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A \( 10\text{ cm}^3 \) sample of a gaseous hydrocarbon was exploded with \( 80\text{ cm}^3 \) (an excess) of oxygen. After cooling to room temperature, the remaining gas volume was \( 65\text{ cm}^3 \). This gas mixture was then shaken with aqueous sodium hydroxide, and the volume decreased to \( 35\text{ cm}^3 \). All gas volumes are measured at room temperature and pressure. What is the molecular formula of the hydrocarbon?

A mixture of anhydrous calcium carbonate (\( \text{CaCO}_3 \), \( M_r = 100.1 \)) and anhydrous magnesium carbonate (\( \text{MgCO}_3 \), \( M_r = 84.3 \)) has a mass of \( 10.00\text{ g} \). The mixture is heated strongly until there is no further change in mass. The remaining solid residue has a mass of \( 5.20\text{ g} \). What is the percentage by mass of \( \text{MgCO}_3 \) in the original mixture?

Propylbenzene, \( \text{C}_6\text{H}_5\text{CH}_2\text{CH}_2\text{CH}_3 \), is reacted with chlorine in the presence of ultraviolet light at boiling point. Which chlorinated organic compound is formed in the greatest yield as the major mono-chlorinated product?

In the electrophilic substitution of benzene to form nitrobenzene, a reaction intermediate is formed. Which statement about this intermediate is correct?

The standard electrode potentials for several half-reactions are shown below. \( \text{Fe}^{3+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Fe}^{2+}(\text{aq}) \quad E^\theta = +0.77\text{ V} \) \( \text{I}_2(\text{aq}) + 2\text{e}^- \rightleftharpoons 2\text{I}^-(\text{aq}) \quad E^\theta = +0.54\text{ V} \) \( \text{Fe}^{2+}(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{Fe}(\text{s}) \quad E^\theta = -0.44\text{ V} \) \( \text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{Cu}(\text{s}) \quad E^\theta = +0.34\text{ V} \) Which of the following reactions is feasible under standard conditions?

A standard cell is set up using the two half-cells: \( \text{Ag}^+(\text{aq}) + \text{e}^- \rightleftharpoons \text{Ag}(\text{s}) \quad E^\theta = +0.80\text{ V} \) \( \text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{Cu}(\text{s}) \quad E^\theta = +0.34\text{ V} \) The overall cell equation is: \( \text{Cu}(\text{s}) + 2\text{Ag}^+(\text{aq}) \rightarrow \text{Cu}^{2+}(\text{aq}) + 2\text{Ag}(\text{s}) \quad E^\theta_{\text{cell}} = +0.46\text{ V} \) The concentration of \( \text{Ag}^+(\text{aq}) \) is decreased from \( 1.0\text{ mol dm}^{-3} \) to \( 0.001\text{ mol dm}^{-3} \) while keeping the concentration of \( \text{Cu}^{2+}(\text{aq}) \) at \( 1.0\text{ mol dm}^{-3} \) and temperature at \( 298\text{ K} \). Which statement correctly describes the effect of this change on the electrode potential of the silver half-cell, \( E_{\text{Ag}^+/\text{Ag}} \), and the cell potential, \( E_{\text{cell}} \)?

The initial rate of the reaction \( 2\text{A} + \text{B} + 2\text{C} \rightarrow \text{products} \) was measured at constant temperature: Experiment 1: \( [\text{A}] = 0.10 \), \( [\text{B}] = 0.10 \), \( [\text{C}] = 0.10 \); Rate = \( 1.2 \times 10^{-3} \text{ mol dm}^{-3}\text{ s}^{-1} \). Experiment 2: \( [\text{A}] = 0.20 \), \( [\text{B}] = 0.10 \), \( [\text{C}] = 0.10 \); Rate = \( 2.4 \times 10^{-3} \text{ mol dm}^{-3}\text{ s}^{-1} \). Experiment 3: \( [\text{A}] = 0.10 \), \( [\text{B}] = 0.20 \), \( [\text{C}] = 0.10 \); Rate = \( 4.8 \times 10^{-3} \text{ mol dm}^{-3}\text{ s}^{-1} \). Experiment 4: \( [\text{A}] = 0.10 \), \( [\text{B}] = 0.10 \), \( [\text{C}] = 0.20 \); Rate = \( 1.2 \times 10^{-3} \text{ mol dm}^{-3}\text{ s}^{-1} \). What are the value and units of the rate constant, \( k \), for this reaction?

The reaction between nitrogen monoxide and hydrogen has the overall equation: \( 2\text{NO(g)} + 2\text{H}_2\text{(g)} \rightarrow \text{N}_2\text{(g)} + 2\text{H}_2\text{O(g)} \) and experimentally determined rate equation: \( \text{Rate} = k[\text{NO}]^2[\text{H}_2] \). A proposed mechanism is: Step 1: \( 2\text{NO(g)} \rightleftharpoons \text{N}_2\text{O}_2\text{(g)} \) (fast equilibrium); Step 2: \( \text{N}_2\text{O}_2\text{(g)} + \text{H}_2\text{(g)} \rightarrow \text{N}_2\text{O(g)} + \text{H}_2\text{O(g)} \) (slow); Step 3: \( \text{N}_2\text{O(g)} + \text{H}_2\text{(g)} \rightarrow \text{N}_2\text{(g)} + \text{H}_2\text{O(g)} \) (fast). Which statement about this mechanism is correct?

