1.	Methanol is made in large quantities as it is used in the production of polymers and in fuels.
The enthalpy of combustion of methanol can be determined theoretically or experimentally.
		CH3OH(l) + 1O2(g) → CO2(g) + 2H2O(g)
 
(a)	Using the information from Table 10 of the Data Booklet, determine the theoretical enthalpy of combustion of methanol.
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(3)

(b)	The enthalpy of combustion of methanol can also be determined experimentally in a school laboratory. A burner containing methanol was weighed and used to heat water in a test tube as illustrated below.
		
	The following data were collected.

Initial mass of burner and methanol / g	80.557	
Final mass of burner and methanol / g	80.034	
Mass of water in test tube / g	20.000	
Initial temperature of water / °C	21.5	
Final temperature of water / °C	26.4	
 
(i)	Calculate the amount, in mol, of methanol burned.
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(2)

(ii)	Calculate the heat absorbed, in kJ, by the water.
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(3)
 
(iii)	Determine the enthalpy change, in kJ mol–1, for the combustion of 1 mole of methanol.
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(2)
 
(c)	The Data Booklet value for the enthalpy of combustion of methanol is –726 kJ mol–1.
Suggest why this value differs from the values calculated in parts (a) and (b).
(i)	Part (a)
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(1)
(ii)	Part (b)
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(1)
(Total 12 marks)

2.	Ethanol is used as a component in fuel for some vehicles. One fuel mixture contains 10 % by mass of ethanol in unleaded petrol (gasoline). This mixture is often referred to as Gasohol E10.
(a)	Assume that the other 90 % by mass of Gasohol E10 is octane. 1.00 kg of this fuel mixture was burned.
	CH3CH2OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)	∆HO = –1367 kJ mol–1
C8H18(l) + 12O2(g) → 8CO2(g) + 9H2O(l)	∆HO = –5470 kJ mol–1
 
(i)	Calculate the mass, in g, of ethanol and octane in 1.00 kg of the fuel mixture.
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(1)
 
(ii)	Calculate the amount, in mol, of ethanol and octane in 1.00 kg of the fuel mixture.
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(1)
 
(iii)	Calculate the total amount of energy, in kJ, released when 1.00 kg of the fuel mixture is completely burned.
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(3)

(b)	If the fuel blend was vaporized before combustion, predict whether the amount of energy released would be greater, less or the same. Explain your answer.
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(2)
(Total 7 marks)

3.	If white anhydrous copper(II) sulfate powder is left in the atmosphere it slowly absorbs water vapour giving the blue pentahydrated solid.
	CuSO4(s) + 5H2O(l) → CuSO4•5H2O(s)
(anhydrous)	(pentahydrated)
	It is difficult to measure the enthalpy change for this reaction directly. However, it is possible to measure the heat changes directly when both anhydrous and pentahydrated copper(II) sulfate are separately dissolved in water, and then use an energy cycle to determine the required enthalpy change value, ∆Hx, indirectly.

(a)	To determine ∆H1 a student placed 50.0 g of water in a cup made of expanded polystyrene and used a data logger to measure the temperature. After two minutes she dissolved 3.99 g of anhydrous copper(II) sulfate in the water and continued to record the temperature while continuously stirring. She obtained the following results.

(i)	Calculate the amount, in mol, of anhydrous copper(II) sulfate dissolved in the 50.0 g of water.
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(1)
 
(ii)	Determine what the temperature rise would have been, in °C, if no heat had been lost to the surroundings.
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(2)

(iii)	Calculate the heat change, in kJ, when 3.99 g of anhydrous copper(II) sulfate is dissolved in the water.
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(2)
 
(iv)	Determine the value of ∆H1 in kJ mol–1.
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(1)
 
(b)	To determine ∆H2, 6.24 g of pentahydrated copper(II) sulfate was dissolved in 47.75 g of water. It was observed that the temperature of the solution decreased by 1.10 °C.
(i)	Calculate the amount, in mol, of water in 6.24 g of pentahydrated copper(II) sulfate.
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(2)
 
(ii)	Determine the value of ∆H2 in kJ mol–1.
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(2)
 
(iii)	Using the values obtained for ∆H1 in (a) (iv) and ∆H2 in (b) (ii), determine the value for ∆Hx in kJ mol–1.
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(1)
 
