chemistry sl c. ii and iii only topic 5 questions - ib relics · chemistry sl topic 5 questions 1....

Chemistry SL Topic 5 Questions

1. What energy changes occur when chemical bonds are formed and broken?

A. Energy is absorbed when bonds are formed and when they are broken.

B. Energy is released when bonds are formed and when they are broken.

C. Energy is absorbed when bonds are formed and released when they are broken.

D. Energy is released when bonds are formed and absorbed when they are broken.

2. The temperature of a 2.0 g sample of aluminium increases from 25°C to 30°C. How many joules of heat energy were added? (Specific heat of Al = 0.90 J g–1K–1)

A. 0.36

B. 2.3

C. 9.0

D. 11

3. Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –390 kJ Mn(s) + O2(g) → MnO2(s) ∆H = –520 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g)

A. 910

B. 130

C. –130

D. –910

4. Which statements about exothermic reactions are correct?

I. They have negative ⊗H values.

II. The products have a lower enthalpy than the reactants.

III. The products are more energetically stable than the reactants.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

5. A sample of a metal is heated. Which of the following are needed to calculate the heat absorbed by the sample?

I. The mass of the sample

II. The density of the sample

III. The specific heat capacity of the sample

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

6. The average bond enthalpies for O—O and O==O are 146 and 496 kJ mol–1 respectively. What is the enthalpy change, in kJ, for the reaction below?

H—O—O—H(g) → H—O—H(g) + ½O==O(g)

A. – 102 B. + 102

C. + 350 D. + 394

7. When the solids Ba(OH)2 and NH4SCN are mixed, a solution is produced and the temperature drops.

Ba(OH)2(s) + 2NH4SCN(s) → Ba(SCN)2(aq) + 2NH3(g) + 2H2O(l)

Which statement about the energetics of this reaction is correct?

A. The reaction is endothermic and ⊗H is negative.

B. The reaction is endothermic and ⊗H is positive.

C. The reaction is exothermic and ⊗H is negative.

D. The reaction is exothermic and ⊗H is positive.

8. Using the equations below

Cu(s) + O2(g) → CuO(s)∆Hο = –156 kJ

2Cu(s) + O2(g) → Cu2O(s)∆Hο = –170 kJ

what is the value of ∆Hο (in kJ) for the following reaction?

2CuO(s) → Cu2O(s) + O2(g)

A. 142 B. 15

C. –15 D. –142

9. Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a catalyst?

Time

A. I only

B. III only

C. I and II only

D. II and III only

10. Consider the following equations.

Mg(s) + O2(g) → MgO(s) ∆Hο = –602 kJ H2(g) + O2(g) → H2O(g) ∆Hο = –242 kJ

What is the ∆H° value (in kJ) for the following reaction?

MgO(s) + H2(g) → Mg(s) + H2O(g)

A. –844 B. –360

C. +360 D. +844

11. For which of the following is the sign of the enthalpy change different from the other three?

A. CaCO3(s) → CaO(s) + CO2(g)

B. Na(g) → Na+(g) + e–

C. CO2(s) → CO2(g)

D. 2Cl(g) → Cl2(g)

12. Separate solutions of HCl(aq) and H2SO4(aq) of the same concentration and same volume were completely neutralized by NaOH(aq). X kJ and Y kJ of heat were evolved respectively. Which statement is correct?

A. X = Y B. Y = 2X

C. X = 2Y D. Y = 3X

13. Which statements are correct for an endothermic reaction?

I. The system absorbs heat.

II. The enthalpy change is positive.

III. The bond enthalpy total for the reactants is greater than for the products.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

14. The mass m (in g) of a substance of specific heat capacity c (in J g–1 K–1 ) increases by t°C. What is the heat change in J?

A. mct

B. mc(t + 273)

C.

D.

15. The average bond enthalpy for the C―H bond is 412 kJ mol–1. Which process has an enthalpy change closest to this value?

A. CH4(g) → C(s) + 2H2(g)

B. CH4(g) → C(g) + 2H2(g)

C. CH4(g) → C(s) + 4H(g)

D. CH4(g) → CH3(g) + H(g) (

16. The following equation shows the formation of magnesium oxide from magnesium metal.

2Mg(s) + O2(g)→2MgO(s) ⊗HӨ = –1204kJ

Which statement is correct for this reaction?

A. 1204 kJ of energy are released for every mol of magnesium reacted.

B. 602 kJ of energy are absorbed for every mol of magnesium oxide formed.

C. 602 kJ of energy are released for every mol of oxygen gas reacted.

D. 1204 kJ of energy are released for every two mol of magnesium oxide formed.

17. The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.

C(s) +O2(g) → CO2(g) ⊗HӨ = –x kJ mol–1

CO(g) + O2(g)→ CO2(g) ⊗HӨ = –y kJ mol–l

What is the enthalpy change, in kJ mol–1, for the oxidation of carbon to carbon monoxide?

C(s) + O2(g)→ CO(g)

A. x + y B. – x – y

C. y – x D. x – y

18. A simple calorimeter was used to determine the enthalpy of combustion of ethanol. The experimental value obtained was –920 kJ mol–1. The Data Booklet value is –1371 kJ mol–1. Which of the following best explains the difference between the two values?

A. incomplete combustion of the fuel

B. heat loss to the surroundings

C. poor ventilation in the laboratory

D. inaccurate temperature measurements

19. For the reaction

2H2(g) + O2(g) → 2H2O(g)

the bond enthalpies (in kJ mol–1) are H–H x

O=O y

O–H z

Which calculation will give the value, in kJ mol–1, of ⊗HӨ for the reaction?

A. 2x + y –2z

B. 4z – 2x – y

C. 2x + y – 4z

D. 2z –2x – y

20. Which statement about bond enthalpies is correct?

A. Bond enthalpies have positive values for strong bonds and negative values for weak bonds.

B. Bond enthalpy values are greater for ionic bonds than for covalent bonds.

C. Bond breaking is endothermic and bond making is exothermic.

D. The carbon–carbon bond enthalpy values are the same in ethane and ethene.

21. An equation for a reaction in which hydrogen is formed is

CH4 + H2O → 3H2 + CO ⊗HӨ = +210 kJ

Which energy change occurs when 1 mol of hydrogen is formed in this reaction?

A. 70 kJ of energy are absorbed from the surroundings.

B. 70 kJ of energy are released to the surroundings.

C. 210 kJ of energy are absorbed from the surroundings.

D. 210 kJ of energy are released to the surroundings.

22. The equations and enthalpy changes for two reactions used in the manufacture of sulfuric acid are:

S(s) O2(g) → SO2(g) ⊗HӨ = –300 kJ

2SO2(g) + O2(g) → 2SO3(g) ⊗HӨ = –200 kJ

What is the enthalpy change, in kJ, for the reaction below?

2S(s) + 3O2(g) → 2SO3(g)

A. –100

B. –400

C. –500

D. –800

23. Approximate values of the average bond enthalpies, in kJ mol–1, of three substances are:

H–H 430

F–F 155

H–F 565

What is the enthalpy change, in kJ, for this reaction?

2HF → H2 + F2

A. +545

B. +20

C. –20

D. –545

24. The standard enthalpy change of formation values of two oxides of phosphorus are:

P4(s) + 3O2(g) → P4O6(s) ⊗HӨf= –1600 kJ mol–1

P4(s) + 5O2(g) → P4O10(s) ⊗HӨf= –3000 kJ mol–1

What is the enthalpy change, in kJ mol–1, for the reaction below?

P4O6(s) + 2O2(g) → P4O10(s)

A. +4600

B. +1400

C. –1400

D. –4600

25. Which statement is correct for an endothermic reaction?

A. The products are more stable than the reactants and ⊗H is positive.

B. The products are less stable than the reactants and ⊗H is negative.

C. The reactants are more stable than the products and ⊗H is positive.

D. The reactants are less stable than the products and ⊗H is negative.

26. Which statement is correct about the reaction shown?

2SO2(g) + O2(g) → 2SO3(g) ⊗H = –196 kJ

A. 196 kJ of energy are released for every mole of SO2(g) reacted.

B. 196 kJ of energy are absorbed for every mole of SO2(g) reacted.

C. 98 kJ of energy are released for every mole of SO2(g) reacted.

D. 98 kJ of energy are absorbed for every mole of SO2(g) reacted.

27. Which statements are correct for all exothermic reactions?

I. The enthalpy of the products is less than the enthalpy of the reactants.

II. The sign of ⊗H is negative.

III. The reaction is rapid at room temperature.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

28. Consider the specific heat capacity of the following metals.

Metal Specific heat capacity / J kg–1 K–1

Cu 385

Ag 234

Au 130

Pt 134

Which metal will show the greatest temperature increase if 50 J of heat is supplied to a 0.001 kg sample of each metal at the same initial temperature?