A 10.0 g sample of a mixture containing only calcium carbonate, \(CaCO_3\), and calcium oxide, \(CaO\), is heated strongly until constant mass is achieved. The mass of the solid residue remaining after heating is 7.8 g. What is the percentage by mass of \(CaCO_3\) in the original mixture? [Ar values: \(Ca = 40.1\), \(C = 12.0\), \(O = 16.0\)]

A piece of magnesium ribbon of mass \(0.120\text{ g}\) is added to \(50.0\text{ cm}^3\) of \(0.0500\text{ mol dm}^{-3}\) hydrochloric acid, \(HCl(aq)\). What is the maximum volume of hydrogen gas, in \(\text{cm}^3\), collected at room temperature and pressure (r.t.p.)? [Molar volume of gas at r.t.p. = \(24.0\text{ dm}^3\text{ mol}^{-1}\), \(A_r(Mg) = 24.3\)]

Methylbenzene is treated with \(Cl_2\) in the presence of anhydrous \(AlCl_3\) to produce a major monosubstituted product X. Separately, methylbenzene is heated with alkaline \(KMnO_4\) and then acidified to yield compound Y. When Y is treated with \(Cl_2\) in the presence of anhydrous \(AlCl_3\), the major monosubstituted product Z is obtained. What are the positions of the chlorine atom relative to the pre-existing substituent in X and Z?

Under standard nitration conditions, methylbenzene reacts significantly faster than benzene. Which statement correctly explains this observation?

Use the standard electrode potentials provided below to determine which of the reactions will NOT occur spontaneously under standard conditions: \(Fe^{3+}(aq) + e^- \rightleftharpoons Fe^{2+}(aq)\) \(E^\ominus = +0.77\text{ V}\); \(I_2(aq) + 2e^- \rightleftharpoons 2I^-(aq)\) \(E^\ominus = +0.54\text{ V}\); \(Br_2(aq) + 2e^- \rightleftharpoons 2Br^-(aq)\) \(E^\ominus = +1.07\text{ V}\); \(SO_4^{2-}(aq) + 4H^+(aq) + 2e^- \rightleftharpoons SO_2(g) + 2H_2O(l)\) \(E^\ominus = +0.17\text{ V}\).

An electrochemical cell is set up under standard conditions using the half-cells: \(Ag^+(aq) + e^- \rightleftharpoons Ag(s)\) \(E^\ominus = +0.80\text{ V}\) and \(Cu^{2+}(aq) + 2e^- \rightleftharpoons Cu(s)\) \(E^\ominus = +0.34\text{ V}\). If the concentration of \(Ag^+(aq)\) in the silver half-cell is decreased to \(0.01\text{ mol dm}^{-3}\) while all other conditions are kept standard, how do the electrode potential of the silver half-cell and the overall cell potential change?

For the reaction \(P + Q + R \rightarrow\text{products}\), initial rate experiments at constant temperature yielded the following data: Run 1: \([P] = 0.10\), \([Q] = 0.10\), \([R] = 0.10\text{ mol dm}^{-3}\), Rate = \(1.2 \times 10^{-3}\text{ mol dm}^{-3}\text{ s}^{-1}\); Run 2: \([P] = 0.20\), \([Q] = 0.10\), \([R] = 0.10\text{ mol dm}^{-3}\), Rate = \(2.4 \times 10^{-3}\text{ mol dm}^{-3}\text{ s}^{-1}\); Run 3: \([P] = 0.10\), \([Q] = 0.20\), \([R] = 0.10\text{ mol dm}^{-3}\), Rate = \(4.8 \times 10^{-3}\text{ mol dm}^{-3}\text{ s}^{-1}\); Run 4: \([P] = 0.10\), \([Q] = 0.10\), \([R] = 0.20\text{ mol dm}^{-3}\), Rate = \(1.2 \times 10^{-3}\text{ mol dm}^{-3}\text{ s}^{-1}\). What is the overall order of the reaction and the unit of the rate constant, \(k\)?

A first-order reaction has a rate constant \(k = 4.50 \times 10^{-4}\text{ s}^{-1}\). If the initial concentration of the reactant is \(0.800\text{ mol dm}^{-3}\), how long will it take for the concentration of the reactant to fall to \(0.100\text{ mol dm}^{-3}\)?

A sample of \( 0.250\text{ g} \) of a carbonate of a Group 2 metal, \( \text{MCO}_3 \), reacted completely with excess dilute hydrochloric acid. The carbon dioxide gas evolved was collected and measured at room temperature and pressure (r.t.p.) to be \( 60.0\text{ cm}^3 \). What is the identity of the Group 2 metal, \( \text{M} \)? [Take the molar volume of a gas at r.t.p. as \( 24.0\text{ dm}^3\text{ mol}^{-1} \).]

When 10 cm\(^3\) of a gaseous hydrocarbon is completely combusted in 100 cm\(^3\) of oxygen (which is in excess), the total volume of gas remaining after cooling to room temperature and pressure is 85 cm\(^3\). On passing this gaseous mixture through concentrated aqueous sodium hydroxide, the volume decreases to 55 cm\(^3\). What is the molecular formula of the hydrocarbon? [Assume all gas volumes are measured at the same temperature and pressure.]