(c)	The magnitude (the value without the + or – sign) found in a data book for ∆Hx is 78.0 kJ mol–1.
(i)	Calculate the percentage error obtained in this experiment. (If you did not obtain an answer for the experimental value of ∆Hx then use the value 70.0 kJ mol–1, but this is not the true value.)
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(1)
 
(ii)	The student recorded in her qualitative data that the anhydrous copper(II) sulfate she used was pale blue rather than completely white. Suggest a reason why it might have had this pale blue colour and deduce how this would have affected the value she obtained for ∆Hx.
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(2)
(Total 14 marks)

4.	Brass is a copper containing alloy with many uses. An analysis is carried out to determine the percentage of copper present in three identical samples of brass. The reactions involved in this analysis are shown below.
	Step 1: Cu(s) + 2HNO3(aq) + 2H+(aq) → Cu2+(aq) + 2NO2(g) + 2H2O(l)
	Step 2: 4I–(aq) + 2Cu2+(aq) → 2CuI(s) + I2(aq)
	Step 3: I2(aq) + 2S2O32–(aq) → 2I–(aq) + S4O62–(aq)

(a)	(i)	Deduce the change in the oxidation numbers of copper and nitrogen in step 1.
	Copper:
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	Nitrogen:
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(2)
 
(ii)	Identify the oxidizing agent in step 1.
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(1)
 
(b)	A student carried out this experiment three times, with three identical small brass nails, and obtained the following results.
		Mass of brass = 0.456 g ± 0.001 g

Titre	1	2	3	
Initial volume of 0.100 mol dm–3 S2O32– (± 0.05 cm3)	0.00	0.00	0.00	
Final volume of 0.100 mol dm–3 S2O32– (± 0.05 cm3)	28.50	28.60	28.40	
Volume added of 0.100 mol dm–3 S2O32– (± 0.10 cm3)	28.50	28.60	28.40	
Average volume added of 0.100 mol dm–3 S2O32–(± 0.10 cm3)	28.50	
 
(i)	Calculate the average amount, in mol, of S2O32– added.
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(2)

(ii)	Calculate the amount, in mol, of copper present in the brass.
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(1)
 
(iii)	Calculate the mass of copper in the brass.
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(1)
 
(iv)	Calculate the percentage by mass of copper in the brass.
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(1)
 
(v)	The manufacturers claim that the sample of brass contains 44.2 % copper by mass.
Determine the percentage error in the result.
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(1)
 
(c)	With reference to its metallic structure, describe how brass conducts electricity.
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(1)
(Total 10 marks)

5.	The percentage by mass of calcium carbonate in eggshell was determined by adding excess hydrochloric acid to ensure that all the calcium carbonate had reacted. The excess acid left was then titrated with aqueous sodium hydroxide.
(a)	A student added 27.20 cm3 of 0.200 mol dm–3 HCl to 0.188 g of eggshell. Calculate the amount, in mol, of HCl added.
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(1)
 
(b)	The excess acid requires 23.80 cm3 of 0.100 mol dm–3 NaOH for neutralization.
Calculate the amount, in mol, of acid that is in excess.
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(1)
 
(c)	Determine the amount, in mol, of HCl that reacted with the calcium carbonate in the eggshell.
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(1)
 
(d)	State the equation for the reaction of HCl with the calcium carbonate in the eggshell.
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(2)
 
(e)	Determine the amount, in mol, of calcium carbonate in the sample of the eggshell.
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(2)
 
(f)	Calculate the mass and the percentage by mass of calcium carbonate in the eggshell sample.
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(3)
 
(g)	Deduce one assumption made in arriving at the percentage of calcium carbonate in the eggshell sample.
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(1)
(Total 11 marks)

6.	Biodiesel makes use of plants’ ability to fix atmospheric carbon by photosynthesis.
Many companies and individuals are now using biodiesel as a fuel in order to reduce their carbon footprint. Biodiesel can be synthesized from vegetable oil according to the following reaction.

