1.	(a)	State Hess’s Law.
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(2)
 
(b)	Define the term standard enthalpy change of combustion.
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(3)
 
(c)	The equation for the combustion of ethanol in air is
	C2H5OH(l) + 3O2(g) ® 2CO2(g) + 3H2O(l)
and the structural representation of this is:

(i)	Calculate the enthalpy change for this reaction using the average bond enthalpy values given below.
Bond	Average bond
enthalpy/kJ mol–1	Bond	Average bond
enthalpy/kJ mol–1	
C¾H	+412	C¾C	+348	
C¾O	+360	O¾H	+463	
O¾O	+496	C¾O	+743	
(3)

(ii)	Draw and label an enthalpy level diagram to represent this reaction.

(2)
(Total 10 marks)

2.	In an experiment to find the enthalpy change when copper is displaced from a solution  of copper ions excess zinc was added to 50.0 cm3 of 1.00 mol dm–3 aqueous copper(ii) sulphate in a plastic cup.
	Zn(s) + Cu2+(aq) ® Zn2+(aq) + Cu(s)
	The temperature of the solution in the cup was measured every minute for 10 minutes with the zinc being added after 3.5 minutes.

The temperature readings are shown on the graph below.

(a)	Suggest two reasons why a series of temperature readings is taken rather than simply initial and final readings.
First reason..............................................................................................................
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Second reason..........................................................................................................
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(2)

(b)	Use the graph to calculate the maximum temperature change, DT, as the reaction takes place.

DT = ...........................................°C
(2)
 
(c)	Calculate the enthalpy change for the reaction using the formula below, giving your answer to an appropriate number of significant figures.
	DH = –4.18 × DT kJ mol–1

(2)
(Total 6 marks)

3.	(a)	(i)	Define the term standard enthalpy of formation, DHf.
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(3)

(ii)	The following table shows some values of standard enthalpy of formation.
Name	Formula	DHf/kJ mol–1	
ethene	C2H4(g)	+52.3	
hydrogen bromide	HBr(g)	–36.2	
bromoethane	C2H5Br(g)	–60.4	
 
	Use the data in the table above to calculate the standard enthalpy change for the following reaction.
C2H4(g) + HBr(g) ® C2H5Br(g)

(2)
 
(iii)	State the significance of the sign of the value obtained in part (a)(ii) above.
..........................................................................................................................
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(1)

(b)	Enthalpy changes can also be calculated using average bond enthalpy data. 
Bond	Average bond enthalpy/kJ mol–1	
C == C	+612	
C ¾ C	+348	
C ¾ H	+412	
C ¾ Br	+276	
H¾ Br	+366	
 
	Use the data in the table above to recalculate the enthalpy change for the reaction in part (a)(ii).
C2H4(g) + HBr(g) ® C2H5Br(g)

(3)
 
(c)	Suggest why the value obtained in part (b) above is likely to be less accurate than that obtained in part (a)(ii).
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(2)
(Total 11 marks)

4.	The reaction of an acid with a base to give a salt is an exothermic reaction. In an experiment to determine the enthalpy of neutralisation of hydrochloric acid with sodium hydroxide, 50.0cm3 of 1.00 mol dm–3 HCl was mixed with 50.0 cm3 of 1.10 mol dm–3 NaOH. The temperature rise obtained was 6.90 °C.
(a)	Define the term enthalpy of neutralisation.
....................................................................................................................................
....................................................................................................................................
(1)
 
(b)	Assuming that the density of the final solution is 1.00 g cm–3 and that its heat capacity is 4.18 J K–1 g–1, calculate the heat evolved during the reaction.

(3)
 
(c)	0.0500 mol of acid was neutralised in this reaction; calculate DHneutralisation in
kJ mol–1.

(2)

(d)	Suggest why sodium hydroxide is used in slight excess in the experiment.
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(1)
(Total 7 marks)

5.	A student was required to determine the enthalpy change for the reaction between iron and copper sulphate solution.
 
	The student produced the following account of their experiment.

(a)	Suggest two improvements you would make to this experiment. Give a reason for each of the improvements suggested.
Improvement 1 ...........................................................................................................
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Reason 1 ....................................................................................................................
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Improvement 2 ...........................................................................................................
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Reason 2 ....................................................................................................................
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(4)
 
(b)	In an improved version of the same experiment a maximum temperature rise of
15.2 °C occurred when reacting excess iron with 50.0 cm3 of 0.500 mol dm–3 aqueous copper sulphate solution.
(i)	Using this data and taking the specific heat capacity of all aqueous solutions as 4.18 Jg–1 deg–1 calculate the heat change.

(1)
 
(ii)	Calculate the number of moles of copper sulphate used.

(1)

(iii)	Calculate the enthalpy change of this reaction in kJ mol–1.

(2)
(Total 8 marks)

6.	(a)	Define the term standard enthalpy of combustion, making clear the meaning of standard in this context.
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(3)

(b)	Use the enthalpies of combustion given below to find the enthalpy change for the reaction:
2C(graphite) + 2H2(g) + O2(g) ® CH3COOH(l)
	DHcombustion/kJ mol–1	
C(graphite)	–394	
H2(g)	–286	
CH3COOH(l)	–874	

(3)
 
(c)	With reference to ethanoic acid, CH3COOH, what is the enthalpy change obtained in (b) called?
………….…………………………………………………………………………..
(1)

(d)	Draw an enthalpy level diagram to represent the enthalpy change for the combustion of graphite. Show both the enthalpy levels of the reactants and products and an energy profile which represents the activation energy for the reaction.

(3)
(Total 10 marks)

7.	(a)	This question is about finding the formula of copper hydroxide. The method is as follows:

	20.0 cm3 of an aqueous solution of a copper salt of concentration 1.00 mol dm–3 was placed in a polystyrene cup and its temperature measured using a thermometer graduated in 0.1 °C intervals.
	A burette was filled with aqueous sodium hydroxide, of concentration 2.00 mol dm–3.
	2.00 cm3 of sodium hydroxide solution was run into the solution of the copper salt and the temperature was measured immediately.
 
	As soon as possible a further 2.00 cm3 of sodium hydroxide solution was run in and the temperature measured again.
	This process of adding 2.00 cm3 portions of sodium hydroxide solution and measuring the temperature was continued until a total of 36.0 cm3 of the sodium hydroxide solution had been added.
 
	The temperature readings are shown in the graph below.

(i)	Explain why the temperature reaches a maximum and then falls slightly on addition of further sodium hydroxide solution.
………….……………….……………………………………………………..
………….……………….……………………………………………………..
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(2)
 
(ii)	From the graph, what volume of the aqueous sodium hydroxide was required for complete reaction?
………….……………….……………………………………………………..
(1)
 
(iii)	Calculate the amount (number of moles) of sodium hydroxide in this volume of solution.

(1)
 
(iv)	Calculate the amount (number of moles) of copper ions that have reacted.

(1)
 
(v)	Write the ratio of moles of copper ions to hydroxide ions reacting.

(1)
 
(vi)	Write the formula of the copper hydroxide that is produced.

(1)
 
(b)	The data call be used to find the enthalpy change for the reaction between sodium hydroxide and the copper salt.
(i)	Use the graph to find the temperature rise that occurs for complete reaction.
………….……………….……………………………………………………..
(1)
 
(ii)	Find the heat change, q, that occurs in the polystyrene cup for complete reaction.
Use the formula
q = 168 × DT joules

(1)
 
(iii)	Use your results from (a)(iv) and (b)(ii) above, to find the molar enthalpy change, DH, for the reaction. Give the correct sign and units to the answer.