Methylbenzene reacts with a mixture of concentrated nitric acid and concentrated sulfuric acid at \( 55\text{ }^\circ\text{C} \) to form a mixture of monosubstituted nitro-isomers. Which statement correctly explains why the reaction of methylbenzene is faster than that of benzene under the same conditions?

A student carries out the halogenation of methylbenzene using chlorine gas. Depending on the conditions used, different products are formed. Under which set of conditions is (chloromethyl)benzene, \( \text{C}_6\text{H}_5\text{CH}_2\text{Cl} \), the major organic product?

The standard electrode potentials for three redox systems are shown below.
\( \text{Fe}^{3+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Fe}^{2+}(\text{aq}) \quad E^\ominus = +0.77\text{ V} \)
\( \text{I}_2(\text{aq}) + 2\text{e}^- \rightleftharpoons 2\text{I}^-(\text{aq}) \quad E^\ominus = +0.54\text{ V} \)
\( \text{S}_4\text{O}_6^{2-}(\text{aq}) + 2\text{e}^- \rightleftharpoons 2\text{S}_2\text{O}_3^{2-}(\text{aq}) \quad E^\ominus = +0.09\text{ V} \)

Which species is the strongest reducing agent under standard conditions?

An electrochemical cell is set up under standard conditions using a standard hydrogen electrode (SHE) as one half-cell and a standard copper half-cell, \( \text{Cu}^{2+}(\text{aq}) | \text{Cu}(\text{s}) \), as the other.
\( \text{Cu}^{2+}(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{Cu}(\text{s}) \quad E^\ominus = +0.34\text{ V} \)

Which statement about this cell is correct when it is operating to deliver current?

For the reaction \( 2\text{A} + \text{B} \rightarrow \text{C} + \text{D} \), the rate equation is determined experimentally to be:
\[ \text{rate} = k[\text{A}][\text{B}]^2 \]

When the initial concentration of \( \text{A} \) is doubled and the initial concentration of \( \text{B} \) is halved, how does the initial rate of reaction change?

The table below shows initial rate data for the reaction:
\[ \text{P} + \text{Q} + \text{R} \rightarrow \text{products} \]

At a constant temperature:

| Experiment | Initial [P] / mol dm\(^{-3}\) | Initial [Q] / mol dm\(^{-3}\) | Initial [R] / mol dm\(^{-3}\) | Initial rate / mol dm\(^{-3}\) s\(^{-1}\) |
| :---: | :---: | :---: | :---: | :---: |
| 1 | 0.10 | 0.10 | 0.10 | \( 1.2 \times 10^{-3} \) |
| 2 | 0.20 | 0.10 | 0.10 | \( 2.4 \times 10^{-3} \) |
| 3 | 0.10 | 0.20 | 0.10 | \( 4.8 \times 10^{-3} \) |
| 4 | 0.10 | 0.10 | 0.20 | \( 1.2 \times 10^{-3} \) |

What is the overall order of the reaction?

\( 10\text{ cm}^3 \) of a gaseous hydrocarbon is reacted with \( 70\text{ cm}^3 \) of oxygen (an excess). After complete combustion and cooling to room temperature, the total volume of dry gas remaining is \( 50\text{ cm}^3 \). Passing this remaining gas through aqueous sodium hydroxide reduces its volume to \( 20\text{ cm}^3 \).

What is the molecular formula of the hydrocarbon?

A \( 5.00\text{ g} \) sample of a mixture containing anhydrous sodium carbonate (\( \text{Na}_2\text{CO}_3 \), \( M_{\text{r}} = 106.0 \)) and sodium chloride (\( \text{NaCl} \), \( M_{\text{r}} = 58.5 \)) is reacted completely with excess dilute hydrochloric acid. The volume of carbon dioxide gas collected at room temperature and pressure (r.t.p.) is \( 840\text{ cm}^3 \).

What is the percentage by mass of sodium chloride in the original sample?

[Assume 1 mol of gas occupies \( 24.0\text{ dm}^3 \) at r.t.p.]

Which of the following compounds undergoes electrophilic substitution of the aromatic ring at the fastest rate when reacted with chlorine in the presence of an anhydrous iron(III) chloride catalyst?

What are the major organic products formed when propylbenzene is treated under the following reaction conditions?

Reaction 1: heating with alkaline \( \text{KMnO}_4 \), followed by acidification
Reaction 2: boiling with \( \text{Cl}_2 \) in the presence of UV light

The following standard electrode potentials are given:

\( \text{Fe}^{3+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Fe}^{2+}(\text{aq}) \quad E^\ominus = +0.77\text{ V} \)
\( \text{I}_2(\text{aq}) + 2\text{e}^- \rightleftharpoons 2\text{I}^-(\text{aq}) \quad E^\ominus = +0.54\text{ V} \)
\( \text{S}_4\text{O}_6^{2-}(\text{aq}) + 2\text{e}^- \rightleftharpoons 2\text{S}_2\text{O}_3^{2-}(\text{aq}) \quad E^\ominus = +0.09\text{ V} \)
\( \text{Fe}^{2+}(\text{aq}) + 2\text{e}^- \rightleftharpoons \text{Fe}(\text{s}) \quad E^\ominus = -0.44\text{ V} \)

Which reaction is thermodynamically feasible under standard conditions?

