1. The Born-Haber cycle for the formation of sodium chloride from sodium and chlorine may be represented by a series of stages labelled A to F as shown. N a (g ) + C l(g ) + e N a (g ) + C l (g ) + e N a (g ) + C l (g ) N a (s) + C l (g ) + + + – 2 2 2 1 2 1 2 1 2 A B C D F E N a (g ) + C l (g ) N aC l(s) – (a) (i) Write the letters A to F next to the corresponding definition in the table below definition letter H/kJ mol –1 1 st ionisation energy of sodium +494 1 st electron affinity of chlorine –364 the enthalpy of atomisation of sodium +109 the enthalpy of atomisation of chlorine +121 the lattice enthalpy of sodium chloride –770 the enthalpy of formation of sodium chloride (3) (ii) Calculate the enthalpy of formation of sodium chloride from the data given. 1
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1. The Born-Haber cycle for the formation of sodium chloride from sodium and chlorine may be represented by a series of stages labelled A to F as shown.
N a (g ) + C l(g ) + e
N a (g ) + C l (g ) + e
N a(g ) + C l (g )
N a(s ) + C l (g )
+
+
+ –2
2
212
12
12
A
B
C
D
F
E
N a (g ) + C l (g )
N aC l(s )
–
(a) (i) Write the letters A to F next to the corresponding definition in the table below
definition letter H/kJ mol–1
1st ionisation energy of sodium
+494
1st electron affinity of chlorine
–364
the enthalpy of atomisation of sodium
+109
the enthalpy of atomisation of chlorine
+121
the lattice enthalpy of sodium chloride
–770
the enthalpy of formation of sodium chloride
(3)
(ii) Calculate the enthalpy of formation of sodium chloride from the data given.
(2)
1
(b) The lattice enthalpies can be calculated from theory as well as determined experimentally.
Experimental
H/kJ mol–1Theoretical
H/kJ mol–1
Sodium chloride –770 –766
Silver iodide –889 –778
Why is the experimental value of the lattice enthalpy of silver iodide (–889kJmol–1) so different from the value calculated theoretically?
2. (a) The Born-Haber cycle for the formation of sodium chloride is shown below.
N a (g ) + C l(g ) + e
N a (g ) + C l (g ) + e
N a (g ) + C l (g )
N a (s) + C l (g )
N a (g ) + C l (g )
N a C l (s )
–
–
–
–
+
+
+
+
1
1
1
2
2
2
2
2
2
Use the data below to calculate the lattice enthalpy of sodium chloride.
Enthalpy changeValue of the
enthalpy change
/kJ mol–1
Enthalpy of atomisation of sodium +109
1st ionisation energy of sodium +494
Enthalpy of formation of sodium chloride –411
Enthalpy of atomisation of chlorine +121
Electron affinity of chlorine –364
(2)
3
(b) Sodium chloride and magnesium oxide have very similar crystal lattices. Suggest why the lattice enthalpy of magnesium oxide is very much larger than that of sodium chloride.
(c) The lattice enthalpy of silver iodide can be calculated but the experimental value does not match the calculated value as well as those for sodium chloride match each other.
Explain why the calculated and experimental values for silver iodide are different.
9. A student was required to determine the enthalpy change for the reaction between iron and copper sulphate solution.
14
The student produced the following account of their experiment.
A piece of ir on, mass about 3 g, was placed in a glass beaker . 50 cm of0.5 mol dm aqueous copper sulphat e solut ion was measur ed using ameasur ing cylinder and added t o t he beaker . T he t emper at ur e of t hemixt ur e was measur ed immediat ely bef or e t he addit ion and ever y minut eaf t er war ds unt il no f ur t her change t ook place.
F e + CuS O F eS O + Cu4 4
–3
3
T iming bef or eaddit ion
1 min 2 mins 3 mins 4 mins 5 mins
T emper at ur e/ °C 22 27 29 26 24 22
(a) Suggest two improvements you would make to this experiment. Give a reason for each of the improvements suggested.
(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)
16
10. (a) Define the term standard enthalpy of combustion, making clear the meaning of standard in this context.
………….…………………………………………………………………………..
………….…………………………………………………………………………..
………….…………………………………………………………………………..
………….…………………………………………………………………………..
………….…………………………………………………………………………..(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)
Hcombustion/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)
17
(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)
11. (a) This question is about finding the formula of copper hydroxide. The method is as follows:
18
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.