(3)
 
(c)	Identify one potential source of error in this experiment, and say what you would do to reduce its effect.
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(2)
(Total 14 marks)

8.	An excess of zinc powder was added to 20.0 cm3 of a solution of copper(II) sulphate of concentration 0.500 mol dm–3. The temperature increased by 26.3 °C.
(a)	How many moles of copper(II) sulphate were used in this experiment?

(1)
 
(b)	Calculate the enthalpy change, ΔH, in kJ mol–1 for this reaction given that:
energy change =	specific
heat capacity	´ mass of
solution	´ temperature
change	
/J	/J g-1 K-1	/g	/K	
	Assume that the mass of solution is 20.0 g and the specific heat capacity of the solution is 4.18 J g–1K–1.

(2)
(Total 5 marks)

9.	A sample of titanium (atomic number 22) is made up of five isotopes. The sample has the following percentage composition:
Mass number
46
47
48
49
50	% composition
8.0
7.3
74.0
5.5
5.2	

(a)	(i)	What is the average relative atomic mass of titanium? Give your answer to three significant figures.

(2)
 
(ii)	What instrument would have been used to find this percentage composition?
............................................................................................................................
(1)
 
(b)	(i)	Give the electronic configuration of a titanium atom, using s p d notation.
............................................................................................................................
(2)
(ii)	Name the part of the Periodic Table where titanium appears.
............................................................................................................................
(1)
 
(c)	Titanium occurs naturally as rutile, TiO2. One possible method of obtaining pure titanium is to heat rutile with carbon.
TiO2(s)  +  2C(s)  ®  Ti(s)  +  2CO(g)
(i)	What type of reaction is this?
............................................................................................................................
(1)

(ii)	Calculate ΔH for this reaction given that
ΔHfο [TiO2(s)] = –940 kJ mol–1
ΔHfο [CO(g)] = –110 kJ mol–1
	Include a sign and units in your answer.

(3)
 
(iii)	Name the law you have used in your calculation.
............................................................................................................................
(1)
 
(iv)	When titanium is manufactured by this method, explain what pollution problem arises.
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(2)
(Total 13 marks)

10.	This question is about nitrogen trifluoride, NF3, and nitrogen trichloride, NCl3, which are covalent compounds. Van der Waals attractions and permanent dipole–dipole attractions exist between molecules of both compounds in the liquid state.
(a)	(i)	Describe how van der Waals attractions are caused.
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(1)
 
(ii)	In which of the two compounds would you expect there to be greater van der Waals attractions? Justify your answer.
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(1)
 
(iii)	In which of the two compounds would you expect there to be the greater permanent dipole–dipole attractions? Justify your answer.
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 (1)
 
(iv)	The boiling point of nitrogen trichloride is much higher than that of nitrogen trifluoride. How might this be explained in terms of the two types of intermolecular attractions?
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(1)
 
(b)	The standard enthalpy change for the formation of gaseous nitrogen trifluoride is
–125 kJ mol–1.
N2 (g)  +  1F2 (g)  ®  NF3 (g)  DHοf = –125 kJ mol–1
	The standard molar enthalpy changes of atomisation of nitrogen, N2, and of fluorine, F2, are given below.
	ΔHat/kJ mol–1
N2 (g) ® N (g)	+ 473
F2 (g) ® F (g)	  + 79

	This information can be represented on a Hess cycle in the following way, and then used to calculate bond energies.

(i)	Insert formulae, showing the correct quantities of each element, into the appropriate boxes.
(1)
(ii)	Insert arrows between the boxes and write the correct numerical data alongside the appropriate arrows.
(2)
 
(iii)	Use the cycle to calculate the N ¾ F bond energy in nitrogen trifluoride.

(2)
(Total 9 marks)

11.	Urea, which is used as a fertillser in much of mainland Europe, Asia and Africa, is manufactured by the reaction of ammonia and carbon dioxide.
2NH3(g)  +  CO2(g)  ®  NH2CONH2(s)  +  H2O(l)

(a)	Define the term standard enthalpy of formation, DHf , of urea.
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(3)
 
(b)	Calculate the enthalpy change, DH for the reaction above, given the following standard enthalpies of formation.
Substance	DHf / kJ mol–1	
NH3(g)	–46.2	
CO2(g)	–393.5	
NH2CONH2(s)	–632.2	
H2O(l)	–285.8	

(3)
(Total 6 marks)

12.	(a)	Halogenoalkanes react with many nucleophiles.
	Define the term nucleophile.
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(2)
 
(b)	(i)	Identify the reagent and conditions necessary for the conversion of iodoethane to ethylamine, C2H5NH2.
Reagent:.........................................................................................................
Conditions:.....................................................................................................
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(3)
 
(ii)	State why the rate of reaction would be slower if bromobutane were used in place of iodoethane, with all other conditions remaining the same.
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(1)

(c)	Iodoethane reacts with water to form ethanol and hydrogen iodide.
C2H5I + H2O ® C2H5OH + HI	DHf  = +36 kJ mol–1
	Use some or all of the data below to calculate the C¾I bond enthalpy.
Bond	Bond enthalpy
/ kJ mol–1	Bond	Bond enthalpy
/ kJ mol–1	
C¾H	413	H¾I	298	
C¾C	347	C¾O	358	
H¾O	464			

(3)
 
(d)	Ethanol was heated under reflux with an excess of a mixture of potassium dichromate(VI) and dilute sulphuric acid. Draw the full structural formnula of the organic product.

(1)
(Total 10 marks)

13.	The apparatus shown in the diagram below may be used to find the enthalpy of combustion of alcohols.

	Using the apparatus, a student recorded the results included in the table below.
Alcohol = ethanol, C2H5OH
Molar Mass (C2H5OH) = 46.0 g mol–1
Volume of water in beaker = 200 cm3
\ mass of water in beaker = 200 g
Weighings
Spirit lamp + ethanol before combustion = 198.76 g
Spirit lamp + ethanol after combustion = 197.68 g
Temperatures
Water before heating = 19.5 °C
Water after heating = 38.1 °C
Specific heat capacity of water = 4.18 J g–1 °C–1	
 
(a)	What assumption is the student making about water to be able to state that its mass is numerically equal to its volume?
...............................................................................................................................................
(1)

(b)	Calculate the heat gained by the water. Give your answer in kJ.

(2)
 
(c)	Calculate the amount (number of moles) of ethanol used.

(2)
 
(d)	Using your values from (b) and (c), calculate the enthalpy of combustion of ethanol. Give your answer to a number of significant figures consistent with the readings in the table. Include a sign and units in your answer.

(3)

(e)	The student’s evaluation of the experiment is given below.
My calculated value of the enthalpy of combustion was
numerically much less than the data book value. The
reasons for my low value include:
1	heat losses to the surrounding air;
2	when I re-checked the mass of the spirit lamp and
ethanol after combustion, I noticed that it had lost
mass even when it was not being used;
3	a black solid which formed on the base of the beaker.	
 
(i)	Explain why the spirit lamp and ethanol lost mass even when not in use.
.....................................................................................................................................
.....................................................................................................................................
(1)
 
(ii)	Suggest the identity of the black solid. Explain why its formation will lead to a low value for the enthalpy of combustion.
Identity ........................................................................................................................
Explanation ................................................................................................................
.....................................................................................................................................
(2)
(Total 11 marks)

14.	This question is about a self-heating can of coffee.
	The bottom of the can has a compartment containing copper(II) nitrate solution. When a button on the bottom of the can is pressed, magnesium powder is released into the compartment where it reacts with the copper(II) nitrate solution.

(a)	(i)	Write an ionic equation for the reaction between magnesium powder and copper(II) ions. Include state symbols, but omit any spectator ions.