A half-cell is set up containing a silver electrode in contact with a solution of silver ions, \( \text{Ag}^+(\text{aq}) \). The standard electrode potential is:

\( \text{Ag}^+(\text{aq}) + \text{e}^- \rightleftharpoons \text{Ag}(\text{s}) \quad E^\ominus = +0.80\text{ V} \)

At \( 298\text{ K} \, the measured electrode potential of this half-cell is \) +0.68\text{ V} \). What is the concentration of \( \text{Ag}^+(\text{aq}) \) in the solution?

The initial rates of the reaction \( \text{A} + 2\text{B} + \text{C} \rightarrow \text{Products} \) were measured at different concentrations of reactants at a constant temperature:

| Experiment | \( [\text{A}] / \text{mol dm}^{-3} \) | \( [\text{B}] / \text{mol dm}^{-3} \) | \( [\text{C}] / \text{mol dm}^{-3} \) | Initial rate / \( \text{mol dm}^{-3}\text{ s}^{-1} \) |
|---|---|---|---|---|
| 1 | 0.10 | 0.10 | 0.10 | \( 1.2 \times 10^{-3} \) |
| 2 | 0.20 | 0.10 | 0.10 | \( 2.4 \times 10^{-3} \) |
| 3 | 0.10 | 0.20 | 0.10 | \( 4.8 \times 10^{-3} \) |
| 4 | 0.10 | 0.10 | 0.20 | \( 1.2 \times 10^{-3} \) |

What are the units of the rate constant, \( k \)?

A chemical reaction is first-order with respect to a reactant P. When the initial concentration of P is \( 1.60\text{ mol dm}^{-3} \), it takes exactly \( 40\text{ minutes} \) for the concentration of P to fall to \( 0.20\text{ mol dm}^{-3} \).

Under the same conditions, how long will it take for the concentration of P to fall from \( 0.80\text{ mol dm}^{-3} \) to \( 0.10\text{ mol dm}^{-3} \)?

A 1.20 g sample of a Group 2 metal carbonate, \(MCO_3\), is heated strongly until it completely decomposes to form the metal oxide and carbon dioxide. The volume of carbon dioxide gas collected at room temperature and pressure (r.t.p.) is 288 cm\(^3\). What is the identity of the metal \(M\)? [Molar volume of gas at r.t.p. = 24.0 dm\(^3\) mol\(^{-1\)}]

A 25.0 cm\(^3\) sample of 0.0200 mol dm\(^{-3}\) of a metal ion \(M^{n+}\) reacts completely with exactly 10.0 cm\(^3\) of 0.0200 mol dm\(^{-3}\) of acidified potassium manganate(VII), \(KMnO_4\). In this reaction, the manganate(VII) ions are reduced to \(Mn^{2+}\) ions, and the metal ions are oxidized to \(M^{5+}\) ions. What is the initial oxidation state, \(n\), of the metal ion?

When a dimethylbenzene isomer is reacted with bromine in the presence of an anhydrous \(AlBr_3\) catalyst, only one monobrominated organic product is formed as the major product. Which compound is this starting isomer?

An unknown hydrocarbon with the molecular formula \(C_9H_{12}\) is heated under reflux with excess alkaline potassium manganate(VII). After acidification of the reaction mixture, a single organic product, benzoic acid (\(C_6H_5COOH\)), is obtained as a white crystalline solid. Which structure could represent the starting hydrocarbon?

The standard electrode potentials for three redox systems are shown below: 1. \(Ag^+(aq) + e^- \rightleftharpoons Ag(s) \quad E^\theta = +0.80\text{ V}\) 2. \(Cu^{2+}(aq) + 2e^- \rightleftharpoons Cu(s) \quad E^\theta = +0.34\text{ V}\) 3. \(Fe^{3+}(aq) + e^- \rightleftharpoons Fe^{2+}(aq) \quad E^\theta = +0.77\text{ V}\) Which reaction is thermodynamically feasible under standard conditions?

The standard electrode potential, \(E^\theta\), for the half-cell \(Cu^{2+}(aq) + 2e^- \rightleftharpoons Cu(s)\) is +0.34 V at 298 K. The Nernst equation for this half-cell is: \(E = E^\theta + \frac{0.059}{z}\log[Cu^{2+}(aq)]\). What is the electrode potential of this half-cell when the concentration of \(Cu^{2+}(aq)\) is \(1.0 \times 10^{-3}\text{ mol dm}^{-3}\) at 298 K?

The reaction between peroxodisulfate ions, \(S_2O_8^{2-}\), and iodide ions, \(I^-\), is represented by the equation: \(S_2O_8^{2-}(aq) + 2I^-(aq) \rightarrow 2SO_4^{2-}(aq) + I_2(aq)\). Initial rate data for this reaction at a constant temperature are shown below. Exp 1: \([S_2O_8^{2-}] = 0.10\text{ mol dm}^{-3}\), \([I^-] = 0.10\text{ mol dm}^{-3}\), rate = \(2.2 \times 10^{-4}\text{ mol dm}^{-3}\text{ s}^{-1}\). Exp 2: \([S_2O_8^{2-}] = 0.20\text{ mol dm}^{-3}\), \([I^-] = 0.10\text{ mol dm}^{-3}\), rate = \(4.4 \times 10^{-4}\text{ mol dm}^{-3}\text{ s}^{-1}\). Exp 3: \([S_2O_8^{2-}] = 0.20\text{ mol dm}^{-3}\), \([I^-] = 0.20\text{ mol dm}^{-3}\), rate = \(8.8 \times 10^{-4}\text{ mol dm}^{-3}\text{ s}^{-1}\). What are the orders of reaction with respect to \(S_2O_8^{2-}\) and \(I^-\)?