A. Cu

B. Ag

C. Au

D. Pt

29. Consider the following reactions.

S(s) + O2(g) → SO3(g) ⊗HӨ = −395 kJ mol−1

SO2(s) + O2(g) → SO3(g) ⊗HӨ = −98 kJ mol−1

What is the ⊗HӨ value (in kJ mol–1) for the following reaction?

S(s) + O2(g) → SO2(g)

A. –297

B. +297

C. – 493

D. +493

30. Which statement is correct for an endothermic reaction?

A. Bonds in the products are stronger than the bonds in the reactants.

B. Bonds in the reactants are stronger than the bonds in the products.

C. The enthalpy of the products is less than that of the reactants.

D. The reaction is spontaneous at low temperatures but becomes non-spontaneous at high temperatures.

31. According to the enthalpy level diagram below, what is the sign for ⊗H and what term is used to refer to the reaction?

⊗H reaction

A. positive endothermic

B. negative exothermic

C. positive exothermic

D. negative endothermic

32. When 40 joules of heat are added to a sample of solid H2O at –16.0°C the temperature increases to –8.0°C. What is the mass of the solid H2O sample?

[Specific heat capacity of H2O(s) = 2.0 J g–1K–1]

A. 2.5 g

B. 5.0 g

C. 10 g

D. 160 g

33. The ⊗HӨ values for the formation of two oxides of nitrogen are given below.

N2(g) + O2(g) → NO2(g) ⊗HӨ = –57 kJ mol–1

N2(g) + 2O2(g) → N2O4(g) ⊗HӨ = +9 kJ mol–1

Use these values to calculate ⊗HӨ for the following reaction (in kJ):

2NO2(g) → N2O4(g)

A. –105

B. – 48

C. +66

D. +123

34. How much energy, in joules, is required to increase the temperature of 2.0 g of aluminium from 25 to 30°C? (Specific heat of Al = 0.90 J g–1 K–1).

A. 0.36

B. 4.5

C. 9.0

D. 54

35. Which combination is correct for a chemical reaction that absorbs heat from the surroundings?

Type of reaction ΔH at constant pressure

A. Exothermic Positive

B. Exothermic Negative

C. Endothermic Positive

D. Endothermic Negative

36. Using the equations below:

C(s) + O2(g) → CO2(g) ∆Hο = –394 kJ mol–1

Mn(s) + O2(g) → MnO2(s) ∆Hο = –520 kJ mol–1

What is ∆H, in kJ, for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g)

A. 914

B. 126

C. –126

D. –914

37.

The diagram shows the distribution of energy for the molecules in a sample of gas at a given temperature, T1.

(a) In the diagram Ea represents the activation energy for a reaction. Define this term.

……………………………………………………………………………………….

………………………………………………………………………………………. (1)

(b) On the diagram above draw another curve to show the energy distribution for the same gas at a higher temperature. Label the curve T2.

(2)

(c) With reference to your diagram, state and explain what happens to the rate of a reaction when the temperature is increased.

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(Total 5 marks)

38. (a) Define the term average bond enthalpy, illustrating your answer with an equation for methane, CH4.

..................................................................................................................................... (3)

(b) The equation for the reaction between methane and chlorine is

CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)

Use the values from Table 10 of the Data Booklet to calculate the enthalpy change for this reaction.

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(c) Explain why no reaction takes place between methane and chlorine at room temperature unless the reactants are sparked, exposed to UV light or heated.

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(d) Draw an enthalpy level diagram for this reaction.

(2)

(Total 10 marks)

39. In aqueous solution, potassium hydroxide and hydrochloric acid react as follows.

KOH(aq) + HCl(aq) → KCl(aq)+ H2O(l)

The data below is from an experiment to determine the enthalpy change of this reaction.

50.0 cm3 of a 0.500 mol dm–3 solution of KOH was mixed rapidly in a glass beaker with 50.0 cm3 of a 0.500 mol dm–3 solution of HCl.

Initial temperature of each solution = 19.6°C Final temperature of the mixture = 23.1°C

(a) State, with a reason, whether the reaction is exothermic or endothermic.

..................................................................................................................................... (1)

(b) Explain why the solutions were mixed rapidly.

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(c) Calculate the enthalpy change of this reaction in kJ mol–1. Assume that the specific heat capacity of the solution is the same as that of water.

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(d) Identify the major source of error in the experimental procedure described above. Explain how it could be minimized.

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(e) The experiment was repeated but with an HCl concentration of 0.510 mol dm–3 instead of 0.500 mol dm–3. State and explain what the temperature change would be.

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..................................................................................................................................... (2)

(Total 10 marks)

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

Mass of sample of sucrose, m = 0.4385 g

Heat capacity of the system, Csystem = 10.114 kJ K–1

(a) Calculate ΔT, for the water, surrounding the chamber in the calorimeter.

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(b) Determine the amount, in moles, of sucrose.

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(c) (i) Calculate the enthalpy change for the combustion of 1 mole of sucrose.

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(ii) Using Table 12 of the Data Booklet, calculate the percentage experimental error based on the data used in this experiment.

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(d) A hypothesis is suggested that TNT, 2-methyl-1,3,5-trinitrobenzene, is a powerful explosive because it has:

• a large enthalpy of combustion • a high reaction rate • a large volume of gas generated upon combustion

Use your answer in part (c)(i) and the following data to evaluate this hypothesis:

Equation for combustion Relative rate

of combustion

Enthalpy of combustion / kJ mol–

1

Sucrose C12H22O11(s) + 12O2(g) → 12CO2(g) + 11H2O(g) Low

TNT 2C7H5N3O6(s) → 7CO(g) + 7C(s) + 5H2O(g) + 3N2(g) High 3406

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(Total 7 marks)

41. (a) Define the term average bond enthalpy.

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(b) Use the information from Table 10 of the Data Booklet to calculate the enthalpy change for the complete combustion of but-1-ene, according to the following equation.

C4H8(g) + 6O2(g) → 4CO2(g) + 4H2O(g)

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..................................................................................................................................... (3)

(Total 5 marks)

42. Given the following data:

C(s) + F2(g) → CF4(g); ∆H1 = –680 kJ mol–1 F2(g) → 2F(g); ∆H2 = +158 kJ mol–1 C(s) → C(g); ∆H3 = +715 kJ mol–1

calculate the average bond enthalpy (in kJ mol–1) for the C––F bond.

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43. Two reactions occurring in the manufacture of sulfuric acid are shown below:

reaction I S(s) +O2(g) → SO2(g) ⊗HӨ = –297 kJ

reaction II SO2(g) + O2(g) SO3(g) ⊗HӨ = –92 kJ

(i) State the name of the term ⊗HӨ. State, with a reason, whether reaction I would be accompanied

by a decrease or increase in temperature. (3)

(ii) At room temperature sulfur trioxide, SO3, is a solid. Deduce, with a reason, whether the ⊗HӨ value would be more negative or less negative if SO3(s) instead of SO3(g) were formed in reaction II.

(2)

(iii) Deduce the ⊗HӨ value of this reaction:

S(s) + O2(g) → SO3(g) (1)

(Total 6 marks)

44. (i) Define the term average bond enthalpy. (3)

(ii) Explain why Br2 is not suitable as an example to illustrate the term average bond enthalpy. (1)

(iii) Using values from Table 10 of the Data Booklet, calculate the enthalpy change for the following reaction:

CH4(g) + Br2(g) → CH3Br(g) + HBr(g) (3)

(iv) Sketch an enthalpy level diagram for the reaction in part (iii). (2)

(v) Without carrying out a calculation, suggest, with a reason, how the enthalpy change for the following reaction compares with that of the reaction in part (iii):

CH3Br(g) + Br2(g) → CH2Br2(g) + HBr(g) (2)

(Total 11 marks)

45. But–1–ene gas, burns in oxygen to produce carbon dioxide and water vapour according to the following equation.

C4H8 + 6O2 → 4CO2 + 4H2O

(a) Use the data below to calculate the value of ⊗HӨ for the combustion of but-1-ene.

Bond C−C C=C C−H O=O C=O O–H

Average bond enthalpy / kJ mol–1

348 612 412 496 743 463

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(b) State and explain whether the reaction above is endothermic or exothermic.