(2)
 
(ii)	Show how the standard enthalpy change for this reaction could be calculated from the standard enthalpies of formation of copper(II) ions and magnesium ions. You should include a Hess cycle in your answer.

(3)
 
(b)	The can contains 150 g of a solution of coffee in water.
	The temperature of the solution needs to increase by 60 °C to produce a hot drink.
(i)	Calculate the energy change needed to produce a temperature increase of 60 °C in the coffee, using the relationship 
Energy change = 4.2 × mass of solution × temperature change.
	Remember to include a unit in your answer.

(2)

(ii)	The standard enthalpy change for this reaction is –530 kJ mol–1.
	Calculate the number of moles of reactants needed to produce the energy change in (i).

(1)
 
(iii)	A solution of copper(II) nitrate of concentration 8.0 mol dm–3 is used.
	Use your answer to (ii) to calculate the volume, in cm3, of copper(II) nitrate solution needed.
	Your answer should be given to two significant figures.

(1)
 
(c)	Suggest TWO reasons why the temperature of the coffee may not increase by as much as 60 °C.
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(2)
(Total 11 marks)

15.	In two similar, separate experiments the enthalpy changes for the reactions of sodium hydrogencarbonate and sodium carbonate with excess dilute hydrochloric acid were determined.
(a)	The first experiment was to find the enthalpy change, DH1, for the reaction
NaHCO3(s) + HCl(aq) ® NaCl(aq) + CO2(g) + H2O(l)
Measurement	Reading	
Mass of solid sodium hydrogencarbonate added to
hydrochloric acid.	5.00 g	
Volume of hydrochloric acid	50.0 cm3	
Temperature of hydrochloric acid before addition
of solid sodium hydrogencarbonate	22.0 °C	
Final temperature of solution	15.5 °C	
Molar mass of sodium hydrogencarbonate	84.0 g mol–1	
Specific heat capacity of solution	4.18 J g–1 °C–1	
 
(i)	Calculate the amount (moles) of sodium hydrogencarbonate used.

(1)

(ii)	Calculate the heat absorbed in the reaction in kJ.
[Assume that 1 cm3 of solution has a mass of 1 g]

(2)
 
(iii)	Calculate the value of DH1 in kJ mol–1. Include a sign in your answer expressing it to a number of significant figures suggested by the data in the table.

(2)

(b)	In the second experiment the enthalpy change for the reaction between sodium carbonate and dilute hydrochloric acid was measured.
Na2CO3(s) + 2HCl (aq) ® 2NaCI(aq) + CO2(g) + H2O(l)
	The molar enthalpy change, DH2, was calculated to be –35.6 kJ mol–1
(i)	Give TWO ways in which the temperature change differs when equal molar amounts of sodium hydrogencarbonate and sodium carbonate react separately with the same volume of hydrochloric acid.
.....................................................................................................................................
.....................................................................................................................................
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(2)
 
(ii)	Give ONE assumption that has been made in calculating the values of DH1, and DH2 from experimental results.
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(1)
(Total 8 marks)

16.	In the manufacture of beer, brewers often add small amounts of salts of Group 2 elements to the water used. These salts influence the chemical reactions during the brewing process.
Two such salts are calcium sulphate and magnesium sulphate.
(a)	A flame test can be used to confirm that a sample of a salt contains calcium ions.
(i)	Describe how you would carry out a flame test.
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(3)
 
(ii)	A positive test results in a brick-red flame colour. Describe the changes that occur in calcium ions to produce a colour.
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(2)
 
(iii)	Impurities in the salt may lead to other colours being observed in the flame.
What metal ion is likely to be present if a yellow flame is seen?
............................................................................................................................
(1)

(b)	Magnesium sulphate can be used in its anhydrous form, MgSO4(s), or in its hydrated form, MgSO4.7H2O(s).
	An experiment was carried out to find the enthalpy change when hydrated magnesium sulphate dissolved completely in water.
MgSO4.7H2O(s)    MgSO4(aq)  +  7H2O(l)
	12.3 g of hydrated magnesium sulphate was added to 100 g of water in a simple calorimeter and the temperature was found to fall by 1.1 °C.
 
(i)	Calculate the energy change, in joules, that occurred in the experiment, using the relationship
Energy change (J) = 4.18 × mass of water × temperature change

(2)
 
(ii)	Calculate the number of moles of hydrated magnesium sulphate used in the experiment. Use the Periodic Table as a source of data.

(2)
 
(iii)	Use your answers to (i) and (ii) to calculate the enthalpy change for the reaction.
Include a sign and units in your final answer, which should be given to 2 significant figures.

(2)
 
(c)	The enthalpy change as hydrated magnesium sulphate is converted to anhydrous magnesium sulphate is very difficult to measure. The Hess Cycle below can be used to find this enthalpy change, ΔHr.

(i)	Use the cycle to write an expression for ΔHr using ΔH1 and ΔH2.

(1)
 
(ii)	Use your expression in (c)(i) and your answer from (b)(iii) to calculate ΔHr.
	Include a sign and units in your final answer, which should be given to 2 significant figures.

(2)
(Total 15 marks)

17.	Phosphine, PH3, is a hydride of the Group 5 element, phosphorus.
(a)	(i)	Draw a ‘dot-and-cross’ diagram of a phosphine molecule. You should include only outer shell electrons.

(1)

(ii)	Draw the shape you would expect for the phosphine molecule, suggesting a value for the HPH bond angle.

HPH bond angle .......................................................................................................
(2)
 
(iii)	Explain the shape of the phosphine molecule you have given in your answer in (ii).
	Justify your value for the HPH bond angle.
.............................................................................................................................
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(2)
 
(b)	(i)	Write a balanced equation, including state symbols, for the atomisation of phosphine gas.
.............................................................................................................................
(1)

(ii)	Use your answer to (i) and the data below to calculate the standard enthalpy change of atomisation of phosphine at 298 K. Include a sign and units in your answer.
ΔHοf[PH3(g)]	=	+	5.4 kJ mol-1
ΔHοat[½H2(g)]	=	+	218.0 kJ mol-1
ΔHοat[P(s)]	=	+	314.6 kJ mol-1

(3)
 
(iii)	Calculate a value for the bond energy of the bond between phosphorus and hydrogen, using your answer to (ii).

(1)
(Total 10 marks)

18.	Methane, CH4, is used as a domestic and industrial fuel and as a reagent in the petrochemical industry.
(a)	Define the term standard enthalpy of combustion.
.......................................................................................................................................
.......................................................................................................................................
.......................................................................................................................................
(3)

(b)	Methane burns in oxygen according to the equation:

	Use the average bond enthalpy data shown below to calculate the enthalpy change of this reaction.
Bond	Bond enthalpy/kJ mol–1	
C¾H	+435	
O==O	+498	
C==O	+805	
H¾O	+464	

(3)

(e)	Methane is the feedstock in the manufacture of hydrogen according to the equation:
CH4(g) + 2H2O(g) ® CO2(g) + 4H2(g)
	Given the enthalpy of formation data below, draw a labelled Hess’s law cycle and use it to calculate the enthalpy change of this reaction.
Substance	Enthalpy of formation/kJ mol–1	
CH4(g)	–75	
CO2(g)	–394	
H2O(g)	–242	

(4)
(Total 10 marks)

19.	In an experiment to find the enthalpy of neutralisation of a monobasic acid, HX, with an alkali, the following procedure was followed:
Step 1	25.0 cm3 of 1.00 mol dm–3 dilute aqueous acid, HX, was measured into a polystyrene cup.
Step II	A 0-100 °C thermometer was placed in the acid. The temperature of the acid was immediately read and recorded.
Step III	5.00 cm3 portions of aqueous sodium hydroxide were added to the acid from a burette. After each addition, the temperature of the solution was read and recorded. The thermometer was removed and rinsed with water between each addition. A total of 50.0 cm3 of aqueous sodium hydroxide was added.