A chemical reaction follows the rate equation: \(\text{rate} = k[X][Y]^2\). In an experiment, when the concentration of \(X\) is \(0.40\text{ mol dm}^{-3}\) and the concentration of \(Y\) is \(0.50\text{ mol dm}^{-3}\), the rate of reaction is \(1.5 \times 10^{-3}\text{ mol dm}^{-3}\text{ s}^{-1}\). What is the value and units of the rate constant \(k\)?

A sample of 5.62 g of a hydrated transition metal salt, MSO4.7H2O, is heated to constant mass, driving off all water of crystallisation. 3.10 g of anhydrous MSO4 is obtained. (a)(i) Calculate the number of moles of water lost during heating. (ii) Calculate the molar mass of the anhydrous salt, MSO4, and hence identify the transition metal M. (iii) Write the full chemical formula of the hydrated transition metal salt. (b) A solution of this salt, MSO4(aq), reacts with aqueous sodium carbonate, Na2CO3(aq), to form a precipitate of MCO3. (i) Write an ionic equation, with state symbols, for this precipitation reaction. (ii) Calculate the mass of dry precipitate obtained when 50.0 cm3 of 0.250 mol dm-3 MSO4(aq) is reacted with excess sodium carbonate solution. (c)(i) In another reaction, MCO3(s) is heated to decompose into MO(s) and CO2(g). Calculate the volume of CO2(g), in dm3, produced at r.t.p. when 2.38 g of MCO3 is completely decomposed. [Take the volume of 1 mole of gas at r.t.p. to be 24.0 dm3] (ii) Explain how the thermal stability of Group 2 carbonates compares with that of transition metal carbonates like MCO3.

Benzene can be converted into methylbenzene, which can then undergo further reactions to produce a variety of organic compounds. (a) Write an equation for the reaction of benzene with a suitable reagent to produce methylbenzene, stating the catalyst used. (b)(i) Describe the mechanism for the electrophilic substitution reaction of benzene with chloromethane in the presence of an aluminium chloride catalyst. Your answer should show the generation of the electrophile, the structure of the intermediate, and the recovery of the catalyst. (ii) Explain why methylbenzene reacts faster than benzene in electrophilic substitution reactions. (c) Methylbenzene can be converted into other organic compounds. (i) State the reagents and conditions required to oxidise methylbenzene to benzoic acid. (ii) Complete the structure of the organic product formed when methylbenzene is reacted with a mixture of concentrated nitric acid and concentrated sulfuric acid at 30 degrees Celsius, and state the type of reaction. (iii) Draw the skeletal structure of the product formed when methylbenzene reacts with chlorine in the presence of ultraviolet light.

Standard electrode potentials, E-standard, are useful for predicting redox processes. (a) Define the term standard electrode potential. (b) An electrochemical cell is constructed from the following half-cells: Half-cell A: A copper electrode in Cu2+(aq). [Cu2+(aq) + 2e- <=> Cu(s), E-standard = +0.34 V] Half-cell B: A platinum electrode in Fe3+(aq) and Fe2+(aq). [Fe3+(aq) + e- <=> Fe2+(aq), E-standard = +0.77 V] (i) State the material used as the inert electrode in Half-cell B, and describe the standard conditions required for this half-cell. (ii) Write the overall cell reaction that occurs when this cell operates. (iii) Calculate the standard cell potential, E-standard of the cell. (iv) Draw a fully labelled diagram of this electrochemical cell, showing the direction of electron flow in the external circuit. (c) The Nernst equation can be used to calculate electrode potentials under non-standard conditions: E = E-standard + (0.059/z) * log([oxidised species]/[reduced species]). Calculate the electrode potential, E, for Half-cell A if the concentration of Cu2+(aq) is decreased to 0.0150 mol dm-3.

The reaction between peroxodisulfate ions, S2O8 2-(aq), and iodide ions, I-(aq), is represented by: S2O8 2-(aq) + 2I-(aq) -> 2SO4 2-(aq) + I2(aq). Experimental data at constant temperature: Experiment 1: [S2O8 2-] = 0.0500 mol dm-3, [I-] = 0.0800 mol dm-3, Initial Rate = 1.10 x 10^-4 mol dm-3 s-1. Experiment 2: [S2O8 2-] = 0.1000 mol dm-3, [I-] = 0.0800 mol dm-3, Initial Rate = 2.20 x 10^-4 mol dm-3 s-1. Experiment 3: [S2O8 2-] = 0.0500 mol dm-3, [I-] = 0.1600 mol dm-3, Initial Rate = 2.20 x 10^-4 mol dm-3 s-1. (a)(i) Deduce the order of reaction with respect to S2O8 2-. Explain your reasoning. (ii) Deduce the order of reaction with respect to I-. Explain your reasoning. (iii) Write the rate equation for this reaction. (iv) Calculate the rate constant, k, for this reaction, using the data from Experiment 1. State its units. (b)(i) Explain why this reaction is very slow in the absence of a catalyst. (ii) Write two equations to show how Fe2+(aq) acts as a homogeneous catalyst for this reaction. (iii) Explain why the catalyst in this reaction is described as homogeneous. (c)(i) Sketch a Boltzmann distribution curve on appropriate axes to show how an increase in temperature increases the rate of reaction. Label your curves T1 and T2, where T2 > T1, and label the activation energy, Ea.