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(Total 4 marks)

46. Calculate the enthalpy change, ⊗H4 for the reaction

C + 2H2 + O2 → CH3OH ⊗H4

using Hess’s Law and the following information.

CH3OH + O2 → CO2 + 2H2O ⊗H1 = −676 kJ mol−1 C + O2 → CO2 ⊗H2 = −394 kJ mol−1 H2 + O2 → H2O ⊗H3 = −242 kJ mol−1

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47. Methylamine can be manufactured by the following reaction.

CH3OH(g) + NH3(g) → CH3NH2(g) + H2O(g)

(a) Define the term average bond enthalpy.

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(b) Use information from Table 10 of the Data Booklet to calculate the enthalpy change for this reaction.

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(Total 6 marks)

48. (a) Define the term average bond enthalpy.

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(b) Use the information from Table 10 in the Data Booklet to calculate the enthalpy change for the complete combustion of but-1-ene according to the following equation

C4H8(g) → 4CO2(g) + 4H2O(g)

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(c) Predict, giving a reason, how the enthalpy change for the complete combustion of but-2-ene would compare with that of but-1-ene based on average bond enthalpies.

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(d) The enthalpy level diagram for a certain reaction is shown below.

State and explain the relative stabilities of the reactants and products.

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(Total 8 marks)

49. The reaction between ethene and hydrogen gas is exothermic.

(i) Write an equation for this reaction. (1)

(ii) Deduce the relative stabilities and energies of the reactants and products. (2)

(iii) Explain, by referring to the bonds in the molecules, why the reaction is exothermic. (2)

(Total 5 marks) 50. (i) Define the term average bond enthalpy.

(2)

(ii) The equation for the reaction of ethyne and hydrogen is:

C2H2(g) + 2H2(g) → C2H6(g)

Use information from Table 10 of the Data Booklet to calculate the change in enthalpy for the reaction.

(2)

(iii) State and explain the trend in the bond enthalpies of the C–Cl, C–Br and C–I bonds. (2)

(Total 6 marks) IB Chemistry SL Topic 5 Answers

1. D

2. C

3. B

4. D

5. B

6. A

7. B"

8. A"

9. C

10. C

11. D

12. B 13. A"

14. A"

15. D"

16. D"

17. C"

18. B 19. C"

20. C

21. A

22. D

23. A

24. C

25. C

26. C

27. A

28. C

29. A

30. B

31. B

32. A

33. D

34. C

35. C

36. B

37. (a) activation energy = minimum energy required for a reaction to occur; 1

(b) curve moved to the right; peak lower, 2 Deduct [1] if shaded area smaller at T2 or if T2 line touches the x-axis

(c) rate increased; as more molecules with energy ≥ Ea; 2 [5]

38. (a) energy for the conversion of a gaseous molecule into (gaseous) atoms; (average values) obtained from a number of similar bonds/compounds/OWTTE; CH4(g) → C(g) + 4H(g); 3

State symbols needed.

(b) (bond breaking) = 1890/654; (bond formation) = 2005/769; enthalpy = –115(kJ mol–1) 3 Award [3] for correct final answer. Penalize [1] for correct answer with wrong sign.

(c) molecules have insufficient energy to react (at room temperature)/ wrong collision geometry/unsuccessful collisions; extra energy needed to overcome the activation energy/Ea for the reaction; 2

(d)

exothermic shown;

activation energy/Ea shown; 2 [10]

39. (a) exothermic because temperature rises/heat is released; 1

(b) to make any heat loss as small as possible/so that all the heat will be given out very quickly; 1

Do not accept “to produce a faster reaction”.

(c) heat released = mass×specific heat capacity×temp increase/q = mc∆T =/ 100×4.18×3.5; = 1463 J/1.463 kJ; (allow 1.47 kJ if specific heat = 4.2) amount of KOH/HCl used = 0.500×0.050 = 0.025 mol; ∆H = (1.463÷0.025) = –58.5 (kJ mol–1); (minus sign needed for mark) 4

Use ECF for values of q and amount used. Award [4] for correct final answer. Final answer of 58.5 or +58.5 scores [3]. Accept 2,3 or 4 significant figures.

(d) heat loss (to the surroundings); insulate the reaction vessel/use a lid/draw a temperature versus time graph; 2

(e) 3.5°C/temperature change would be the same; amount of base reacted would be the same/excess acid would not react/ KOH is the limiting reagent; 2

[10]

40. (a) ΔT = 23.70 – 23.03 = 0.67 (°C/K); 1

(b) = 1.281×10–3; 1

(c) (i) ΔHc = (C ΔT)/n = = –5.3×103 kJ mol–1; 1 Use ECF for values of ⊗T and n.

(ii) Percentage experimental error == 5.4%; 1 Use ECF for values of ΔHc.

(d) enthalpy change of combustion of sucrose > TNT, and therefore not important; rate of reaction for TNT is greater than that of sucrose, so this is valid; amount of gas generated (in mol) for sucrose > than that of TNT (according to the given equation), so this is not important; 3

[7]

41. (a) The amount of energy needed to break 1 mole of (covalent) bonds; in the gaseous state; average calculated from a range of compounds; 2 max

Award [1] each for any two points above.

(b) Bonds broken (612) + (2×348) + (8×412) + (6×496)/7580 (kJ mol–1); Bonds made (8×743) + (8×463) / 9648 (kJ mol–1); ⊗H = –2068 (kJ mol–1); 3

Award [3] for the correct answer. Allow full ECF. Allow kJ but no other incorrect units. Even if the first two marks are lost, the candidate can score [1] for a clear correct subtraction for ⊗H.

[5]

42. C(s) + 2F2(g) → CF4(g) ∆H1 = –680 kJ; 4F(g) → 2F2(g) ∆H2 = 2(–158) kJ; C(g) → C(s) ∆H3 = –715 kJ;

Accept reverse equations with +∆H values.

C(g) + 4F(g) → CF4(g) ∆H = –1711 kJ,

so average bond enthalpy =

= –428 kJ mol–1; 4 Accept + or – sign. Lots of ways to do this! The correct answer is very different from the value in the Data Booklet, so award [4] for final answer with/without sign units not needed, but deduct [1] if incorrect units. Accept answer in range of 427 to 428 without penalty for sig figs. If final answer is not correct use following; Award [1] for evidence of cycle or enthalpy diagram or adding of equations. Award [1] for 2F2 (g) → 4F(g) 2×158 seen. Award [1] for dividing 1711 or other value by 4.

[4]

43. (a) (i) standard enthalpy (change) of reaction; (temperature) increase; reaction is exothermic/sign of ⊗H° is negative; 3

(ii) more (negative); heat given out when gas changes to solid/solid has less enthalpy than gas/OWTTE; 2

(iii) –389 kJ; 1 [6]

44. (i) the energy needed to break one bond; (in a molecule in the) gaseous state; value averaged using those from similar compounds; 3

(ii) it is an element/no other species with just a Br-Br bond/OWTTE; 1

(iii) (sum bonds broken =) 412 + 193 = 605; (sum bonds formed =) 276 + 366 = 642; (⊗H% =) –37 kJ; 3

Award [3] for correct final answer. Award [2] for “+ 37”. Accept answer based on breaking and making extra C-H bonds.

(iv)

Enthalpy CH4 + Br2

CH3Br + HBr ;

2 Award [1] for enthalpy label and two horizontal lines, [1] for reactants higher than products. ECF from sign in (iii), ignore any higher energy level involving atoms.

(v) (about) the same/similar; same (number and type of) bonds being broken and formed; 2

[11]

45. (a) (Amount of energy required to break bonds of reactants) 8×412 + 2×348 + 612 + 6×496/7580 (kJ mol−1);

(Amount of energy released during bond formation) 4×2×743 + 4×2×463/9648 (kJ mol−1);

⊗H = −2068 (kJ or kJ mol−1); 3 ECF from above answers. Correct answer scores [3]. Award [2] for (+)2068. If any other units apply −1(U), but only once per paper.

(b) exothermic and ⊗HӨ is negative/energy is released; 1 Apply ECF to sign of answer in part (a). Do not mark if no answer to (a).

[4]

46. −1×⊗H1/676; 1×⊗H2/–394; 2×⊗H3/– 484; ⊗H4 = −202 (kJ mol−1 ); 4

Accept alternative methods. Correct answers score [4].

Award [3] for (+)202 or (+)40 (kJ/kJ mol−1). −1(U) if units incorrect (ignore if absent).