(a)	Suggest ONE change that could be made at Step II and ONE change that could be made at Step III to improve the accuracy of the experiment.
Step II ...................................................................................................................................
...............................................................................................................................................
Step III .................................................................................................................................
...............................................................................................................................................
(2)
 
(b)	The readings of temperature and volume are plotted on the grid. Draw two separate straight lines of best fit, extending the two lines so that they intersect.

(2)

(c)	From the graph, read off the maximum temperature rise, DT, and the volume of aqueous sodium hydroxide added at neutralisation, VN.
DT = ..................................... °C	VN = ................................ cm3
(2)
 
(d)	(i)	Use the formula below to calculate the heat evolved in the neutralisation.
Heat evolved = kJ

(1)
 
(ii)	Given that the amount (moles) of acid neutralised was 0.025 mol, calculate the enthalpy of neutralisation, DHneut, in units of kJ mol–1.

DHneut = ............................... kJ mol–1
(2)
(Total 9 marks)

20.	A reaction of ammonium dichromate(VI) is shown by the following equation.
(NH4)2Cr2O7(s)  N2(g) + 4H2O(g) + Cr2O3(s)
(a)	What type of reaction is this?
......................................................................................................................................
(1)
 
(b)	The enthalpy change for this reaction can be calculated from standard enthalpy changes of formation.
(i)	State fully what is meant by the standard enthalpy change of formation, ΔHfο, 
of a compound.
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
(3)
 
(ii)	Complete the Hess cycle for the reaction so that you can calculate the enthalpy change of the reaction from standard enthalpy changes of formation.

(3)
 
(iii)	What is the value of ΔHfο[N2(g)]? ......................................................................
(1)

(iv)	Calculate ΔHοr for the reaction using the following data. Remember to include a sign and units in your answer.
ΔHfο[(NH4)2Cr2O7(s)]	=	–1810 kJ mol–1
ΔHfο[H2O(g)]	=	–242 kJ mol–1
ΔHfο[Cr2O3(s)]	=	–1140 kJ mol–1

(3)
 
(c)	In this reaction, water vapour is formed which condenses to liquid water on cooling.
Is this reaction H2O(g) ® H2O(l) exothermic or endothermic?
	Justify your answer.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(Total 13 marks)

21.	Chlorine can be converted to the gas chlorine(I) oxide, Cl2O.
	The standard molar enthalpy change of formation of chlorine(I) oxide and the standard molar enthalpy changes of atomisation of chlorine and oxygen are given below:
ΔHfο [Cl2O(g)]	=	+ 80.3 kJ mol–1
ΔHatο [½Cl2(g)]	=	+121.7 kJ mol–1
ΔHatο [½O2(g)]	=	+249.2 kJ mol–1
	A partially completed Hess cycle involving chlorine(I) oxide is shown below:

(i)	Insert the appropriate formulae, showing the correct quantities of each element, into the box above. Include state symbols in your answer.
(1)
(ii)	Insert arrows between the boxes, writing the correct numerical data alongside the appropriate arrows.
(1)

(iii)	Use the cycle to calculate ΔHatο  [Cl2O(g)].

(1)
(iv)	Calculate the Cl—O bond energy in chlorine(I) oxide.

(1)
(Total 14 marks)

22.	(a)	Enthalpy changes can be calculated using average bond enthalpy data.
(i)	The enthalpy change to convert methane into gaseous atoms is shown below.
CH4(g) → C(g) + 4H(g)	∆H = +1664 kJ mol–1
	Calculate the average bond enthalpy of a C—H bond in methane.

(1)

(ii)	Use the data in the table below and your answer to (a)(i) to calculate the enthalpy change for
2C(g) + 2H2(g) + Br2(g) → CH2BrCH2Br(g)
Bond	Average bond
enthalpy / kJ mol–1	Bond	Average bond
enthalpy / kJ mol–1	
C—C	+348	H—H	+436	
Br—Br	+193	C—Br	+276	

(3)
 
(b)	The standard enthalpy of formation of 1,2-dibromoethane, CH2BrCH2Br, is
–37.8 kJ mol–1.
	Suggest the main reason for the difference between this value and your calculated value in (a)(ii).
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(1)
(Total 5 marks)

23.	In an experiment to find the enthalpy change for the reaction
Zn(s) + Cu2+(aq) ® Zn2+(aq) + Cu(s)
	a student was given the following list of instructions:
·	weigh out 5.0 g of zinc powder into a weighing bottle
·	use a measuring cylinder to transfer 50 cm3 of 1.0 mol dm–3 aqueous copper(II) sulphate into a polystyrene cup, firmly held in a 250 cm3 beaker
·	stir the solution with the thermometer and record the temperature to the nearest 0.5 °C
·	continue to stir the solution, recording its temperature every minute
·	at exactly 3.5 minutes, add the zinc powder to the aqueous copper(II) sulphate, stirring continuously
·	record the temperature of the solution every minute from 4.0 to 9.0 minutes.
 
	The temperature readings obtained are shown in the table below.
Time/min	0	1.0	2.0	3.0	4.0	5.0	6.0	7.0	8.0	9.0	
Temperature/°C	20.0	20.0	20.0	20.0	63.0	60.5	59.0	57.0	55.5	53.0	

(a)	(i)	Plot a graph of temperature against time on the grid below.

(2)
 
(ii)	Use the graph to calculate the maximum temperature change, DT. Show clearly on the graph how you obtained your answer.
DT = ................................................°C
(2)

(iii)	Give ONE reason why a series of temperature readings is obtained instead of just the starting and maximum temperatures.
............................................................................................................................
............................................................................................................................
(1)
 
(b)	(i)	Calculate the heat change, in joules.
	The specific heat capacity of the solution is 4.18 J g–1 °C–1.

(1)
 
(ii)	What assumption have you made about the solution in your calculation in (i)?
............................................................................................................................
............................................................................................................................
(1)
 
(iii)	Calculate the amount (moles) of copper (II) sulphate, CuSO4, in 50 cm3 of a 1.0 mol dm–3 solution.

(1)

(iv)	Calculate the enthalpy change for this reaction in kJ mol–1.

(2)
 
(c)	Suggest TWO improvements that could be made to the experimental procedure.
	Give a reason for each.
Improvement 1 ............................................................................................................
......................................................................................................................................
......................................................................................................................................
Reason .........................................................................................................................
......................................................................................................................................
......................................................................................................................................
Improvement 2 ............................................................................................................
......................................................................................................................................
......................................................................................................................................
Reason ..........................................................................................................................
......................................................................................................................................
......................................................................................................................................
(4)
(Total 14 marks)

24.	When solutions of potassium carbonate and calcium chloride are mixed together, the following reaction takes place 
CaCl2(aq) + K2CO3(aq) → CaCO3(s) + 2KCl(aq)
(a)	Re-write the above equation as an ionic equation. Include state symbols, but omit any spectator ions.

(2)
 
(b)	An experiment was carried out to measure the enthalpy change for this reaction. 50 cm3 of a 1.00 mol dm–3 solution of potassium carbonate was added to 50 cm3 of a 1.00 mol dm–3 solution of calcium chloride. The temperature fell by 1.5 °C.
(i)	Calculate the energy taken in from the surroundings using the relationship
		energy	=	mass of	×	specific heat capacity	×	temperature
			solution		of solution		change
	/J		/g		/J g–1 °C–1		/°C
	You may assume that
·	1.0 cm3 of solution has a mass of 1.0 g.
·	The specific heat capacity of the solution is 4.2 J g–1 °C–1.