A sample of 2.50 g of a mixture of calcium carbonate, \( \text{CaCO}_3 \), and inert silica, \( \text{SiO}_2 \), is treated with 50.0 cm\(^3\) of 1.00 mol dm\(^{-3}\) hydrochloric acid (an excess). The resulting solution is made up to 250 cm\(^3\) in a volumetric flask. A 25.0 cm\(^3\) portion of this diluted solution requires 18.40 cm\(^3\) of 0.100 mol dm\(^{-3}\) sodium hydroxide for complete neutralisation. (a) Write a balanced chemical equation for the reaction of calcium carbonate with hydrochloric acid. (b) Calculate the number of moles of sodium hydroxide used in the titration. (c) Calculate the number of moles of excess hydrochloric acid in the 250 cm\(^3\) volumetric flask. (d) Determine the percentage by mass of \( \text{CaCO}_3 \) in the original mixture. Show all your working.

Methylbenzene is an important starting material in organic synthesis. (a) Write an equation for the generation of the nitronium ion electrophile, \( \text{NO}_2^+ \), from a mixture of concentrated nitric acid and concentrated sulfuric acid. (b) Explain why methylbenzene is nitrated much more rapidly than benzene. (c) Methylbenzene can be converted into 3-nitrobenzoic acid. State the reagents and conditions for each step of this two-step synthesis, and explain why the order of these steps is crucial.

An electrochemical cell is set up at 298 K 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. Given standard electrode potentials: \( \text{Fe}^{3+}(\text{aq}) + \text{e}^- \rightleftharpoons \text{Fe}^{2+}(\text{aq}) \quad E^\ominus = +0.77\text{ V} \) and \( \text{Ag}^+(\text{aq}) + \text{e}^- \rightleftharpoons \text{Ag}(\text{s}) \quad E^\ominus = +0.80\text{ V} \). (a) Calculate the standard cell potential, \( E^\ominus_{\text{cell}} \). (b) Write the overall ionic equation for the cell reaction that occurs when current flows under standard conditions. (c) The concentration of \( \text{Ag}^+ \) in the silver half-cell is decreased to \( 0.025\text{ mol dm}^{-3} \), while the \( \text{Fe}^{3+}/\text{Fe}^{2+} \) half-cell remains under standard conditions. Use the Nernst equation, \( E = E^\ominus + 0.059 \log [\text{Ag}^+] \), to calculate the non-standard electrode potential of the \( \text{Ag}^+/\text{Ag} \) half-cell. (d) Determine the cell potential, \( E_{\text{cell}} \), under these non-standard conditions and state whether the direction of electron flow in the external circuit is the same or has reversed.

The reaction between peroxodisulfate ions, \( \text{S}_2\text{O}_8^{2-} \), and iodide ions, \( \text{I}^- \), is represented by the equation: \( \text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow 2\text{SO}_4^{2-}(\text{aq}) + \text{I}_2(\text{aq}) \). Rate data at a constant temperature: Exp 1: \( [\text{S}_2\text{O}_8^{2-}] = 0.010 \), \( [\text{I}^-] = 0.020 \), Rate = \( 1.50 \times 10^{-4} \); Exp 2: \( [\text{S}_2\text{O}_8^{2-}] = 0.020 \), \( [\text{I}^-] = 0.020 \), Rate = \( 3.00 \times 10^{-4} \); Exp 3: \( [\text{S}_2\text{O}_8^{2-}] = 0.010 \), \( [\text{I}^-] = 0.040 \), Rate = \( 3.00 \times 10^{-4} \). All concentrations in mol dm\(^{-3}\), rates in mol dm\(^{-3}\) s\(^{-1}\). (a) Determine the order of reaction with respect to \( \text{S}_2\text{O}_8^{2-} \) and \( \text{I}^- \). Explain your reasoning. (b) Write the rate equation. (c) Calculate the rate constant, \( k \), and state its units. (d) Describe how \( \text{Fe}^{2+} \) ions act as a homogeneous catalyst for this reaction, including two chemical equations.

Hydrocarbons can be characterised by their molecular formulas. (a) A sample of 0.220 g of a liquid alkane, X, was vaporised at 100 degrees Celsius and \( 1.01 \times 10^5 \text{ Pa} \). The volume of vaporised hydrocarbon X was 78.5 cm\(^3\). (i) Calculate the relative molecular mass, \( M_{\text{r}} \), of X using the ideal gas equation. (ii) Determine the molecular formula of X. (b) A 10.0 cm\(^3\) sample of a gaseous hydrocarbon, Y, of molecular formula \( \text{C}_x\text{H}_y \), was mixed with 80.0 cm\(^3\) of \( \text{O}_2 \) (an excess) and exploded. After cooling to room temperature, the total gas volume was 65.0 cm\(^3\). Passing this mixture through aqueous NaOH reduced the volume to 35.0 cm\(^3\). Determine the molecular formula of Y. Show all working.