3 0
2 9
2 8
2 7
2 6
2 5
2 4
2 3
2 2
2 1
2 00 4 8 1 2 1 6 20 2 4 2 8 3 2 3 6 4 0
Volu m e o f N a O H (a q ) / cm
Te m p e ra tu re / ºC
– 3
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(i) Explain why the temperature reaches a maximum and then falls slightly on addition of further sodium hydroxide solution.
………….……………….……………………………………………………..
………….……………….……………………………………………………..
………….……………….……………………………………………………..(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)
20
(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 × T joules
(1)
(iii) Use your results from (a)(iv) and (b)(ii) above, to find the molar enthalpy change, H, 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.
………………...……………….……………………………………………………..
………………...……………….……………………………………………………..
………………...……………….……………………………………………………..
………………...……………….……………………………………………………..
………………...……………….……………………………………………………..(2)
(Total 14 marks)
21
12. 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 = specificheat capacity
mass ofsolution
temperaturechange
/J /J g1 K1 /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)
13. 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)
22
(a) Define the term standard enthalpy of formation, Hf , of urea.
(c) Iodoethane reacts with water to form ethanol and hydrogen iodide.
C2H5I + H2O C2H5OH + HI Hf = +36 kJ mol–1
Use some or all of the data below to calculate the CI bond enthalpy.
Bond Bond enthalpy
/ kJ mol–1Bond Bond enthalpy
/ kJ mol–1
CH 413 HI 298
CC 347 CO 358
HO 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)
25
15. (a) The equation below shows the reaction which occurs when ammonia is dissolved in water.
NH3(g) + H2O(1) NH4 (aq) + OH–(aq)
(i) Explain why water is classified as an acid in this reaction.
(b) Calculate the heat gained by the water. Give your answer in kJ.
(2)
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(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 wasnumerically much less than the data book value. Thereasons 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 lostmass 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.
17. 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)
30
(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)
31
(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.
18. The reaction between chlorine and methane, in the presence of ultraviolet light, involves the formation of free radicals and includes the following steps:
A Cl2 2Cl• ΔΗο = +242 kJ mol–1
B CH4 + Cl• HCl + CH3• ΔΗο = +4 kJ mol–1
C Cl2 + CH3• CH3Cl + Cl• ΔΗο = –97 kJ mol–1
D Cl• + Cl• Cl2
E CH3• + CH3
• CH3CH3
F Cl• + CH3• CH3Cl ΔΗο = –339 kJ mol–1
32
(a) (i) What is meant by a free radical? ....................................................................
19. 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, H1, for the reaction
NaHCO3(s) + HCl(aq) NaCl(aq) + CO2(g) + H2O(l)
Measurement Reading
Mass of solid sodium hydrogencarbonate added tohydrochloric acid.
5.00 g
Volume of hydrochloric acid 50.0 cm3
Temperature of hydrochloric acid before additionof 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)
35
(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 H1 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)
36
(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, H2, 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.
20. 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.
(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) waterexcess 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
39
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)
21. 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)
40
(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).
(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.
ΔHο
f[PH3(g)] = + 5.4 kJ mol1
ΔHοat[½H2(g)] = + 218.0 kJ mol1
ΔHο
at[P(s)] = + 314.6 kJ mol1
(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)
22. 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.
(b) Methane burns in oxygen according to the equation:
H C H (g ) + 2 O O (g ) O C O (g ) + 2 H O H (g )
H
H
Use the average bond enthalpy data shown below to calculate the enthalpy change of this reaction.
Bond Bond enthalpy/kJ mol–1
CH +435
O==O +498
C==O +805
HO +464
(3)
43
(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)
23. 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.
44
(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 ...................................................................................................................................
(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.
Tem p era tu re/°C
2 5
2 0
1 50 1 0 2 0 3 0 4 0 5 0
Vo lu m e o f so d iu m h y d ro x id e ad d ed / cm 3
×
×
×
×
××
×× × ×
×
(2)
45
(c) From the graph, read off the maximum temperature rise, T, and the volume of aqueous sodium hydroxide added at neutralisation, VN.
T = ..................................... C VN = ................................ cm3
(2)
(d) (i) Use the formula below to calculate the heat evolved in the neutralisation.
Heat evolved =
1000
18.4T25VN
kJ
(1)
(ii) Given that the amount (moles) of acid neutralised was 0.025 mol, calculate the
enthalpy of neutralisation, Hneut, in units of kJ mol–1.