[4]

47. (a) energy needed to break (1 mol of) a bond in a gaseous molecule; averaged over similar compounds; 2

(b) bonds broken identified as C−O and N−H; bonds formed identified as C−N and O−H; ⊗H = 748 − 768 (kJ);

= − 20 kJ/kJ mol−1 (units needed for this mark); 4 If wrong bonds identified apply ECF to 3rd and 4th marks. Accept answer based on breaking and making all bonds. Award [4] for correct final answer. Award max [3] if only one bond missed.

Answer of 20 or +20 kJ (mol−1 ) scores [3]. [6]

48. (a) amount of energy needed to break one mole of (covalent) bonds; in the gaseous state; average calculated from a range of compounds; 2

Award [1] each for any two points above.

(b) bonds broken: 161 + 2×348 + 8×412 + 6×496/7580 kJ mol−1;

bonds made: 8×743 + 8×463/9648 kJ mol−1;

(bonds broken − bonds made =) ⊗H = −2068(kJ mol−1); 3 Award [3] for the correct answer. Allow full ECF − 1 mistake equals 1 penalty. Allow kJ but not other wrong units.

(c) same/equal, because the same bonds are being broken and formed; 1

(d) products more stable than reactants; bonds are stronger in products than reactants/HP < HR/enthalpy/stored energy of products less

than reactants; 2 [8]

49. (a) (i) C2H4(g) + H2(g) → C2H6(g); 1 State symbols not required for mark

(ii) products more stable than reactants/reactants less stable than products; products lower in energy/reactants higher in energy; 2

(iii) (overall) bonds in reactants weaker/(overall) bonds in product stronger /all bonds in product are ⌠ bonds/weaker bond broken and a (stronger) ⌠ bond formed;

less energy needed to break weaker bonds/more energy produced to make stronger bonds (thus reaction is exothermic)/OWTTE;

OR

bond breaking is endothermic/requires energy and bond making is exothermic/releases energy; stronger bonds in product mean process is exothermic overall; 2

[5] 50. (i) energy required to break (a mole of) bonds in the gaseous state /energy given out when (a mole of)

bonds are made in the gaseous state; average value from a number of similar compounds; 2

(ii) (⊗HӨreaction = (∑BEbreak − BEmake))

= [(837) + 2(436)] − [(348 + 4(412)];

= − 287(kJ/kJ mol−1); 2 Award [1 max] for 287 or + 287.

(iii) (BE): C−Cl > C−Br > C−I/C−X bond becomes weaker; halogen size/radius increases/bonding electrons further away from the nucleus/bonds

become longer; 2 [6]

IB Chemistry – SL

Topic 10 Questions

1. Which of the structures below is an aldehyde?

A.

B.

C.

D. (Total 1 mark)

2. What product results from the reaction of CH2==CH2 with Br2?

A. CHBrCHBr

B. CH2CHBr

C. CH3CH2Br

D. CH2BrCH2Br (Total 1 mark)

3. What is the final product formed when CH3CH2OH is refluxed with acidified potassium dichromate(VI)?

A. CH3CHO

B. CH2==CH2

C. CH3COOH

D. HCOOCH3 (Total 1 mark)

4. Which substance(s) could be formed during the incomplete combustion of a hydrocarbon?

I. Carbon II. Hydrogen III. Carbon monoxide

A. I only

B. I and II only

C. I and III only

D. II and III only (Total 1 mark)

5. Which formulas represent butane or its isomer?

I. CH3(CH2)2CH3

II. CH3CH(CH3)CH3

III. (CH3)3CH

A. I and II only

B. I and III only

C. II and III only

D. I, II and III (Total 1 mark)

6. Which statement about neighbouring members of all homologous series is correct?

A. They have the same empirical formula.

B. They differ by a CH2 group.

C. They possess different functional groups.

D. They differ in their degree of unsaturation. (Total 1 mark)

7. What is the IUPAC name for CH3CH2CH(CH3)2?

A. 1,1-dimethylpropane

B. 2-methylbutane

C. isopentane

D. ethyldimethylmethane (Total 1 mark)

8. Which compound has the lowest boiling point?

A. CH3CH2CH(CH3)CH3

B. (CH3)4C

C. CH3CH2CH2CH2CH3

D. CH3CH2OCH2CH3 (Total 1 mark)

9. What type of reaction does the equation below represent?

CH2=CH2 + Br2 → BrCH2CH2Br

A. substitution

B. condensation

C. reduction

D. addition (Total 1 mark)

10. Which compound is a member of the same homologous series as 1-chloropropane?

A. 1-chloropropene

B. 1-chlorobutane

C. 1-bromopropane

D. 1,1-dichloropropane (Total 1 mark)

11. Which formula is a correct representation of pentane?

A. CH3CH2CHCH2CH3

B. (CH3CH2)2CH3

C. CH3(CH2)3CH3

D. CH3(CH3)3CH3 (Total 1 mark)

12. How many structural isomers are possible with the molecular formula C6 H14?

A. 4

B. 5

C. 6

D. 7 (Total 1 mark)

13. Which compound is a member of the aldehyde homologous series?

A. CH3COCH3

B. CH3CH2CH2OH

C. CH3CH2COOH

D. CH3CH2CHO (Total 1 mark)

14. Which type of compound can be made in one step from a secondary alcohol?

A. an aldehyde

B. an alkane

C. a carboxylic acid

D. a ketone (Total 1 mark)

15. Which formula represents a tertiary alcohol?

(Total 1 mark)

16. Which reaction type is typical for halogenoalkanes?

A. nucleophilic substitution

B. electrophilic substitution

C. electrophilic addition

D. nucleophilic addition (Total 1 mark)

17. Which substance is not readily oxidized by acidified potassium dichromate(VI) solution?

A. propan-1-ol

B. propan-2-ol

C. propanal

D. propanone (Total 1 mark)

18. What is the correct name of this compound?

A. 1,3-dimethylbutane

B. 2,4-dimethylbutane

C. 2-methylbutane

D. 2-methylpentane (Total 1 mark)

19. Propane, C3H8, undergoes incomplete combustion in a limited amount of air. Which products are most

likely to be formed during this reaction?

A. Carbon monoxide and water

B. Carbon monoxide and hydrogen

C. Carbon dioxide and hydrogen

D. Carbon dioxide and water (Total 1 mark)

20. What is/are the product(s) of the reaction between ethene and hydrogen bromide?

A. CH3CH2Br

B. CH3CH2Br and H2

C. CH2BrCH2Br

D. CH3BrCH2 Br and H2 (Total 1 mark)

21. Which are characteristics typical of a free radical?

I. It has a lone pair of electrons. II. It can be formed by the homolytic fission of a covalent bond. III. It is uncharged.

A. I and II only

B. I and III only

C. II and III only


22. Which of the following products could be formed from the oxidation of ethanol?

I. ethanal

II. ethanoic acid

III. ethane

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

(Total 1 mark)

23. What is the reaction type when (CH3)3CBr reacts with aqueous sodium hydroxide to form (CH3)3COH and NaBr?

A. Addition

B. Elimination

C. SN1

D. SN2 (Total 1 mark)

24. Which species is a free radical?

A. •CH3

B. +CH3

C. –CH3

D. :CH3 (Total 1 mark)

25. Which compound is a tertiary halogenoalkane?

A. (CH3CH2)2CHBr

B. CH3(CH2)3CH2Br

C. (CH3)2CHCH2CH2Br

D. CH3CH2C(CH3)2Br (Total 1 mark)

26. Which species reacts most readily with propane?

A. Br2

B. Br•

C. Br–

D. Br+ (Total 1 mark)

27. An organic compound X reacts with excess acidified potassium dichromate(VI) to form compound Y, which reacts with sodium carbonate to produce CO2(g).

What is a possible formula for compound X?

A. CH3CH2COOH

B. CH3CH2CH2OH

C. CH3CH(OH)CH3

D. (CH3)3COH (Total 1 mark)

28. Which statement about successive members of all homologous series is correct?

A. They have the same empirical formula.

B. They differ by a CH2 group.

C. They have the same physical properties.

D. They differ in their degree of unsaturation. (Total 1 mark)

29. The following is a three-dimensional representation of an organic molecule.

Which statement is correct?