(a)	Identify the organic functional group present in both vegetable oil and biodiesel.
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(1)
 
(b)	For part of her extended essay investigation into the efficiency of the process, a student reacted a pure sample of a vegetable oil (where R = C17H33) with methanol.
The raw data recorded for the reaction is below.
	Mass of oil	=	1013.0 g
Mass of methanol	=	200.0 g
Mass of sodium hydroxide	=	3.5 g
Mass of biodiesel produced	=	811.0 g
	The relative molecular mass of the oil used by the student is 885.6. Calculate the amount (in moles) of the oil and the methanol used, and hence the amount (in moles) of excess methanol.
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(3)
 
(c)	The reversible arrows in the equation indicate that the production of biodiesel is an equilibrium process.
(i)	State what is meant by the term dynamic equilibrium.
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(1)
 
(ii)	Using the abbreviations [vegetable oil], [methanol], [glycerol] and [biodiesel] deduce the equilibrium constant expression (Kc) for this reaction.

(1)
 
(iii)	Suggest a reason why excess methanol is used in this process.
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(1)
 
(iv)	State and explain the effect that the addition of the sodium hydroxide catalyst will have on the position of equilibrium.
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(2)
 
(d)	The reactants had to be stirred vigorously because they formed two distinct layers in the reaction vessel. Explain why they form two distinct layers and why stirring increases the rate of reaction.
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(2)
 
(e)	Calculate the percentage yield of biodiesel obtained in this process.
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(2)
 
(f)	When biodiesel is combusted it produces carbon dioxide. Explain why the use of biodiesel as a fuel does not significantly contribute to global warming.
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(1)
(Total 14 marks)

7.	Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of ethene to form ethane.
		C2H4(g) + H2(g) → C2H6(g)
	John used the average bond enthalpies from Table 10. Marit used the values of enthalpies of combustion from Table 12.
 
(a)	Calculate the value for the enthalpy of hydrogenation of ethene obtained using the average bond enthalpies given in Table 10.
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(2)
 
(b)	Marit arranged the values she found in Table 12 into an energy cycle.

	Calculate the value for the enthalpy of hydrogenation of ethene from the energy cycle.
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(1)

(c)	Suggest one reason why John’s answer is slightly less accurate than Marit’s answer.
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(1)
 
(d)	John then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane.
		C6H10(l) + H2(g) → C6H12(l)
 
(i)	Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene.
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(1)
 
(ii)	The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for ethene. John’s hypothesis was that it would be the same. Determine why the use of average bond enthalpies is less accurate for the cyclohexene equation shown above, than it was for ethene. Deduce what extra information is needed to provide a more accurate answer.
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(2)
(Total 7 marks)

8.	Two students were asked to use information from the Data Booklet to calculate a value for the enthalpy of hydrogenation of ethene to form ethane.
		C2H4(g) + H2(g) → C2H6(g)
	John used the average bond enthalpies from Table 10. Marit used the values of enthalpies of combustion from Table 12.
 
(a)	Calculate the value for the enthalpy of hydrogenation of ethene obtained using the average bond enthalpies given in Table 10.
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(2)
 
(b)	Determine the value for the enthalpy of hydrogenation of ethene using the values for the enthalpies of combustion of ethene, hydrogen and ethane given in Table 12.
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(2)
 
(c)	Suggest one reason why John’s answer is slightly less accurate than Marit’s answer and calculate the percentage difference.
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(2)

(d)	John then decided to determine the enthalpy of hydrogenation of cyclohexene to produce cyclohexane.
		C6H10(l) + H2(g) → C6H12(l)
 
(i)	Use the average bond enthalpies to deduce a value for the enthalpy of hydrogenation of cyclohexene.
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(1)
 
(ii)	The percentage difference between these two methods (average bond enthalpies and enthalpies of combustion) is greater for cyclohexene than it was for ethene. John’s hypothesis was that it would be the same. Determine why the use of average bond enthalpies is less accurate for the cyclohexene equation shown above, than it was for ethene. Deduce what extra information is needed to provide a more accurate answer.
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(2)
(Total 9 marks)

9.	The data below is from an experiment used to determine the percentage of iron present in a sample of iron ore. This sample was dissolved in acid and all of the iron was converted to Fe2+. The resulting solution was titrated with a standard solution of potassium manganate(VII), KMnO4. This procedure was carried out three times. In acidic solution, MnO4– reacts with Fe2+ ions to form Mn2+ and Fe3+ and the end point is indicated by a slight pink colour.