	Energy taken in = .................... J
(1)
 
(ii)	How many moles of calcium chloride are used in this experiment?

(1)
 
(iii)	Calculate the enthalpy change for the reaction, giving your answer to two significant figures. Include a sign and units in your answer.

(2)
 
(iv)	Which measurement is likely to have caused the major source of error in this experiment? Explain your answer.
................................................................................................................................
................................................................................................................................
................................................................................................................................
................................................................................................................................
(1)
 
(v)	What apparatus should be used to contain the reaction mixture during this experiment?
................................................................................................................................
(1)
 
(c)	If the experiment in (b) was repeated, but using only 25 cm3 of each solution, predict what the fall in temperature would be.
.......................................................................................................................................
(1)
(Total 9 marks)

25.	(a)	Calculate the number of atoms in 3.50 g of lithium.
	Use the Periodic Table as a source of data.
	[The Avogadro constant, L = 6.02 × 1023 mol–1]

(2)
 
(b)	The equation for the reaction of lithium with hydrochloric acid is shown below.
2Li(s) + 2HCl(aq) ® 2LiCl(aq) + H2(g)
(i)	Rewrite this equation as an ionic equation, omitting the spectator ions.

(1)

(ii)	Draw a ‘dot and cross’ diagram of lithium chloride showing all the electrons. Indicate charges clearly on your diagram.

(2)
 
(iii)	The value of the standard enthalpy change for the reaction, DHο, is –557 kJ mol–1. State TWO of the reaction conditions necessary for this enthalpy change to be standard.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
(Total 7 marks)

26.	The equation below shows a reaction of magnesium nitrate.
2Mg(NO3)2(s)  2MgO(s) + 4NO2(g) + O2(g)
	The Hess cycle below can be used to find the enthalpy change for the reaction in (a), ΔHr, using enthalpy changes of formation.

(i)	Complete the empty box in the cycle.
(2)
 
(ii)	Use the cycle and the data below to calculate the enthalpy change of the reaction, ΔHr.
	ΔHf
	/ kJ mol–1
Mg(NO3)2	–790.7
MgO	–601.7
NO2	+33.2

 

(3)
(Total 5 marks)

27.	(a)	Define the term standard enthalpy of formation.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
 
(b)	The dissociation of phosphorus pentachloride is a reversible reaction.
PCl5(g)  PCl3(g) + Cl2(g)
 
(i)	Use the values of enthalpy of formation given to calculate ∆H for the forward reaction.
	∆Hf /
kJ mol–1	
PCl5(g) 	– 399 	
PCl3(g) 	– 306 	

(1)
 
(ii)	Explain, with reasons, the effect that raising the temperature would have on the composition of the equilibrium mixture.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)

(iii)	Other than by changing the temperature, suggest how the amount of PCl5 present at equilibrium could be increased. Give a reason for your answer.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
(Total 8 marks)

28.	The enthalpy change for the thermal decomposition of calcium carbonate cannot be measured directly, but can be found by carrying out two reactions as shown in the Hess cycle below.

(a)	Suggest ONE reason why it is difficult to measure Hreaction directly by experiment.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)

(b)	In an experiment to find H1 a student added 2.00 g of finely powdered calcium carbonate to 20.0 cm3 of 2.50 mol dm–3 hydrochloric acid solution (an excess) in a polystyrene container. The temperature rose from 20.5 °C to 23.0 °C.
(i)	Why is the calcium carbonate used in this experiment finely powdered, rather than in lumps? Explain why this is important for an accurate result.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(ii)	Calculate the energy change using the relationship below.
Energy change	=	4.2	×	mass of solution	×	temperature change
	/J	/J g–1 K–1	/g	/K
	Assume that the mass of the solution is 20 g.

(1)

(iii)	Calculate the enthalpy change, DH1. Include a sign and units in your answer.
[The molar mass of CaCO3 is 100 g mol–1]

(3)

(iv)	In another experiment, the value of H2 was found to be –181 kJ mol–1.
	Use this result and your answer to (iii) to calculate the value of Hreaction.

(2)
 
(c)	The student checked the experimental results using information from the Book of data in another Hess cycle.

	Name the enthalpy change represented by DH3.
.....................................................................................................................................
(1)
(Total 10 marks)

29.	The Hess cycle below can be used to estimate the enthalpy change of formation, DHf, of the unstable gaseous compound with the formula HOCl(g).

(a)	(i)	Insert formulae, with state symbols, into the appropriate boxes, to show the correct quantities of each element.
(1)
 
(ii)	Use the cycle to calculate a value for the enthalpy change of formation, [HOCl(g)].

(1)

(iii)	Assuming that the H—O bond energy is +464 kJ mol–1, calculate a value for the O—Cl bond energy.

(1)

(b)	(i)	Draw a ‘dot and cross’ diagram for the HOCl molecule showing outer electrons only.

(2)
(ii)	Predict the HOCl bond angle. Justify your answer.
Angle ................................................................................................................
Justification ......................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)

(c)	HOCl(g) can be made from chlorine(I) oxide by the reversible reaction
Cl2O(g) + H2O(g)  2HOCl(g)
	What effect, if any, would an increase in pressure have on the proportion of HOCl(g) at equilibrium? Justify your answer.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 9 marks)

30.	The enthalpy change for the reaction of anhydrous aluminium chloride, AlCl3, with water can be found as follows:
•	Add about 100 cm3 of distilled water to a weighed polystyrene cup.
•	Measure the steady temperature of the water.
•	Add anhydrous aluminium chloride to the polystyrene cup, with stirring.
•	Measure the highest temperature reached.
•	Re-weigh the polystyrene cup and contents
Data
Mass of anhydrous aluminium chloride	= 4.00 g
Mass of solution	= 104 g
Initial temperature	= 17.5 °C
Highest temperature reached	= 43.5 °C
Specific heat capacity of the solution	= 4.09 J g–1 °C–1	

(i)	Calculate the heat change in this experiment.

(2)
 
(ii)	Assuming that 100 cm3 of water is a large excess, calculate the enthalpy change, in kJ mol–1, when one mole of aluminium chloride reacts. Include a sign and unit in your answer.

(3)
(Total 5 marks)

31.	An experiment was carried out to find the enthalpy change for the reaction of zinc powder with copper(II) sulphate solution.
	Zn(s) + CuSO4(aq) ® ZnSO4(aq) + Cu(s)
	50cm3 of copper(II) sulphate solution, of concentration 1.0 mol dm–3, was put into a polystyrene cup and the temperature of the solution measured. After one minute, 5.0 g of zinc powder was added, the mixture stirred with a thermometer and the temperature measured every 30 s.
 
(a)	(i)	What is meant by a spectator ion?
...........................................................................................................................
...........................................................................................................................
(1)

(ii)	Give the formula of the spectator ion in this reaction.
...........................................................................................................................
(1)
 
(iii)	Write the equation for this reaction, omitting the spectator ion.

(1)
 
(b)	How would you measure the 50 cm3 of copper(II) sulphate solution?
.....................................................................................................................................
.....................................................................................................................................
(1)
 
(c)	Give TWO reasons why it is better to use a polystyrene cup, rather than a metal container, to obtain more accurate results.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)

(d)	Calculate the number of moles of each of the reactants and hence deduce which reactant is completely used up.
Use the Periodic Table as a source of data.
	Moles of zinc powder

	Moles of copper(II) sulphate

	Reactant used up .................................................................
(3)
 
(e)	The following results were obtained.
Time /s	0	60	90	120	150	180	210	
Temperature /°C	22	22	60	65	63	61	59	

(i)	On the graph paper below, plot the results of this experiment.