A. The correct IUPAC name of the molecule is 2-methylpentane.

B. All the bond angles will be approximately 90°.

C. One isomer of this molecule is pentane.

D. The boiling point of this compound would be higher than that of pentane. (Total 1 mark)

30. Which compound forms when hydrogen bromide is added to but-2-ene?

A. 2-bromobutane

B. 2,3-dibromobutane

C. 1-bromobutane

D. 1,2-dibromobutane (Total 1 mark)

31. Which products can be potentially obtained from crude oil and are economically important?

I. Plastics II. Margarine III. Motor fuel

A. I and II only

B. I and III only

C. II and III only


32. Propane, C3H8, undergoes incomplete combustion in a limited amount of air. Which products are most likely to be formed during this reaction?

A. Carbon monoxide and water

B. Carbon monoxide and hydrogen

C. Carbon dioxide and hydrogen

D. Carbon dioxide and water (Total 1 mark)

33. Two reactions of an alkene, B, are shown below.

(i) State the name of A and write an equation for its complete combustion. Explain why the incomplete combustion of A is dangerous.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (5)

(ii) Outline a test to distinguish between A and B, stating the result in each case.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(iii) Write an equation for the conversion of B to C. State the type of reaction taking place and draw the structure of C.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(Total 11 marks

34. For the two compounds HCOOCH2CH3 and HCOOCHCH2: I II

(i) State and explain which of the two compounds can react readily with bromine.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(ii) Compound II can form polymers. State the type of polymerization compound II undergoes, and draw the structure of the repeating unit of the polymer.

(2) (Total 4 marks)

35. The compound C2H4 can be used as a starting material for the preparation of many substances.

(a) Name the compound C2H4 and draw its structural formula.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(b) In the scheme below, state the type of reaction and identify the reagent needed for each reaction.

AB C2H4 → CH3CH2OH → CH3COOH

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(c) C2H4 can be converted into one of the compounds below in a single step reaction.

C2H3Cl C2H4Cl2

Draw the structural formula for each of these compounds and identify the compound which can be formed directly from C2H4.

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(d) One of the two compounds in (c) has an isomer. Draw the structural formula of the isomer and explain why it can not be formed directly from C2H4.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(e) C2H4 can also react to form a polymer. Name this type of polymer and draw the structural formula of a section of this polymer consisting of three repeating units.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(f) Polymers can also be formed in a different type of reaction. Identify this type of reaction and name two different types of such polymers.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 16 marks)

36. The plastic PVC, poly(chloroethene), is made from the monomer chloroethene, C2H3Cl, by a polymerization reaction.

(i) Draw the structural formula of chloroethene.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (1)

(ii) State the type of polymerization reaction that occurs to make poly(chloroethene) and identify the structural feature needed in the monomer.

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(iii) Draw the structure of the repeating unit of poly(chloroethene).

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (1)

(iv) Explain why monomers are often gases or volatile liquids, whereas polymers are solids.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(Total 6 marks)

37. The hydrolysis of 2-iodo-2-methylpropane by 0.10 mol dm–3 KOH(aq) to form 2-methylpropan-2-ol is an example of nucleophilic substitution.

Give equations to illustrate the SN1 mechanism for this reaction.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 2 marks)

38. The molecular formula C4H9Br represents four structural isomers, all of which can undergo nucleophilic substitution reactions with aqueous sodium hydroxide. An equation to represent all these reactions is

C4H9Br + NaOH → C4H9OH + NaBr

(a) Explain what is meant by the term nucleophilic substitution. (2)

(b) The main mechanism for a tertiary halogenoalkane is SN1. Give the equations for this substitution reaction of the tertiary isomer of C4H9Br. Show the structures of the organic reactant and product and use curly arrows to show the movement of electron pairs.

(4)

(c) The main mechanism for a primary halogenoalkane is SN2. Give the mechanistic equation for this substitution reaction of the straight-chain primary isomer of C4H9Br, showing the structures of the organic reactant and product, and using curly arrows to show the movement of electron pairs.

(4)

(d) Give a structural formula for the secondary isomer and for the other primary isomer. State the name of each isomer.

(4) (Total 14 marks)

39. Write equations to show the mechanisms of the following reactions. In each case, show the structure of the intermediate and organic product, and use curly arrows to show the movement of electron pairs.

(i) the reaction between KOH and CH3CH2CH2CH2Cl.

(3)

(ii) the reaction between KOH and (CH3)3CCl.

(2)

(Total 5 marks)

40. Some alcohols are oxidized by heating with acidified potassium dichromate(VI). If oxidation does occur, identify the possible oxidation products formed by each of the alcohols below. Indicate if no oxidation occurs.

Butan-1-ol

..............................................................................................................................................

..............................................................................................................................................

Butan-2-ol

..............................................................................................................................................

..............................................................................................................................................

2-methylpropan-2-ol

..............................................................................................................................................

.............................................................................................................................................. (Total 4 marks)

41. Chlorine and ethane react together to form chloroethane.

(a) State the condition needed for the reaction to occur.

....................................................................................................................................

.................................................................................................................................... (1)

(b) Write equations to represent initiation, propagation and termination steps in the reaction.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (4)

(Total 5 marks)

42. CH3COCH3 is the first member of the ketone homologous series. Draw the full structural formula of the next member of this homologous series and predict how its melting point compares with that of CH3COCH3.

(Total 2 marks)

43. (i) Write an equation for the reaction between but-2-ene and bromine, showing the structure of the organic product.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(ii) State the type of reaction occurring.

.........................................................................................................................

......................................................................................................................... (1)

(Total 3 marks)

44. CH3COCH3 can be prepared in the laboratory from an alcohol. State the name of this alcohol, the type of

reaction occurring and the reagents and conditions needed for the reaction. (Total 5 marks)

45. 2-bromobutane can be converted into butan-2-ol by a nucleophilic substitution reaction. This reaction occurs by two different mechanisms.

(i) Give the structure of the transition state formed in the SN2 mechanism. (2)

(ii) Write equations for the SN1 mechanism. (2)

(Total 4 marks)

46. Ethene is an unsaturated hydrocarbon used as a starting material for many organic chemicals.

(a) Draw the structural formula of ethene and state the meaning of the term unsaturated hydrocarbon. (2)

(b) State an equation for the conversion of ethene to ethanol and identify the type of reaction. (2)

(c) Describe the complete oxidation of ethanol and name the product. Include the conditions, reagents required and any colour changes.

(4)

(d) State an equation for the reaction between ethanol and the product of complete oxidation in (c). Include any other reagent required for this reaction. Name the organic product and state one possible use of this product.


47. Ethene is an unsaturated hydrocarbon used as a starting material for many organic chemicals.

(a) State the meaning of the term unsaturated hydrocarbon. (1)

(b) State an equation for the conversion of ethene to ethanol and identify the type of reaction. (2)

(c) Describe the complete oxidation of ethanol. Include the conditions, reagents required and any colour changes. Name the organic product X.

(4)

(d) State an equation for the reaction between ethanol and compound X. Include any other reagent required. Name the organic compound Z and state one use of this product.


48. The equation for a reaction of ethane is

CH3CH3 + Cl2 → CH3CH2Cl + HCl

The mechanism of this reaction involves initiation, propagation and termination steps. Describe this reaction, including equations for each step and the role of ultraviolet light.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 5 marks)

49. (i) Draw the structural formula of propan-2-ol.

(1)

(ii) Identify the alcohol as primary, secondary, or tertiary.

.................................................................................................................................... (1)

(iii) Identify the organic product formed by the oxidation of this alcohol using acidified potassium dichromate(VI) solution.

..........................................................................................................................

(1) (Total 3 marks)

50. Secondary halogenoalkanes can undergo nucleophilic substitution reactions by both SN1 and SN2 mechanisms. The mechanism showing the formation of the transition state in the reaction between 2-bromobutane and potassium hydroxide can be represented as follows.

(a) Identify the type of mechanism shown.

.................................................................................................................................... (1)

(b) State and explain how the following changes would alter the rate of the reaction by this mechanism.

(i) using water instead of potassium hydroxide.

..........................................................................................................................

..........................................................................................................................

..........................................................................................................................

.......................................................................................................................... (2)

(ii) using bromoethane instead of 2-bromobuture.

..........................................................................................................................

..........................................................................................................................

..........................................................................................................................

.......................................................................................................................... (2)

(Total 5 marks)

51. The following is a computer-generated representation of the molecule, methyl 2-hydroxy benzoate, better known as oil of wintergreen.

(i) Deduce the empirical formula of methyl 2-hydroxy benzoate and draw the full structural formula, including any multiple bonds that may be present. The computer-generated representation shown does not distinguish between single and multiple bonds.

(2)

(ii) In this representation, two of the carbon-oxygen bond lengths shown are 0.1424 nm and 0.1373 nm. Explain why these are different and predict the carbon-oxygen bond length in carbon dioxide.