Hneut = ............................... kJ mol–1
(2)(Total 9 marks)
46
24. A reaction of ammonium dichromate(VI) is shown by the following equation.
(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)
47
(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?
25. 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)
49
(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)
26. (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)
50
(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).
28. The values of the lattice energies of potassium iodide and calcium iodide experimentally determined from Born-Haber cycles and theoretically calculated from an ionic model are shown below.
Experimental latticeenergy
/kJ mol–1
Theoreticallattice energy
/kJ mol–1
Potassium iodide, KI(s) – 651 – 636
Calcium iodide, CaI2(s) –2074 –1905
(i) Explain why the experimental lattice energy of potassium iodide is less exothermic than the experimental lattice energy of calcium iodide.
(ii) Explain why the experimental and theoretical values of the lattice energy are almost the same for potassium iodide, but are significantly different for calcium iodide.
(iii) Hence calculate the concentration of nitric acid, HNO3, in mol dm3.
(2)
(d) (i) Use the data from (b) to calculate the heat change for this reaction.
63
The density of the mixture produced at neutralisation is 1.0g cm–3 and the specific
heat capacity of the mixture is 4.2 J g–1 °C–1.
Heat change = mass × specific heat capacity × T
(2)
(ii) Use your answer from (d)(i) and (c)(iii) to calculate the enthalpy of neutralisation per mole of nitric acid, HNO3. Include a sign and units with your answer.
(3)
(e) The enthalpy of neutralisation found by this method may be less exothermic than the data book value because of heat loss.
Suggest ONE way to reduce the error due to heat loss.
32. (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)
65
(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, Hο, is –557 kJ mol–1. State TWO of the reaction conditions necessary for this enthalpy change to be standard.
34. 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.
C aC O (s ) C aO (s) + C O (g )H re a c tio n
3 2
H 3 H 4
E lem en ts in th e ir s tan d a rd s ta te s
(a) Suggest ONE reason why it is difficult to measure Hreaction directly by experiment.
(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.
(b) In the Haber process, ammonia is manufactured from nitrogen and hydrogen as shown in the equation.
N2(g) + 3H2(g) 2NH3(g)
(i) Use the bond enthalpies below to calculate the standard enthalpy of formation of ammonia.
Bond Bond enthalpy / kJ mol–1
N≡N in N2 +945
H–H in H2 +436
N–H in NH3 +391
(4)
(ii) Draw a labelled enthalpy level diagram for the formation of ammonia in the Haber process.
E n th a lp y
(2)
75
(iii) State the temperature used in the Haber process and explain in terms of the rate of reaction and position of equilibrium, why this temperature is chosen.
(c) The table shows values for the lattice energies of the metal chlorides of some Group 2 metals.
Group 2 metalchloride
MgCl2 CaCl2 SrCl2 BaCl2
Lattice energy/
kJ mol–1 –2526 –2237 –2112 –2018
Explain why these lattice energies become less exothermic from MgCl2 to BaCl2.
(3)(Total 10 marks)
80
39. 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.
(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
82
(i) On the graph paper below, plot the results of this experiment.
(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.
(ii) Use your Hess’s Law cycle to calculate the standard enthalpy of formation of ethanol.
(2)(Total 8 marks)
42. 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
th e rm o m ete r
b eak e r
w a te r
m e th an o l
sp irit lam p
Results
Molar mass (methanol) = 32 g mol–1
Volume of water in beaker = 50 cm3
Mass of water in beaker = 50 g
WeighingsSpirit lamp + methanol before combustion = 163.78 gSpirit lamp + methanol after combustion = 163.44 g
TemperaturesWater before heating = 22.0 °CWater after heating = 43.5 °C
Specific heat capacity of water = 4.18 J g–1 °C–1
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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.
(iii) Use the Born-Haber cycle to calculate the lattice energy of magnesium oxide.
(2)
(b) Magnesium iodide is another compound of magnesium. The radius of the magnesium ion is 0.072 nm, whereas the radius of the iodide ion is much larger and is 0.215 nm.
(i) Describe the effect that the magnesium ion has on an iodide ion next to it in the magnesium iodide lattice.
44. 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.
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(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)
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(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.