(2)
 
(ii)	Explain the shape of your graph
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(iii)	The maximum recorded temperature in this experiment was 65°C. Use your graph to estimate a more accurate maximum temperature.
...........................................................................................................................
(1)
 
(f)	(i)	Calculate the energy change in this experiment using your answer to (e)(iii) and the relationship
	energy change =	mass of	×	specific heat capacity	×	temperature rise
		solution		of solution
	/J	/g		/J °C–1 g–1		/°C
	You may assume that
•	1.0 cm3 of solution has a mass of 1.0 g
•	The specific heat capacity of the solution is 4.2 J °C–1 g–1

(1)
 
(ii)	Use your answers to (d) and (f)(i) to calculate ∆H for this reaction. Include a sign and units in your answer.

(3)
(Total 18 marks)

32.	This question is about the chemistry of propane, C3H8.
	Propane is sold for use as a fuel for camping stoves. On complete combustion it forms carbon dioxide and water.
(a)	The enthalpy change of combustion of propane, ΔHc, can be measured by burning a known mass of propane below a container of water and measuring the temperature rise of the water.
	The heat capacity of the apparatus (the energy required to raise the temperature of the apparatus by 1 °C) is found by calibrating it with a fuel with known enthalpy change of combustion.
	The results of an experiment are shown below.
	Mass of propane burned	0.500 g
Temperature of water at start	21.0 °C
Final temperature of water	39.0 °C
Heat capacity of apparatus	1.35 kJ °C–1
 
(i)	Calculate the number of kilojoules of energy transferred when the 0.500 g sample of propane burns in this experiment.
(1)

(ii)	Use your answer to (i) to calculate ΔHc for propane in kJ mol–1. Give your answer to three significant figures.
	Use the Periodic Table as a source of data.

(2)
(iii)	The Book of data gives the value of ΔHc for propane as –2220 kJ mol–1.
	Calibrating the apparatus means that the answer you calculated in (ii) allows for errors due to heat loss.
	Suggest the other main source of error which makes the experimental result different from the data book value.
...........................................................................................................................
...........................................................................................................................
(1)

(b)	A value of ΔHc for propane can be calculated using mean bond energies and the Hess cycle below.
(i)	Complete the Hess cycle, and use the mean bond energies to calculate ΔH1.
Hence calculate ΔHc.
	Mean bond energies
/kJ mol–1	
C=O	805	
H–O	464	

ΔHc = ................................................................................................................
(3)
 
(ii)	Give ONE reason why the value you calculated in (b)(i) also differs from the value for the heat of combustion of propane in the Book of data.
...........................................................................................................................
...........................................................................................................................
(1)
 
(c)	When propane reacts with chlorine in the presence of ultraviolet light one of the products is 2-chloropropane.
(i)	Name the mechanism and type of this reaction.
Mechanism .......................................................................................................
Type ..................................................................................................................
(2)
 
(ii)	In this reaction a small quantity of an alkane, C6H14, is produced.
	Explain how this occurs. Include an equation in your answer.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(d)	2-Chloropropane and 2-iodopropane are both colourless liquids at room temperature.
They can be distinguished by their reactions with aqueous silver nitrate.
(i)	What would you see when the reaction is carried out with each halogenoalkane?
2-chloropropane ...............................................................................................
2-iodopropane ..................................................................................................
(2)
 
(ii)	Write an ionic equation showing how silver ions react in the mixture made from 2-iodopropane and aqueous silver nitrate. Include state symbols in your answer.

(2)
 
(iii)	Both 2-chloropropane and 2-iodopropane form the same organic product in the reaction with aqueous silver nitrate.
	Name, or give the structural formula of, this organic product.
...........................................................................................................................
(1)
(Total 17 marks)

33.	This question is about some of the chemicals used in car engines and their reactions.
(a)	Compound X, shown below, is one component of petrol.
	
(i)	Name X.
...........................................................................................................................
(1)
(ii)	Give the empirical formula of X.
...........................................................................................................................
(1)
(iii)	X can be made by cracking decane, C10H22.
	Assuming only one other product forms in a cracking reaction, deduce the molecular formula of this other product.

(1)
 
(iv)	What is the sign of the enthalpy change for the reaction in which decane is cracked? Give a reason for your answer.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)

(v)	If the air supply in a car engine is poor, there is not enough air for carbon dioxide to be produced.
	Use this information to suggest ONE possible equation for the combustion of X in this engine. Use the molecular formula of X in your equation.

(2)
 
(b)	When air enters a car engine, as well as the fuel burning, nitrogen and oxygen can react to form nitrogen(II) oxide.
	N2(g) + O2(g)  2NO(g)	ΔH = + 180 kJ mol–1
(i)	What, if any, is the effect on the percentage of nitrogen(II) oxide in an equilibrium mixture of these three gases if the pressure and temperature are increased?
Explain your answers.
	Increase in pressure
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
	Increase in temperature
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(ii)	In a car exhaust pipe, nitrogen(II) oxide passes over a catalytic converter.
The following reaction occurs.
	2NO(g) + 2CO(g) → N2(g) + 2CO2(g)	ΔH = –746 kJ mol–1
	Explain why this reaction speeds up when the car engine has been running for a few minutes.
...........................................................................................................................
...........................................................................................................................
(1)
 
(iii)	A textbook says “The catalytic converter converts the gases coming out of the engine into less harmful ones”.
	State, with a reason, which of the four gases in the equation in (ii) you consider to be least harmful.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)
 
(iv)	The diagram below shows the reaction profile for the change which occurs in the catalytic converter.

	On the diagram, show the activation energy, EA.
	Add a line showing the reaction profile if no catalyst is present.
(2)
(Total 12 marks)

34.	(a)	Define the term standard enthalpy of combustion.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
 
(b)	The following standard enthalpies of combustion are needed to calculate the standard enthalpy of formation of ethanol, C2H5OH.
Substance	Standard enthalpy of
combustion /kJ mol–1	
carbon, C (s, graphite)	–394	
hydrogen, H2(g)	–286	
ethanol, C2H5OH (l)	–1371	
 
(i)	Complete the Hess’s Law cycle by filling in the box and labelling the arrows with the enthalpy changes.

(3)

(ii)	Use your Hess’s Law cycle to calculate the standard enthalpy of formation of ethanol.

(2)
(Total 8 marks)

35.	The apparatus used and the recordings made by a student, carrying out an experiment to determine the enthalpy of combustion of methanol, are shown below.
	Diagram
	
	Results
Molar mass (methanol) = 32 g mol–1
Volume of water in beaker = 50 cm3
Mass of water in beaker = 50 g
Weighings
Spirit lamp + methanol before combustion = 163.78 g
Spirit lamp + methanol after combustion = 163.44 g
Temperatures
Water before heating = 22.0 °C
Water after heating = 43.5 °C
Specific heat capacity of water = 4.18 J g–1 °C–1	

	Observations
•	When the spirit lamp was being weighed its mass was continually falling.
•	A black substance formed on the bottom of the beaker as the methanol burned.
 
(a)	(i)	Calculate the amount (moles) of methanol, CH3OH, burned.

(2)
 
(ii)	Calculate the heat gained by the water. Give your answer in kJ.

(2)
 
(iii)	Use your values from (i) and (ii) to calculate the enthalpy of combustion of methanol in kJ mol–1. Include a sign with your answer.

	∆H = ........................................ kJ mol–1
(2)

(b)	(i)	The thermometer used in the experiment can be read to an accuracy of ±0.5 °C.
Calculate the percentage error in the temperature change.

(1)
 
(ii)	Calculate the maximum temperature change that could have occurred during the experiment.