(2)

(iii) Name all the functional groups present in the molecule. (2)

(Total 6 marks)

52. (i) State and explain the trend in the boiling points of the first six alkanes involving straight-chains. (2)

(ii) Write an equation for the reaction between methane and chlorine to form chloromethane. Explain this reaction in terms of a free-radical mechanism.

(5) (Total 7 marks)

53. (i) Identify the formulas of the organic products, A–E, formed in the reactions, I–IV:

I. CH3(CH2)8OH + K2Cr2O7

II. (CH3)3CBr + NaOH C

III. (CH3)2CHOH + K2Cr2O7 D

IV. H2C=CH2 + Br2 E (5)

(ii) H2C=CH2 can react to form a polymer. Name this type of polymer and draw the structural formula of a section of this polymer consisting of three repeating units.

(2) (Total 7 marks)

54. Ethene, propene and but-2-ene are members of the alkene homologous series.

(a) Describe three features of members of a homologous series. (3)

(b) State and explain which compound has the highest boiling point. (3)

(c) Draw the structural formula and give the name of an alkene containing five carbon atoms. (2)

(d) Write an equation for the reaction between but-2-ene and hydrogen bromide, showing the structure of the organic product. State the type of reaction occurring.

(3)

(e) Propene can be converted to propanoic acid in three steps:

step1 step 2 step 3 propene propan-1-ol propanal propanoic acid

State the type of reaction occurring in steps 2 and 3 and the reagents needed. Describe how the conditions of the reaction can be altered to obtain the maximum amount of propanal, and in a separate experiment, to obtain the maximum amount of propanoic acid.

(5)

(f) Identify the strongest type of intermolecular force present in each of the compounds propan-1-ol, propanal and propanoic acid. List these compounds in decreasing order of boiling point.


55. (a) State two characteristics of a homologous series.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Describe a chemical test to distinguish between alkanes and alkenes, giving the result in each case.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (3)

(Total 5 marks)

56. The following transition state is formed during the reaction of a halogenoalkane with aqueous sodium hydroxide:

(a) Deduce the structure of the halogenoalkane. Classify it as primary, secondary or tertiary, giving a reason for your choice.

....................................................................................................................................

.................................................................................................................................... (2)

(b) The mechanism of this reaction is described as SN2. Explain what is meant by the symbols in SN2. Predict a rate expression for this reaction.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (3)

(c) The same halogenoalkane reacts with sodium hydroxide by an SN1 mechanism. Deduce the structure of the intermediate formed in this reaction.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (1)

(Total 6 marks) 57. (a) List two characteristics of a homologous series.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (1)

(b) Ethanol and ethanoic acid can be distinguished by their melting points. State and explain which of the two compounds will have a higher melting point.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(c) Draw the three isomers containing the alcohol functional group of formula C4H9OH.

(2)

(Total 5 marks

IB Chemistry – SL

Topic 10 Answers

1. A [1]

2. D [1]

3. C [1]

4. C [1]

5. D [1]

6. B [1]

7. B [1]

8. B [1]

9. D [1]

10. B [1]

11. C [1]

12. B [1]

13. D [1]

14. D [1]

15. C [1]

16. A [1]

17. D [1]

18. D [1]

19. A

[1]

20. A [1]

21. C [1]

22. A [1]

23. C [1]

24. A [1]

25. D [1]

26. B [1]

27. B [1]

28. B [1]

29. C [1]

30. A [1]

31. B [1]

32. A [1]

33. (i) butane; C4H10(g) + O2(g) → 4CO2(g) + 5H2O(l); (ignore state symbols, accept balancing using 13O2 ) [1] for all formulas and [1] for balancing equation. CO produced; CO is poisonous/combines with hemoglobin/OWTTE;

or C; which causes respiratory problems; 5

(ii) add Br2 (water); valid test needed to score further marks.

A – no effect; B – would decolorise Br2 (do not accept discolour); 3

(iii) CH3CH==CHCH3 + HBr → CH3CHBrCH2CH3; 3

[1] for HBr in balanced equation, [1] for structure of product.

addition; [11]

34. (i) II reacts with Br2 II is an alkene/has unsaturated R group/CC present, I contains only saturated R groups; 2

(ii) addition polymerization;

2 [4]

35. (a) ethene; 2

(b) A addition/hydration; H2O/water/steam; B oxidation; acidified K2Cr2O7 Accept acidified KMnO4. 4

(c)

;

The compound formed directly may be circled or indicated by some 3 other means. Accept any other structure showing a Cl atom on each C atom.

(d)

;

addition across a double bond occurs at both C atoms/OWTTE; 2 If 1,1-dichloroethane is given in (c) accept 1,2-dichloroethane as the isomer as ECF but Award [1] max;

(e) addition polymer; 2

(f) condensation polymer; polyesters; polyamides; 3 [16]

36. (i) CH2CHCl/CH2 = CHCl/; 1

(ii) addition (polymerization); (carbon-carbon) double bond/unsaturation/OWTTE; 2

(iii)

; 1

(iv) monomers have smaller molecules/surface area than polymers; with weaker intermolecular/Van

der Waals’ forces; 2 Accept opposite argument for polymers.

[6]

37. (CH3)3CI → (CH3)3C+ + I−;

(CH3)C+ + OH− → (CH3)3COH; 2 Do not allow SN 2 reaction.

[2]

38. (a) replacement of atom/group (in a molecule)/OWTTE; Do not accept substitution.

by a species with a lone pair of electrons/species attracted to an electron-deficient carbon atom; 2

(b) correct structure of (CH3)3CBr; curly arrow showing C(Br bond fission; correct structure of (CH3)3C+; curly arrow showing attack by OH− on correct C atom; correct structure of (CH3)3COH; 4

Award [1] each for any four.

(c) correct structure of CH3CH2CH2CH2Br; curly arrow showing C(Br bond fission; correct structure of transition state showing charge and all bonds; curly arrow showing attack by OH− on correct C atom; correct structure of CH3CH2CH2CH2OH; 4

Award [1] each for any four.

(d) secondary CH3CHBrCH2CH3; 2-bromobutane; other primary (CH3)2CHCH2Br; 1-bromo-2-methylpropane; 4

[14]

39. (i) (SN2 mechanism) 3

Intermediate structure showing overall negative charge and partial bonds. Accept negative charge to be indicated as delocalised between the HO-CH2-Cl.

→ HO − CH2CH2CH2CH3 + Cl−;

(ii) (SN1 mechanism) 2

[5]

40. butan-1-ol: butanal; butanoic acid;

butan-2-ol: butanone;

2 methylpropan-2-ol: no oxidation; Also accept correct structures. Where both name and structure given structure must be correct and name largely correct.

[4]

41. (a) UV light/sunlight (present); 1

(b) Throughout accept radical with or without • initiation reaction(s):

Cl2 → 2Cl•; 1

propagation reactions:

Cl• + CH3CH3 → CH3CH2• + HCl; CH3CH2• + Cl2 → CH3CH2Cl + Cl•; 2

termination reactions:

CH3CH2• + Cl• → CH3CH2Cl; 2Cl• → Cl2; 2CH3CH2• → CH3CH2CH2CH3; 1

Award [1] for any termination reaction. If initiation, propagation, termination not labelled or incorrectly labelled award [3] max.

[5]

42.

[2]

43. (i) C4H8 + Br2 → C4H8Br2; Equation scores [1].

CH3CHBrCHBrCH3; 2 Accept more detailed formula.

(ii) addition; 1 [3]

44. propan-2-ol; Accept 2-propanol.

oxidation/redox; (potassium/sodium) dichromate(VI)/potassium manganate(VII); Accept just dichromate, permanganate, KMnO4 , Mn, K2Cr2O7, Cr2 (sulfuric) acid; heat under reflux;

[5]

45. (i)

all five groups around C correct; negative charge and dotted lines to OH and Br correct; 2

Do not award 2nd mark if bond from OH (i.e. OH-----).

(ii) CH3CH2CH(CH3)Br → CH3CH2CH(CH3)+ + Br−; CH3CH2CH(CH3)+ + OH− → CH3CH2CH(CH3)OH; 2

Accept C4H9 instead of CH3CH2CH(CH3) throughout. [4]

46. (a)

Allow CH2=CH2. a hydrocarbon that contains at least one C=C (or C≡C)/carbon-carbon double bond (or triple

bond)/carbon to carbon multiple bond; 2 Do not accept just “double bond”.