Phenol, \( \text{C}_6\text{H}_5\text{OH} \), is a weak acid. (a) Arrange the following compounds in order of increasing acidity: ethanol, phenol, water. Explain this trend by comparing the stability of their conjugate bases. (b) State what is observed when phenol reacts with: (i) bromine water at room temperature, (ii) dilute nitric acid at room temperature. (c) Suggest a two-step synthesis to convert phenylamine, \( \text{C}_6\text{H}_5\text{NH}_2 \), into phenol. Include all reagents and conditions.

The thermal decomposition of calcium carbonate is represented by: \( \text{CaCO}_3(\text{s}) \rightarrow \text{CaO}(\text{s}) + \text{CO}_2(\text{g}) \). Standard thermodynamic data at 298 K: \( \text{CaCO}_3(\text{s}) \): \( \Delta H^\ominus_{\text{f}} = -1207\text{ kJ mol}^{-1} \), \( S^\ominus = 93\text{ J K}^{-1}\text{ mol}^{-1} \); \( \text{CaO}(\text{s}) \): \( \Delta H^\ominus_{\text{f}} = -635\text{ kJ mol}^{-1} \), \( S^\ominus = 40\text{ J K}^{-1}\text{ mol}^{-1} \); \( \text{CO}_2(\text{g}) \): \( \Delta H^\ominus_{\text{f}} = -394\text{ kJ mol}^{-1} \), \( S^\ominus = 214\text{ J K}^{-1}\text{ mol}^{-1} \). (a) Explain why the standard entropy change of this reaction is positive. (b) Calculate the standard enthalpy change, \( \Delta H^\ominus_{\text{reaction}} \), in \( \text{kJ mol}^{-1} \). (c) Calculate the standard entropy change, \( \Delta S^\ominus_{\text{reaction}} \), in \( \text{J K}^{-1}\text{ mol}^{-1} \). (d) Determine the minimum temperature, in Kelvin, at which the decomposition of calcium carbonate becomes thermodynamically feasible. State any assumption made in your calculation.

Transition metal complexes display a wide range of structures and isomerism. (a) Define the term bidentate ligand, and name a neutral bidentate ligand. (b) The octahedral complex \( [\text{Co}(\text{en})_2\text{Cl}_2]^+ \) (where en is 1,2-diaminoethane) has cis and trans stereoisomers. (i) Describe the structural difference between the cis and trans isomers. (ii) State which of these two isomers is optically active, and explain why. (c) Describe how the d-orbitals of a transition metal ion split in an octahedral complex, and explain how this splitting leads to the complex being coloured.

**Part (a)**

Define the term *standard electrode potential*, \(E^\ominus\). [2]

**Part (b)**

An electrochemical cell is set up consisting of a standard \(\text{Cu}^{2+}(aq)/\text{Cu}(s)\) half-cell and a standard \(\text{Fe}^{3+}(aq)/\text{Fe}^{2+}(aq)\) half-cell.

(i) Write the two reduction half-equations that represent these standard half-cells and state their standard electrode potentials, \(E^\ominus\). [2]

(ii) Deduce the equation for the overall cell reaction when the cell is discharging and calculate the standard cell potential, \(E^\ominus_{\text{cell}}\). [2]

**Part (c)**

The concentration of the species in the iron half-cell are adjusted such that \([\text{Fe}^{3+}(aq)] = 0.85\text{ mol dm}^{-3}\) and \([\text{Fe}^{2+}(aq)] = 0.015\text{ mol dm}^{-3}\).

Calculate the electrode potential, \(E\), of this non-standard iron half-cell at \(298\text{ K}\) using the Nernst equation:

\[E = E^\ominus + \frac{0.059}{z}\log\left(\frac{[\text{oxidised species}]}{[\text{reduced species}]}\right)\]

Show your working. [3]

**Part (d)**

The copper half-cell concentration of \(\text{Cu}^{2+}(aq)\) is now decreased to \(0.050\text{ mol dm}^{-3}\).

Explain, with reference to the \(\text{Cu}^{2+}/\text{Cu}\) half-cell electrode potential and the overall cell system, how this change affects the value of the cell potential, \(E_{\text{cell}}\). [2.1]