(1)
 
(c)	(i)	Give a reason why the mass of the spirit lamp fell as it was being weighed.
...........................................................................................................................
...........................................................................................................................
(1)
 
(ii)	Suggest the identity of the black substance that forms on the beaker. State the effect on the value of the enthalpy of combustion obtained.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
(Total 11 marks)

36.	Calcium hydroxide decomposes on strong heating to form calcium oxide and water.
	Ca(OH)2(s) → CaO(s) + H2O(l)
	Two samples of calcium hydroxide were taken, each weighing exactly 1.00 g.
	The first sample was cautiously added to 25.0 cm3 of dilute hydrochloric acid contained in a glass beaker. The temperature rise was measured and found to be 16.5 °C.
	The other sample was heated for some time. It was then allowed to cool and then added to another 25.0 cm3 portion of hydrochloric acid as before. In this case the temperature rose by 25.5 °C.
	In both cases, the acid used was an excess.
 
(a)	(i)	Calculate the energy produced by the reaction of each solid with the acid.
	Use the relationship
	Energy produced	=	mass of solution	×	4.2	×	temperature rise
	/ J		/ g		/ J °C–1 g–1	/°C
	You may assume that 1.0 cm3 of solution has a mass of 1.0 g. Ignore the mass of the solid.
	For the solid calcium hydroxide

	For the solid calcium oxide

(1)

(ii)	How many moles of calcium hydroxide were used in each experiment?
[Molar mass of Ca(OH)2 = 74.0 g mol–1]

(1)
(iii)	Using your answers to (a)(i) and (ii), calculate the enthalpy changes for each reaction.
	Give your answers to two significant figures. Include the sign and units for each answer.
	For the solid calcium hydroxide, ΔH1

	For the solid calcium oxide, ΔH2

(2)
 
(b)	A Hess cycle for all these reactions is shown below.

(i)	Use this Hess cycle and your answers in (a)(iii) to calculate ΔHreaction. Include a sign and units.

(2)
 
(ii)	Apart from the approximations involved in using the equation given in (a)(i), give TWO other potential sources of error which are likely to affect the accuracy of the results.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
 
(iii)	Suggest why ΔHreaction is difficult to determine directly by experiment.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)
(Total 9 marks)

37.	(i)
Formula of substance	Standard enthalpy of formation / kJ mol–1	
CH4 (g)	–75	
O2 (g)	0	
CO2 (g)	–394	
H2O (g)	–242	
	The equation for the complete combustion of methane at 150 °C is:
	CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
	Use the given data to calculate the enthalpy of combustion of methane under these conditions.

(2)
 
(ii)	The standard enthalpy of combustion of methane is –891 kJ mol–1.
	Explain why this is very different from the value you have calculated in (a)(i).
.....................................................................................................................................
.....................................................................................................................................
(1)

(iii)	Here are some data about different fuels:
Name	Formula	Standard state	Energy evolved
per gram / kJ	Relative cost
per kJ	
hydrogen	H2	gas	143	2	
octane	C8H18	liquid	48.4	70	
	State and explain the advantages and disadvantages of using hydrogen compared with octane as a fuel for aeroplanes.
	Use the data in the table above and your knowledge of the combustion of fuels to help you.
.....................................................................................................................................
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.....................................................................................................................................
(4)
(Total 7 marks)

38.	The enthalpy change for the reaction between aqueous sodium hydroxide solution and aqueous hydrochloric acid was determined by the following method:
•	Aqueous hydrochloric acid was titrated against 25.0 cm3 of 1.50 mol dm–3 aqueous sodium hydroxide solution using a suitable indicator. The mean (or average) titre was 22.75 cm3.
•	25.0 cm3 of the sodium hydroxide solution was carefully measured into a polystyrene cup and 22.75 cm3 of the hydrochloric acid was transferred to a clean dry beaker.
Both solutions were allowed to stand for five minutes before their temperatures were noted.
•	The hydrochloric acid was then added to the sodium hydroxide solution, the mixture stirred thoroughly and the highest temperature noted.
•	The experiment was repeated three times giving an average temperature change of +10.5°C.

(a)	(i)	Calculate the heat produced in the reaction, in joules.
	Use the approximations that the density of the final solution is 1.00 g cm–3 and its specific heat capacity is 4.18 J g–1 K–1.

(2)

(ii)	Calculate the enthalpy change for the reaction, in kJ mol–1.

(3)

(b)	State ONE assumption made when calculating this enthalpy change, other than those stated in (a)(i).
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(1)
(Total 6 marks)

39.	Two experiments were carried out in order to calculate the enthalpy change of formation of magnesium carbonate, MgCO3.
	A Hess cycle for these reactions is shown below.

(a)	Complete the Hess cycle above for the formation of magnesium carbonate from its elements by balancing the equations and adding state symbols.
(2)

(b)	In Experiment 1 the temperature of 100 cm3 of hydrochloric acid was measured.
After one minute, 0.100 g of magnesium was added to the excess acid and the temperature measured every minute. The following results were obtained:
Time / min	0	1	2	3	4	5	6	
Temp / °C	21.0	21.0	25.3	25.1	24.9	24.8	24.7	

(i)	How many moles of magnesium were used in this experiment?
	Use the Periodic Table as a source of data.
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(1)

(ii)	The initial concentration of the hydrochloric acid was 2.00 mol dm–3.
	Calculate the number of moles of hydrochloric acid at the start and hence the number remaining at the end of the experiment.

(3)

(iii)	Plot the graph of temperature against time.

(2)

(iv)	Calculate the energy change in this experiment assuming the temperature rise is 4.5 °C. Use the expression
	Energy change (J) = 4.2 × mass of solution × temperature change
	[Assume that 1 cm3 of solution has a mass of 1 g]

(1)

(v)	Use your answer to (iv) to calculate ∆H1 for one mole of magnesium reacting with hydrochloric acid. Include a sign and units in your answer.

(2)

(vi)	Suggest why a temperature rise of 4.5 °C was used in the calculation in (iv).
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(1)

(c)	2.2 g of magnesium carbonate was added to 100 cm3 of the same acid in Experiment 2.
	The temperature changed from 21.0 °C to 23.5 °C resulting in an energy change of 1.05 kJ.
(i)	Calculate the mass of one mole of magnesium carbonate, MgCO3 and hence the number of moles of magnesium carbonate used in this experiment.
Use the Periodic Table as a source of data.

(2)

(ii)	Using the method in part (b)(v), calculate ∆H3.

(1)

(d)	Using your answers to (b)(v) and (c)(ii), calculate the enthalpy change of formation, ∆Hf, of magnesium carbonate, MgCO3.
Include a sign and units in your answer.

(2)
(e)	Why is it impossible to measure ∆Hf of MgCO3(s) directly?
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(1)
(Total 18 marks)

40.	Calculate the standard enthalpy change of formation of gaseous silicon tetrachloride,
ΔHοf [SiCl4(g)].
	Your answer should include a sign and units.
	Use the Hess cycle below and the following data at 298 K.
	[Si(s)]	= +455.6 kJ mol–1
	[½Cl2]	= +121.7 kJ mol–1
	Bond energy, E (Si-Cl)	= +407.4 kJ mol–1
	
(Total 3 marks)

41.	This question is about the manufacture of sulphuric acid, H2SO4.
(a)	The first stage in the manufacture of sulphuric acid is the combustion of sulphur.
	The following equation shows the reaction taking place when the standard enthalpy of combustion of sulphur is measured.
	S(s) + O2(g) → SO2(g)
	Define the term standard enthalpy of combustion.
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(3)

(b)	In the second stage of the manufacture of sulphuric acid, sulphur dioxide is oxidised to sulphur trioxide as shown in the following equation:
	2SO2(g) + O2(g)  2SO3(g)	ΔH = –196 kJ mol–1
(i)	Explain, in terms of collision theory, why the rate of a reaction is increased by increasing the temperature and by the addition of a catalyst.
Temperature
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Catalyst
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(4)
 