(b) C2H4 + H2O → C2H5OH; addition/hydration reaction; 2

(c) heat under reflux;

EITHER

potassium dichromate(VI)/K2Cr2O7/Cr2O72–and acidified/H++;

orange to green;

OR potassium permanganate/manganate(VII)/KMnO4/MnO4

– and acidified/H++; purple to colourless;

Penalize wrong oxidation state, but not missing oxidation state.

ethanoic acid; 4

(d) CH3COOH + C2H5OH → CH3COOCH2CH3 + H2O;

Accept CH3COOC2H5 sulfuric acid/H2SO4/(ortho)phosphoric acid/H3PO4;

ethyl ethanoate; solvent/flavouring/perfumes/plasticizers;. 4

[12]

47. (a) a hydrocarbon that contains at least one C=C (or C≡C)/carbon-carbon double bond (or triple bond)/carbon to carbon multiple bond; 1

Do not accept just “double bond”.

(b) C2H4 + H2O → C2H5OH; addition/hydration reaction; 2

(c) heat under reflux;

EITHER

potassium dichromate(VI)/K2Cr2O7 / Cr2O72– and acidified/H+;

orange to green;

OR

potassium permanganate/manganate(VII)/KMnO4 / MnO4– and acidified/H+;

purple to colourless; Penalize wrong oxidation state, but not missing oxidation state.

ethanoic acid; 4

(d) CH3COOH + C2H5OH CH3COOCH2CH3 + H2O; accept equations including H+. Reversible arrow not required for the mark.

sulfuric acid/H2SO4/(ortho)phosphoric acid/H3PO4; Z – ethyl ethanoate; solvent/flavouring/perfumes/plasticizers; 4

[11]

48. ultraviolet light causes Cl−Cl bond to split; Cl2 → 2Cl•; Cl• + CH3CH3 → CH3CH2• + HCl; CH3CH2• + Cl2 → CH3CH2Cl + Cl• CH3CH2• + Cl• → CH3CH2Cl/other correct termination step; 5

Penalize missing � symbol once only. If different alkane used, then deduct [1]. No penalty for not labelling steps, but deduct [1] if any wrongly labelled.

[5]

49. (i)

1 Allow bond to HO rather than OH or halfway between the two

(ii) secondary; 1

(iii) CH3COCH3/propanone/acetone; 1 Allow ECF from a different alcohol drawn in (i)

[3]

50. (a) SN2 / bimolecular; 1

(b) (i) reaction slower; neutral/uncharged/less polar/electrons donated less easily in H2O; 2

(ii) reaction faster; less bulky group/reduced steric hindrance; 2

[5]

51. (i) (Empirical formula =) C8H8O3; 2 Allow double bonds on arene in alternate positions, or allow delocalized representation (of pi electrons).

(ii) the bond at 0.1373 nm is a double bond and the bond at 0.1424 nm is a single bond; in CO2(g) both bonds are double bonds and would have a value around 0.137 nm; 2

(iii) Ester; Arene/benzene ring; Alcohol; 2 Award 2 for any three correct, award [1] for any two correct. Do not accept alkane as a type of functional group in this molecule.

[6]

52. (i) boiling point increases as the number of carbons increases/OWTTE; Greater Mr and hence greater van der Waals’/London/dispersion forces present; 2

(ii) CH4 + Cl2 CH3Cl + HCl; Do not award mark if hv/uv light is not given.

Initiation step: Cl2 2Cl•; Do not award mark if hv/uv light is not given. Penalize once only.

Propagation step: CH4 + Cl• → CH3• + HCl; CH3• + Cl2 → CH3Cl + Cl•; Termination step: Cl• + Cl• → Cl2 or Cl• + CH3• → CH3Cl or CH3• + CH3• → CH3CH3; 5

Allow fish-hook half-arrow representations i.e. use of . Penalize use of full curly arrows once only. Penalize missing dots on radicals once only.

[7]

53. (i) A. = CH3(CH2)7CHO; B. = CH3(CH2)7COOH/CH3(CH2)7CO2H; C. = (CH3)3COH; D. =

(CH3)2CO; E. = BrCH2CH2Br; 5 Allow correct structural formulas.

(ii) addition; /-(CH2-CH2)3-/-(CH2)6-; 2 [7]

54. (a) same general formula/CnH2n; formulas of successive members differ by CH2; similar chemical properties/same functional group; gradation/gradual change in physical properties; 3

Award [1] each for any three.

(b) but-2-ene; Accept 2-butene.

strongest intermolecular/van der Waals’ forces; largest (molecular) mass/size/surface area/area of contact; 3

(c) CH2CHCH2CH2CH3/CH3CHCHCH2CH3/any correct branched structure; Accept more detailed formula.

pent-1-ene/pent-2-ene; 2 Name must match formula. Accept 1-pentene/2-pentene.

(d) C4H8 + HBr → CH3CH2CHBrCH3; Award [1] for all molecular formulas correct and [1] for correct product structure. Award [1] for completely correct equation starting with but-1-ene.

addition; 3

(e) oxidation/redox; (potassium) dichromate(VI)/; (sulfuric) acid; distilling off propanal as it is formed; heating under reflux (to obtain propanoic acid); 5

(f) (propan-1-ol) hydrogen bonding; (propanal) dipole-dipole attractions; (propanoic acid) hydrogen bonding; propanoic acid > propan-1-ol > propanal; 4

[20]

55. (a) same general formula; successive members differ by CH2;

Do not allow elements or just “they”. similar chemical properties;

Allow same/constant. gradual change in physical properties;

Do not allow change periodically. same functional group; 2

Award [1] each for any two.

(b) add bromine (water); alkanes − no change/stays or turns brown;

Allow red-brown or any combination of brown, orange or yellow. alkenes − bromine (water) decolorizes;

Do not allow clear or discoloured.

or

add (acidified) KMnO4; alkanes − no change; alkenes − KMnO4 decolorizes/brown/black; 3

[5]

56. (a) (CH3)2CHBr/more detailed formula;

secondary/2° because two alkyl groups attached to C with Br; 2

(b) nucleophilic substitution; bimolecular/molecularity of two/two species in rate-determining step;

Accept second order.

rate = k [(CH3)2CHBr][OH−]; 3 No penalty for incorrect halogenoalkane formula.

(c) (CH3)2CH+/more detailed formula; 1 [6]

57. (a) one general formula/same general formula; differ by CH2; similar chemical properties; gradual

change in physical properties; 1 Award [1] for any two of the above characteristics.

(b) ethanol lower/ethanoic acid higher;

due to larger mass of ethanoic acid/stronger van der Waals’/ London/dispersion forces; due to stronger hydrogen bonding/2 hydrogen bonds per molecule; 2

Accept either answer

(c) 2 Allow condensed structural formulas such as CH3CH2CH2CH2OH. Award [2] for all three correct isomers, [1] for any two correct isomers.

[5]

Topic 9 Questions

1. Which statement is correct?

A. Oxidation involves loss of electrons and a decrease in oxidation state.

B. Oxidation involves gain of electrons and an increase in oxidation state.

C. Reduction involves loss of electrons and an increase in oxidation state.

D. Reduction involves gain of electrons and a decrease in oxidation state. (Total 1 mark)

2. What occurs during the operation of a voltaic cell based on the following reaction?

Ni(s) + Pb2+(aq) → Ni2+(aq) + Pb(s)

External circuit Ion movement in solution

A. electrons move from Ni to Pb Pb2+(aq) move away from Pb(s)

B. electrons move from Ni to Pb Pb2+(aq) move toward Pb(s)

C. electrons move from Pb to Ni Ni2+(aq) move away from Ni(s)

D. electrons move from Pb to Ni Ni2+(aq) move toward Ni(s) (Total 1 mark)

3. The oxidation number of chromium is the same in all the following compounds except

A. Cr(OH)3

B. Cr2O3

C. Cr2(SO4)3

D. CrO3 (Total 1 mark)

4. Magnesium is a more reactive metal than copper. Which is the strongest oxidizing agent?

A. Mg

B. Mg2+

C. Cu

D. Cu2+ (Total 1 mark)

5. Which processes occur during the electrolysis of molten sodium chloride?

I. Sodium and chloride ions move through the electrolyte.

II. Electrons move through the external circuit.

III. Oxidation takes place at the positive electrode (anode).

A. I and II only

B. I and III only

C. II and III only


6. What happens to the Cr3+(aq) ion when it is converted to CrO42–(aq)?

A. Its oxidation number decreases and it undergoes reduction.

B. Its oxidation number decreases and it undergoes oxidation.

C. Its oxidation number increases and it undergoes reduction.

D. Its oxidation number increases and it undergoes oxidation. (Total 1 mark)

7. The following reactions are spontaneous as written.

Fe(s) + Cd2+(aq) → Fe2+(aq) + Cd(s)

Cd(s) + Sn2+(aq) → Cd2+(aq) + Sn(s)

Sn(s) + Pb2+(aq) → Sn2+(aq) + Pb(s)

Which of the following pairs will react spontaneously?