A student investigates the kinetics of the reaction between peroxodisulfate ions, \( \text{S}_2\text{O}_8^{2-}(\text{aq}) \), and iodide ions, \( \text{I}^-(\text{aq}) \): \( \text{S}_2\text{O}_8^{2-}(\text{aq}) + 2\text{I}^-(\text{aq}) \rightarrow 2\text{SO}_4^{2-}(\text{aq}) + \text{I}_2(\text{aq}) \). This reaction is studied using a clock reaction method. A small, constant amount of sodium thiosulfate, \( \text{Na}_2\text{S}_2\text{O}_3 \), and starch indicator are added to the reaction mixture. The time, \( t \), taken for the mixture to turn blue-black is recorded. Five experiments are carried out at 298 K. In each experiment, the volume of \( 0.100 \text{ mol dm}^{-3} \) \( \text{K}_2\text{S}_2\text{O}_8 \) is kept constant at \( 20.0 \text{ cm}^3 \), and the volume of \( 0.010 \text{ mol dm}^{-3} \) \( \text{Na}_2\text{S}_2\text{O}_3 \) is kept constant at \( 5.0 \text{ cm}^3 \). The volume of \( 0.200 \text{ mol dm}^{-3} \) \( \text{KI} \), \( x \), is varied, and distilled water is added to keep the total volume of \( \text{KI} \) and water constant at \( 25.0 \text{ cm}^3 \). Starch indicator (\( 1.0 \text{ cm}^3 \)) is added to each mixture, making the total volume of the mixture \( 51.0 \text{ cm}^3 \) in all cases. (a) (i) Explain why the total volume of the reaction mixture must be kept constant at \( 51.0 \text{ cm}^3 \). (ii) Calculate the concentration of \( \text{I}^-(\text{aq}) \) in the reaction mixture for Experiment 2, where \( x = 10.0 \text{ cm}^3 \). (b) The times, \( t \), recorded for the five experiments are shown below: Experiment 1: \( x = 5.0 \text{ cm}^3 \), \( t = 120.0 \text{ s} \); Experiment 2: \( x = 10.0 \text{ cm}^3 \), \( t = 60.0 \text{ s} \); Experiment 3: \( x = 15.0 \text{ cm}^3 \), \( t = 50.0 \text{ s} \); Experiment 4: \( x = 20.0 \text{ cm}^3 \), \( t = 30.0 \text{ s} \); Experiment 5: \( x = 25.0 \text{ cm}^3 \), \( t = 24.0 \text{ s} \). (i) Calculate the value of \( 1/t \) for each of the five experiments. Identify which experiment gave an anomalous result and suggest a reason for this anomaly. (ii) State and explain how the student should treat this anomalous result when analyzing the data graphically. (iii) Use the relationship between \( 1/t \) and \( x \) to deduce the order of reaction with respect to \( \text{I}^-(\text{aq}) \). (iv) Calculate the percentage uncertainty in the time measurement for Experiment 5, given that the uncertainty of the stopwatch is \( \pm 0.2 \text{ s} \). (c) Describe how the student could modify this experimental procedure to determine the activation energy, \( E_a \), of the reaction. State the measurements that would be recorded and how the data would be processed graphically.

An electrochemical cell is set up to investigate the Nernst equation for the \( \text{Cu}^{2+}(\text{aq})/\text{Cu}(\text{s}) \) half-cell. The cell is represented as: \( \text{Zn}(\text{s}) \mid \text{Zn}^{2+}(\text{aq}, 1.00 \text{ mol dm}^{-3}) \parallel \text{Cu}^{2+}(\text{aq}, C \text{ mol dm}^{-3}) \mid \text{Cu}(\text{s}) \). The standard cell potential, \( E^{\ominus}_{\text{cell}} \), is \( 1.10 \text{ V} \) at 298 K. The cell potential, \( E_{\text{cell}} \), is measured for various concentrations, \( C \), of \( \text{Cu}^{2+}(\text{aq}) \) at 298 K. (a) (i) Draw a fully labelled diagram of the experimental setup used to measure \( E_{\text{cell}} \). (ii) State the role of the salt bridge and explain why potassium nitrate, \( \text{KNO}_3 \), is a suitable electrolyte for it. (b) Describe how a student would prepare \( 100 \text{ cm}^3 \) of \( 0.0100 \text{ mol dm}^{-3} \) \( \text{CuSO}_4(\text{aq}) \) from a stock solution of \( 1.00 \text{ mol dm}^{-3} \) \( \text{CuSO}_4(\text{aq}) \). State the names and capacities of the volumetric apparatus required. (c) The table below shows the results of the experiment: \( [\text{Cu}^{2+}] \) in \( \text{mol dm}^{-3} \): \( 1.00, 0.100, 0.0100, 0.00100, 0.000100 \). Measured \( E_{\text{cell}} \) in \( \text{V} \): \( 1.10, 1.07, 1.04, 1.01, 0.98 \). (i) Calculate the value of \( \log[\text{Cu}^{2+}] \) for each concentration and describe the trend in \( E_{\text{cell}} \) as \( \log[\text{Cu}^{2+}] \) decreases. (ii) Determine the gradient of the graph of \( E_{\text{cell}} \) against \( \log[\text{Cu}^{2+}] \). Show your working. (iii) The theoretical gradient of this graph is equal to \( \frac{2.303RT}{zF} \), where \( R = 8.31 \text{ J K}^{-1} \text{ mol}^{-1} \), \( T = 298 \text{ K} \), \( F = 96500 \text{ C mol}^{-1} \), and \( z \) is the number of electrons transferred per copper ion. Use your calculated gradient to determine the experimental value of \( z \) to two significant figures. (d) Predict how the measured cell potential, \( E_{\text{cell}} \), would change, if at all, under the following conditions. Explain your predictions. (i) The concentration of \( \text{Zn}^{2+}(\text{aq}) \) in the zinc half-cell is increased to \( 2.00 \text{ mol dm}^{-3} \), while \( [\text{Cu}^{2+}(\text{aq})] \) is kept at \( 1.00 \text{ mol dm}^{-3} \). (ii) The copper electrode in the copper half-cell is replaced with a copper sheet of twice the surface area.