(ii)	State and explain the effect, if any, of increasing the temperature on the equilibrium yield of sulphur trioxide.
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(2)
 
(iii)	State and explain the effect, if any, of an increased pressure on the equilibrium yield of sulphur trioxide.
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(2)
 
(c)	The following equation represents the overall reaction for the manufacture of sulphuric acid from sulphur, oxygen and water.
	2S(s) + H2O + 3O2 ® 2H2SO4(l)
	Use the data below to calculate the enthalpy change for this reaction.
Substance	ΔHfο
/ kJ mol–1	
H2O (l)	–286	
H2SO4 (l)	–814	

(2)
(Total 13 marks)

42.	An experiment was carried out to measure the enthalpy change for the reaction of zinc with aqueous copper(II) sulphate.
	The equation for the reaction is
	Zn + CuSO4 → ZnSO4 + Cu
•	A measuring cylinder was used to transfer separate 50 cm3 samples of 1.25 mol dm–3 copper(II) sulphate solution into polystyrene cups.
•	Weighed amounts of zinc powder were added to each sample in turn.
•	Each mixture was stirred thoroughly and the temperature rise noted with a thermometer accurate to 0.5 °C.
	The results of this experiment are summarised on the graph below.

(a)	Explain why the graph initially shows a rise in temperature and then levels off.
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(2)

(b)	(i)	Suggest why the mass of metal is not used in the calculation of the heat change.
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(1)
 
(ii)	The graph shows that the maximum temperature change is 63.5 °C. Use this value to calculate the maximum heat change, in joules, in this reaction.
	You should assume that the density of the solution is 1.00 g cm–3 and its heat capacity is the same as water, 4.18 J g–1 °C–1.

(1)
(iii)	From the heat change calculated in (b)(ii) calculate the enthalpy change, in kJ mol–1, for the reaction. Include the appropriate sign and give your answer to three significant figures.

(4)

(c)	(i)	It is suggested that the precision of the experiment would be improved by using a thermometer accurate to 0.1 °C.
	Explain why this suggestion is incorrect.
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(1)
 
(ii)	Suggest a simple practical change to the method that would make the experiment more accurate.
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(1)
(Total 10 marks)

43.	The Hess cycle below can be used to find the enthalpy change, ∆Hr, for the reaction between hydrogen sulphide and sulphur dioxide, using standard enthalpy changes of formation.

(i)	Complete the cycle by filling in the empty box.
(2)

(ii)	What is meant by the standard enthalpy change of formation, ∆Hfο, of a compound?
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(3)
 
(iii)	Use the cycle and the data below to calculate the enthalpy change of the reaction, ∆Hr.
	∆Hfο / kJ mol–1	
SO2 (g)	–296.8	
H2S (g)	–20.6	
H2O (l)	–285.8	
(2)
(Total 7 marks)

44.	Which of the equations shown below represents the reaction for which DH is the standard enthalpy change of formation, DHοf 298, for ethanol, C2H5OH. Ethanol melts at 156 K and boils at 352 K.
A	2C(g) + 6H(g) + O(g) ® C2H5OH(g)
B	2C(s) + 3H2(g) + O2(g) ® C2H5OH(l)
C	2C(s) + 3H2(g) + O(g) ® C2H5OH(g)
D	2C(s) + 3H2(g) + ½O2(g) ® C2H5OH(l)
(Total 1 mark)

45.	Copper(II) sulfate solution can be prepared from solid copper(II) carbonate by reaction with hot dilute sulfuric acid.
(a)	Write the equation for the reaction, including state symbols.

(1)
 
(b)	The experiment was carried out using 0.025 moles of sulfuric acid of concentration 2.0 mol dm–3. What volume of this sulfuric acid was used?

(1)
 
(c)	(i)	It is usual to react the sulfuric acid with a slight excess of copper(II) carbonate.
	Calculate the mass of copper(II) carbonate needed if a 10% excess is required.
	[Molar mass of copper(II) carbonate = 123.5 g mol–1]

(2)

(ii)	A student doing this experiment chose to use a balance reading to 0.01g in an attempt to work accurately.
	Was this choice of balance necessary from the point of view of accuracy? Explain your answer.
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(1)
 
(d)	The sulfuric acid is heated to boiling and the copper(II) carbonate is added in small portions.
	State the next step needed to prepare pure copper(II) sulfate solution. Justify your answer.
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(1)
 
(e)	When the solution of copper(II) sulfate is allowed to crystallise, the crystals which are produced have the formula CuSO4.5H2O.
(i)	What is the molar mass of CuSO4.5H2O?

(1)

(ii)	3.98 g of CuSO4.5H2O crystals were obtained. Calculate the percentage yield in this experiment.

(2)
(Total 9 marks)

46.	Airbags, used as safety features in cars, contain sodium azide, NaN3. An airbag requires a large volume of gas to be produced in a few milliseconds. The gas is produced in this reaction:
	2NaN3(s) ® 2Na(s) + 3N2(g)	DH is positive
	When the airbag is fully inflated, 50 dm3 of nitrogen gas is produced.
 
(a)	Calculate the number of molecules in 50 dm3 of nitrogen gas under these conditions.
	[The Avogadro constant = 6.02 × 1023 mol–1. The molar volume of nitrogen gas under the conditions in the airbag is 24 dm3 mol–1].

(2)

(b)	Calculate the mass of sodium azide, NaN3, that would produce 50 dm3 of nitrogen gas.

(3)
 
(c)	What will happen to the temperature in the airbag when the reaction occurs?
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(1)
 
(d)	The airbag must be strong enough not to burst in an accident. An airbag which has burst in an accident is hazardous if the sodium azide in it has decomposed.
	Explain why this is so.
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(2)
(Total 8 marks)

47.	A student investigated a reaction which could be used to warm up coffee in self-heating cans.
	Mg(s) + Cu(NO3)2(aq) ® Mg(NO3)2(aq) + Cu(s)
	In the self-heating cans, the bottom has a compartment containing copper(II) nitrate solution. When a button on the bottom of the can is pressed, the magnesium powder is released into the compartment where it reacts with the copper(II) nitrate solution.

(a)	A student investigated the enthalpy change for this reaction by measuring
	50.0 cm3 of 0.300 mol dm–3 copper(II) nitrate solution into a 100 cm3 beaker and adding 1g (an excess) of magnesium powder.
	The results are shown below.
	Temperature of copper(II) nitrate solution at start	= 22 °C
Temperature of mixture after reaction	= 43 °C
(i)	Calculate the energy change which took place. The specific heat capacity of the solution is 4.20 J g–1 K–1.
	Which is the correct value for the energy change in joules?

(1)
 
(ii)	How many moles of copper(II) nitrate were used in the experiment?

(1)
 
(iii)	Calculate the enthalpy change for the reaction. You should include a sign and units in your answer.

(2)

(iv)	Suggest two changes you would make to the equipment used in order to improve the accuracy of the result.
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(2)
 
(b)	The ionic equation for the reaction is shown below:
	Mg(s) + Cu2+(aq) ® Mg2+(aq) + Cu(s)	DH= –532 kJ mol–1
	Would the following affect the value of the experimental result?
	Explain your answer, stating the effect, if any, on the value of the enthalpy change obtained.
(i)	The student used 2 g rather than 1g of magnesium.
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(2)
 
(ii)	The heat losses that occurred from the student’s beaker.
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(2)

(c)	The temperature in the self-heating can needs to increase by 60 °C to produce a hot drink.
	Suggest a change you could make to the mixture in the experiment in (a) to produce a greater temperature rise. You are not expected to do a calculation.
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(1)
(Total 11 marks)