C a (O H ) 2 (s ) C aO (s) + H 2O (l)
C a C l2 (aq ) + 2 H 2O (l)
H reac tio n
H 1 H 2
2 H C l (a q ) 2 H C l (aq )
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(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.
(b) (i) Use the average (mean) bond enthalpy data to calculate a value for the enthalpy change for this reaction. You are reminded to show all your working.
BondAverage bond enthalpy
/ kJ mol–1
N≡N 944
H—H 436
N—H 388
(3)
(ii) The actual standard enthalpy change for this reaction is –92 kJ mol–1. Explain why the value you calculated in (b)(i) is not the same as this.
(c) The manufacturer of ammonia would like to achieve a high rate of reaction and a high equilibrium yield of product.
(i) State and explain, in terms of collision theory, TWO ways to increase the rate of the reaction. An increase in pressure does not alter the rate in this process.
46. 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.
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(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).
48. 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.
H 3
E x p er im en t 1
M g + C + O 2 M g C O (s)3
M g C l2 (aq ) + H 2 (g ) + C + O 2
H f
H 1
H 2– 1= – 6 8 0 k J m o l
+ 2 H C l (aq )
+ 2 H C l (aq ) E x p er im en t 2
M g C l2 (aq ) + H 2O (l) + C O 2 (g )
(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?
(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)
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(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?
51. (a) The following data were collected to use in a Born-Haber cycle for silver fluoride, AgF.
Value
/kJ mol–1
enthalpy of atomisation of silver +285
first ionisation energy of silver +731
enthalpy of atomisation of fluorine +79
enthalpy of formation of silver fluoride –205
lattice energy of silver fluoride –958
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On the following outline of a Born-Haber cycle, complete boxes A and B by adding the formula and state symbol for the appropriate species. Write the name of the enthalpy change at C.
A g (g )+
A g (s) + ½ F 2 (g )
F (g )– A g F (s)+
B o x B
B o x A
C ............................................
............................................
(3)
(b) ΔHlatt (theoretical) is the lattice energy calculated assuming the crystal lattice is completely ionic.ΔHlatt (experimental) is the lattice energy determined experimentally using the Born-Haber cycle.
Values for the silver halides are listed below.
Formula of halide ΔHlatt
(theoretical)
/ kJ mol–1
ΔHlatt
(experimental)
/ kJ mol–1
ΔHlatt (theoretical)minus
ΔHlatt (experimental)
/ kJ mol–1
AgF –920 –958 38
AgCl –833 –905 72
AgBr –816 –891 75
AgI –778 –889 111
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(i) Explain why the theoretical lattice energies become less exothermic from AgF to AgI.
52. 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.
S O 2 (g ) + 2 H 2 S (g ) 3 S (s) + 2 H 2O (l)
H r
H 2H 1
(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?
(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)
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53. Which of the equations shown below represents the reaction for which H is the standard
enthalpy change of formation, Hοf 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)
54. 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) H 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)
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(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?
(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.
55. 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.
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(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 °CTemperature 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)
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(iv) Suggest two changes you would make to the equipment used in order to improve the accuracy of the result.
(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.
56. The following data can be used in a Born-Haber cycle for copper(II) bromide, CuBr2.
Enthalpy change of atomisation of bromine Hοat[½Br2(l)] +111.9 kJ mol–1
Enthalpy change of atomisation of copper, Hοat[Cu(s)] +338.3 kJ mol–1
First ionisation energy of copper, Em1[Cu(g)] +746.0 kJ mol–1
Second ionisation energy of copper, Em2 [Cu(g)] +1958.0 kJ mol–1
Electron affinity of bromine, Eaff[Br(g)] –342.6 kJ mol–1
Enthalpy change of formation of CuBr2(s), Hοf [CuBr2(s)] –141.8 kJ mol–1
(a) On the following outline of a Born-Haber cycle complete the boxes A, B, and C by putting in the formula and state symbol for the appropriate species and writing the name of the enthalpy change D.
C u (g ) C u B r2 +2(s)2 B r (g )–
C u (s) + B r 2(l)
B
A
C D ...........................................
+
(3)
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(b) Use the data to calculate a value for the lattice energy of copper(II) bromide.
Give a sign and units in your answer.
(3)
(c) When the lattice energy of copper(II) bromide is calculated from ionic radii and charges, the result is a value numerically about 10% less than the one obtained from the Born-Haber cycle.
(i) What does this suggest about the nature of the bonding in copper(II) bromide?