}}\), of a complex ion.

(ii) Write an expression for the stability constant, \(K_{\text{stab}}\), of the following ligand exchange reaction, and state its units.

\[[\text{Fe}(\text{H}_2\text{O})_6]^{3+}(\text{aq}) + \text{SCN}^-(\text{aq}) \rightleftharpoons [\text{Fe}(\text{SCN})(\text{H}_2\text{O})_5]^{2+}(\text{aq}) + \text{H}_2\text{O}(\text{l})\]

(b) Cobalt(II) ions exist in aqueous solution as the pink hexaaquacobalt(II) ion, \[[\text{Co}(\text{H}_2\text{O})_6]^{2+}\]. When concentrated hydrochloric acid is added, a ligand exchange reaction occurs to form a blue tetrahedral complex:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{Cl}^-(\text{aq}) \rightleftharpoons [\text{CoCl}_4]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\]

The stability constant, \(K_{\text{stab}}\), for this equilibrium is \(3.2 \times 10^1\text{ dm}^{12}\text{ mol}^{-4}\).

At equilibrium, a solution of cobalt(II) chloride in concentrated hydrochloric acid has a total concentration of cobalt species of \(0.10\text{ mol dm}^{-3}\) and a free chloride ion concentration, \[[\text{Cl}^-]\], of \(2.5\text{ mol dm}^{-3}\).

Calculate the equilibrium concentration of the remaining hexaaquacobalt(II) ions, \[[[\text{Co}(\text{H}_2\text{O})_6]^{2+}]\).

(c) The bidentate ligand 1,2-diaminoethane is represented by \(\text{en}\).

The reaction of hexaaquacobalt(II) with \(\text{en}\) has a very large stability constant:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 3\text{en} \rightleftharpoons [\text{Co}(\text{en})_3]^{2+} + 6\text{H}_2\text{O} \quad \log K_{\text{stab}} = 13.8\]

In contrast, the corresponding reaction with the monodentate ligand ammonia, \(\text{NH}_3\), has a much lower stability constant:

\[[\text{Co}(\text{H}_2\text{O})_6]^{2+} + 6\text{NH}_3 \rightleftharpoons [\text{Co}(\text{NH}_3)_6]^{2+} + 6\text{H}_2\text{O} \quad \log K_{\text{stab}} = 5.1\]

Explain this difference in the stability of the two complexes in terms of entropy change, \(\Delta S^\theta\), and Gibbs free energy change, \(\Delta G^\theta\).

A student wants to determine the activation energy, \(E_a\), of the reaction between peroxodisulfate ions and iodide ions:
\[ \text{S}_2\text{O}_8^{2-}(aq) + 2\text{I}^-(aq) \rightarrow 2\text{SO}_4^{2-}(aq) + \text{I}_2(aq) \]

The rate of this reaction can be monitored using a 'clock reaction' by adding a small, fixed amount of sodium thiosulfate, \(\text{Na}_2\text{S}_2\text{O}_3\), and starch indicator.

The thiosulfate ions instantly reduce any iodine formed back to iodide:
\[ \text{I}_2(aq) + 2\text{S}_2\text{O}_3^{2-}(aq) \rightarrow 2\text{I}^-(aq) + \text{S}_4\text{O}_6^{2-}(aq) \]
Once all the thiosulfate is consumed, the remaining iodine reacts with starch, turning the solution blue-black.
The time, \(t\), taken for the blue-black color to appear is measured. The rate of reaction can be assumed to be proportional to \(1/t\).

**(a)** Identify:
(i) the independent variable.
(ii) the dependent variable.
(iii) two variables that must be kept constant.
[3 marks]

**(b)** Design a detailed experimental procedure to investigate how the rate of this reaction varies with temperature in the range \(20\ ^\circ\text{C}\) to \(60\ ^\circ\text{C}\). Your plan should include:
- details of how the temperature will be varied and controlled.
- the volumes and concentrations of reagents you would use (suggested reagents: \(0.100\text{ mol dm}^{-3}\ \text{K}_2\text{S}_2\text{O}_8\), \(0.200\text{ mol dm}^{-3}\ \text{KI}\), \(0.0100\text{ mol dm}^{-3}\ \text{Na}_2\text{S}_2\text{O}_3\), and starch solution). Total volume of the reaction mixture should be \(50.0\text{ cm}^3\).
- how you will ensure that the reagents are at the correct temperature before mixing.
[5 marks]