Titre	1	2	3	
Initial burette reading / cm3	1.00	23.60	10.00	
Final burette reading / cm3	24.60	46.10	32.50	

Mass of iron ore / g	3.682 × 10–1	
Concentration of KMnO4 solution / mol dm–3	2.152 × 10–2	
 
(a)	Deduce the balanced redox equation for this reaction in acidic solution.
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 (2)
 
(b)	Identify the reducing agent in the reaction.
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(1)
 
(c)	Calculate the amount, in moles, of MnO4– used in the titration.
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(2)
 
(d)	Calculate the amount, in moles, of Fe present in the 3.682 × 10–1 g sample of iron ore.
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(2)
 
(e)	Determine the percentage by mass of Fe present in the 3.682 × 10–1 g sample of iron ore.
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(2)
(Total 9 marks)

10.	The percentage of iron(II) ions, Fe2+, in a vitamin tablet can be estimated by dissolving the tablet in dilute sulfuric acid and titrating with standard potassium manganate(VII) solution, KMnO4(aq). During the process iron(II) is oxidized to iron(III) and the manganate(VII) ion is reduced to the manganese(II) ion, Mn2+(aq). It was found that one tablet with a mass of 1.43 g required 11.6 cm3 of 2.00 × 10–2 mol dm–3 KMnO4(aq) to reach the end-point.
(a)	(i)	State the half-equation for the oxidation of the iron(II) ions.
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(1)
 
(ii)	State the half-equation for the reduction of the MnO4– ions in acidic solution.
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(1)
 
(iii)	Deduce the overall redox equation for the reaction.
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(1)
 
(b)	(i)	Calculate the amount, in moles, of MnO4– ions present in 11.6 cm3 of
2.00 × 10–2 mol dm–3 KMnO4(aq).
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(1)
 
(ii)	Calculate the amount, in moles, of Fe2+ ions present in the vitamin tablet.
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(1)

(iii)	Determine the percentage by mass of Fe2+ ions present in the vitamin tablet.
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(2)
(Total 7 marks)

11.	The data below is from an experiment used to measure the enthalpy change for the combustion of 1 mole of sucrose (common table sugar), C12H22O11(s). The time-temperature data was taken from a data-logging software programme.

	Mass of sample of sucrose, m = 0.4385 g
	Heat capacity of the system, Csystem = 10.114 kJ K–1

(a)	Calculate ΔT, for the water, surrounding the chamber in the calorimeter.
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(1)
 
(b)	Determine the amount, in moles, of sucrose.
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(1)
 
(c)	(i)	Calculate the enthalpy change for the combustion of 1 mole of sucrose.
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(1)
 
(ii)	Using Table 12 of the Data Booklet, calculate the percentage experimental error based on the data used in this experiment.
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(1)

(d)	A hypothesis is suggested that TNT, 2-methyl-1,3,5-trinitrobenzene, is a powerful explosive because it has:
•	a large enthalpy of combustion
•	a high reaction rate
•	a large volume of gas generated upon combustion
	Use your answer in part (c)(i) and the following data to evaluate this hypothesis:
	Equation for combustion	Relative
rate of
combustion	Enthalpy of
combustion
/ kJ mol–1	
Sucrose	C12H22O11(s) + 12O2(g) ® 12CO2(g) + 11H2O(g)	Low		
TNT	2C7H5N3O6(s) ® 7CO(g) + 7C(s) + 5H2O(g) + 3N2(g)	High	3406	
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(3)
(Total 7 marks)

12.	Alex and Hannah were asked to investigate the kinetics involved in the iodination of propanone.
They were given the following equation by their teacher.
		CH3COCH3(aq) + I2(aq)  CH2ICOCH3(aq) + HI(aq)
	Alex’s hypothesis was that the rate will be affected by changing the concentrations of the propanone and the iodine, as the reaction can happen without a catalyst. Hannah’s hypothesis was that as the catalyst is involved in the reaction, the concentrations of the propanone, iodine and the hydrogen ions will all affect the rate.
	They carried out several experiments varying the concentration of one of the reactants or the catalyst while keeping other concentrations and conditions the same, and obtained the results below.