**(c)** The Arrhenius equation is:
\[ \ln k = -\frac{E_a}{RT} + \text{constant} \]
Since rate \(\propto 1/t\), the rate constant \(k\) can be approximated by \(1/t\), giving:
\[ \ln\left(\frac{1}{t}\right) = -\frac{E_a}{RT} + \text{constant} \]
Draw a template table that the student would use to record raw and processed data to plot a straight-line graph to determine \(E_a\). Include appropriate column headings with units.
[3 marks]

**(d)** Describe how the student would use the gradient of the graph to determine the activation energy, \(E_a\), in \(\text{kJ mol}^{-1}\). State the value of any constant(s) used.
[2 marks]

**(e)** A student carried out the experiment and obtained the following data point:
At temperature \(\theta = 45.0\ ^\circ\text{C}\), the time taken for the blue-black color to appear was \(t = 34.0\text{ s}\).
Calculate the value of:
(i) \(1/T\) in \(\text{K}^{-1}\)
(ii) \(\ln(1/t)\)
Show your working and express your answers to an appropriate number of significant figures.
[2 marks]

A student wants to determine the enthalpy change of hydration of anhydrous magnesium chloride, \(\text{MgCl}_2(s)\):
\[ \text{MgCl}_2(s) + 6\text{H}_2\text{O}(l) \rightarrow \text{MgCl}_2\cdot 6\text{H}_2\text{O}(s) \quad \Delta H_{hyd} \]

To do this, the student plans to determine the enthalpy changes of solution of both anhydrous magnesium chloride and hydrated magnesium chloride.

**(a)** The student first plans to measure the enthalpy change of solution of anhydrous \(\text{MgCl}_2\).
(i) Describe how the student would weigh the anhydrous \(\text{MgCl}_2\) to ensure the exact mass added to the water is known. [1 mark]
(ii) Show by calculation that if \(50.0\text{ cm}^3\) of water is used, a mass of approximately \(1.50\text{ g}\) of anhydrous \(\text{MgCl}_2\) will produce a temperature rise of at least \(5.0\ ^\circ\text{C}\).
[Assume density of water is \(1.00\text{ g cm}^{-3}\), specific heat capacity is \(4.18\text{ J g}^{-1}\text{ K}^{-1}\), \(M_r\) of \(\text{MgCl}_2 = 95.3\text{ g mol}^{-1}\), and approximate \(\Delta H_{sol} \approx -155\text{ kJ mol}^{-1}\)]. [3 marks]
(iii) In the experiment, the mixture is stirred continuously, and the temperature is recorded at regular intervals. Explain why the maximum recorded temperature is not the theoretical maximum temperature change, and state how the student can find the correct maximum temperature change. [2 marks]

**(b)** Draw a Hess's Law cycle that allows the calculation of the enthalpy change of hydration, \(\Delta H_{hyd}\), of anhydrous magnesium chloride using the enthalpy changes of solution of anhydrous and hydrated magnesium chloride. Write an equation relating these three enthalpy changes. [2 marks]

**(c)** In the second experiment, the student dissolves \(10.15\text{ g}\) of hydrated magnesium chloride, \(\text{MgCl}_2\cdot 6\text{H}_2\text{O}\), in \(50.0\text{ cm}^3\) of water in a polystyrene cup.
The temperature of the water decreased by \(1.80\ ^\circ\text{C}\).
(i) Calculate the number of moles of \(\text{MgCl}_2\cdot 6\text{H}_2\text{O}\) used. [\(M_r\) of \(\text{MgCl}_2\cdot 6\text{H}_2\text{O} = 203.3\text{ g mol}^{-1}\)] [1 mark]
(ii) Calculate the enthalpy change of solution of \(\text{MgCl}_2\cdot 6\text{H}_2\text{O}(s)\) in \(\text{kJ mol}^{-1}\). Include the correct sign. [3 marks]

**(d)** Using your answer to **(c)(ii)** and a value of \(\Delta H_{sol}(\text{MgCl}_2(s)) = -160.0\text{ kJ mol}^{-1}\), calculate the enthalpy change of hydration, \(\Delta H_{hyd}\), of anhydrous magnesium chloride. [2 marks]

**(e)** A second student used a weighing bottle containing anhydrous magnesium chloride. After transferring the solid, they noticed that the inside of the weighing bottle was wet.
Predict and explain the effect of this on the calculated enthalpy change of solution of the anhydrous magnesium chloride. [1 mark]