	Composition by volume of mixture / cm3	Initial rate	
Experiment	1.00 mol dm–3
CH3COCH3(aq)	Water	1.00 mol dm–3
H+(aq)	5.00 × 10–3
mol dm–3 I2 in KI	/ mol dm–3 s–1	
1	10.0	60.0	10.0	20.0	4.96 × 10–6	
2	10.0	50.0	10.0	30.0	5.04 × 10–6	
3	5.0	65.0	10.0	20.0	2.47 × 10–6	
4	10.0	65.0	5.0	20.0	2.51 × 10–6	
 
(a)	Explain why they added water to the mixtures.
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(1)
 
(b)	(i)	Deduce the order of reaction for each substance and the rate expression from the results.
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(2)

(ii)	Comment on whether Alex’s or Hannah’s hypothesis is correct.
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(1)
 
(c)	Using the data from Experiment 1, determine the concentration of the substances used and the rate constant for the reaction including its units.
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(3)
 
(d)	(i)	This reaction uses a catalyst. Sketch and annotate the Maxwell-Boltzmann energy distribution curve for a reaction with and without a catalyst on labelled axes below.

(3)
 
(ii)	Describe how a catalyst works.
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(1)
(Total 11 marks)

13.	Consider the following graph of ln k against  (temperature in Kelvin) for the second order decomposition of N2O into N2 and O.
		N2O → N2 + O
		
		 K–1
 
(a)	State how the rate constant, k varies with temperature, T.
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(1)

(b)	Determine the activation energy, Ea, for this reaction.
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(3)
 
(c)	The rate expression for this reaction is rate = k [N2O]2 and the rate constant is
0.244 dm3 mol–1 s–1 at 750 °C.
	A sample of N2O of concentration 0.200 mol dm–3 is allowed to decompose. Calculate the rate when 10 % of the N2O has reacted.
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(2)
(Total 6 marks)

14.	(a)	The production of ammonia is an important industrial process.
		N2(g) + 3H2(g)  2NH3(g)
 
(i)	Using the average bond enthalpy values in Table 10 of the Data Booklet, determine the standard enthalpy change for this reaction.
(3)

(ii)	The standard entropy values, S, at 298 K for N2(g), H2(g) and NH3(g) are 193, 131 and 192 JK–1 mol–1 respectively. Calculate ∆SO for the reaction and with reference to the equation above, explain the sign of ∆SO.
(4)
 
(iii)	Calculate ∆GO for the reaction at 298 K.
(1)
 
(iv)	Describe and explain the effect of increasing temperature on the spontaneity of the reaction.
(2)
 
(b)	The reaction used in the production of ammonia is an equilibrium reaction. Outline the characteristics of a system at equilibrium.
(2)
 
(c)	Deduce the equilibrium constant expression, Kc, for the production of ammonia.
(1)
 
(d)	(i)	0.20 mol of N2(g) and 0.20 mol of H2(g) were allowed to reach equilibrium in a 1 dm3 closed container. At equilibrium the concentration of NH3(g) was 0.060 mol dm–3. Determine the equilibrium concentrations of N2(g) and H2(g) and calculate the value of Kc.
(3)
 
(ii)	Predict and explain how increasing the temperature will affect the value of Kc.
(2)
 
(e)	Describe how increasing the pressure affects the yield of ammonia.
(2)
 
(f)	In practice, typical conditions used in the Haber process are a temperature of 500 °C and a pressure of 200 atmospheres. Outline why these conditions are used rather than those that give the highest yield.
(2)

(g)	A catalyst of iron is used in the Haber process. State and explain how the catalyst affects Kc and the position of equilibrium.
(3)
(Total 25 marks)

15.	Consider the following reaction.
		2CH3OH(g) + H2(g) → C2H6(g) + 2H2O(g)
 
(a)	The standard enthalpy change of formation for CH3OH(g) at 298 K is –201 kJ mol–1 and for H2O(g) is –242 kJ mol–1. Using information from Table 11 of the Data Booklet, determine the enthalpy change for this reaction.
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(2)
 
(b)	The standard entropy for CH3OH(g) at 298 K is 238 J K–1 mol–1, for H2(g) is
131 J K–1 mol–1 and for H2O(g) is 189 J K–1 mol–1. Using information from Table 11 of the Data Booklet, determine the entropy change for this reaction.
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(2)

(c)	Calculate the standard change in free energy, at 298 K, for the reaction and deduce whether the reaction is spontaneous or non-spontaneous.
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(3)
(Total 7 marks)