I. Sn(s) + Fe2+(aq)

II. Cd(s) + Pb2+(aq)

III. Fe(s) + Pb2+(aq)

A. I only

B. II only

C. III only


8. What species are produced at the positive and negative electrodes during the electrolysis of molten sodium chloride?

Positive electrode Negative electrode

A. Na+(l) Cl2(g)

B. Cl–(l) Na+(l)

C. Na(l) Cl2(g)

D. Cl2(g) Na(l) (Total 1 mark)

9. Consider the following reaction.

H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq)


A. H2SO3 is the reducing agent because it undergoes reduction.

B. H2SO3 is the reducing agent because it undergoes oxidation.

C. Sn4+ is the oxidizing agent because it undergoes oxidation.

D. Sn4+ is the reducing agent because it undergoes oxidation. (Total 1 mark)

10. In which change does oxidation occur?

A. CH3CHO → CH3CH2OH

B. CrO42–→ Cr2O7

2–

C. SO42–→ SO3

2–

D. NO2– → NO3

– (Total 1 mark)

11. What happens at the positive electrode in a voltaic cell and in an electrolytic cell?

Voltaic cell Electrolytic cell A. Oxidation Reduction B. Reduction Oxidation C. Oxidation Oxidation D. Reduction Reduction

(Total 1 mark)

12. What are the oxidation numbers of the elements in sulfuric acid, H2SO4?

Hydrogen Sulfur Oxygen

A. +1 +6 –2

B. +1 +4 –2

C. +2 +1 +4

D. +2 +6 –8 (Total 1 mark)

13. A voltaic cell is made from copper and zinc half-cells. The equation for the reaction occurring in the cell is

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Which statement is correct when the cell produces electricity?

A. Electrons are lost from zinc atoms.

B. The mass of the copper electrode decreases.

C. Electrons flow from the copper half-cell to the zinc half-cell.

D. Negative ions flow through the salt bridge from the zinc half-cell to the copper half-cell. (Total 1 mark)

14. What happens when molten sodium chloride is electrolysed in an electrolytic cell?

A. Chlorine is produced at the positive electrode.

B. Sodium ions lose electrons at the negative electrode.

C. Electrons flow through the liquid from the negative electrode to the positive electrode.

D. Oxidation occurs at the negative electrode and reduction at the positive electrode. (Total 1 mark)

15. Which equations represent reactions that occur at room temperature?

I. 2Br–(aq) + Cl2(aq) → 2Cl–(aq) + Br2(aq)

II. 2Br–(aq) + I2(aq) → 2I–(aq) + Br2(aq)

III. 2I–(aq) + Cl2(aq) → 2Cl–(aq) + I2(aq)

A. I and II only

B. I and III only

C. II and III only


16. Which equation represents a redox reaction?

A. KOH(aq) + HCl(aq) → KCl(aq) + H2O(l)

B. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

C. CuO(s) + 2HCl(aq) → CuCl2(aq) + H2O(l)

D. ZnCO3(s) + 2HCl(aq) → ZnCl2(aq) + CO2(g) + H2O(l) (Total 1 mark)

17. The following information is given about reactions involving the metals X, Y and Z and solutions of their sulfates.

X(s) + YSO4(aq) → no reaction Z(s) + YSO4(aq) → Y(s) + ZSO4(aq)

When the metals are listed in decreasing order of reactivity (most reactive first), what is the correct order?

A. Z > Y > X

B. X > Y > Z

C. Y > X > Z

D. Y > Z > X (Total 1 mark)

18. What are the oxidation numbers of the elements in the compound phosphoric acid, H3PO4?

Hydrogen Phosphorus Oxygen

A. +1 +1 –2

B. +1 +5 –2

C. +3 +1 –4

D. +3 +5 –8 (Total 1 mark)

19. A voltaic cell is made from magnesium and iron half-cells. Magnesium is a more reactive metal than iron. Which statement is correct when the cell produces electricity?

A. Electrons are lost from magnesium atoms.

B. The concentration of Fe2+ ions increases.

C. Electrons flow from the iron half-cell to the magnesium half-cell.

D. Negative ions flow through the salt bridge from the magnesium half-cell to the iron half-cell. (Total 1 mark)

20. Which are examples of reduction?

I. Fe3+ becomes Fe2+

II. Cl– becomes Cl2

III. CrO3 becomes Cr3+

A. I and II only

B. I and III only

C. II and III only


21. Which statement is correct for the electrolysis of a molten salt?

A. Positive ions move toward the positive electrode.

B. A gas is produced at the negative electrode.

C. Only electrons move in the electrolyte.

D. Both positive and negative ions move toward electrodes. (Total 1 mark)

22. Which statement about the following reaction is correct?

2Br–(aq) + Cl2(aq) → Br2(aq) + 2Cl–(aq)

A. Br–(aq) is reduced and gains electrons.

B. Cl2(aq) is reduced and loses electrons.

C. Br–(aq) is oxidized and loses electrons.

D. Cl2(aq) is oxidized and gains electrons. (Total 1 mark)

23. Consider the following spontaneous reactions.

Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s) Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s) Zn(s) + Fe2+(aq) → Zn2+(aq) + Fe(s)

Which is the correct combination of strongest oxidizing agent and strongest reducing agent?

Strongest oxidizing agent Strongest reducing agent

A. Ag(s) Zn(s)

B. Ag+(aq) Zn(s)

C. Zn2+(aq) Ag(s)

D. Zn(s) Ag+(aq) (Total 1 mark)

24. In which change does nitrogen undergo oxidation?

A. NO2 → N2O4

B. NO3– → NO2

C. N2O5→ NO3–

D. NH3 → N2 (Total 1 mark)

25. Which statement is correct?

A. Spontaneous redox reactions produce electricity in an electrolytic cell.

B. Electricity is used to carry out a non-spontaneous redox reaction in a voltaic cell.

C. Oxidation takes place at the negative electrode in a voltaic cell and the positive electrode in an electrolytic cell.

D. Oxidation takes place at the negative electrode in a voltaic cell and reduction takes place at the positive electrode in an electrolytic cell.

(Total 1 mark)

26. The compound [Co(NH3)5Br]SO4 is isomeric with the compound [Co(NH3)5SO4]Br. What is the oxidation state of cobalt in these compounds?

[Co(NH3)5Br]SO4 [Co(NH3)5SO4]Br

A. +3 +3

B. +2 +1

C. +3 +2

D. +2 +3 (Total 1 mark)

27. What happens to vanadium during the reaction VO2+(aq) → VO3–(aq)?

A. It undergoes oxidation and its oxidation number changes from +4 to +5.

B. It undergoes oxidation and its oxidation number changes from +2 to +4.

C. It undergoes reduction and its oxidation number changes from +2 to –1.

D. It undergoes reduction and its oxidation number changes from +4 to +2. (Total 1 mark)

28. What occurs during the electrolysis of a molten salt?

A. Electricity is produced by a spontaneous redox reaction.

B. Electricity is used to cause a non-spontaneous redox reaction to occur.

C. Electrons flow through the molten salt.

D. Electrons are removed from both ions of the molten salt. (Total 1 mark)

29. Which statement is correct about an oxidizing agent in a chemical reaction?

A. It reacts with oxygen.

B. It reacts with H+ ions.

C. It loses electrons.

D. It undergoes reduction. (Total 1 mark)

30. Which formula represents an aldehyde?

A. CH3CH2CHO

B. CH3COCH3

C. CH3CH2COOH

D. CH3COOCH3 (Total 1 mark)

31. What is the reducing agent in this reaction?

Cu(s) + (aq) + 4H+(aq) → Cu2+(aq) + 2NO2(g) + 2H2O(l)

A. Cu(s)

B. (aq)

C. Cu2+(aq)

D. H+(aq) (Total 1 mark)

32. A particular voltaic cell is made from magnesium and iron half-cells. The overall equation for the reaction occurring in the cell is

Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

Which statement is correct when the cell produces electricity?

A. Magnesium atoms lose electrons.

B. The mass of the iron electrode decreases.

C. Electrons flow from the iron half-cell to the magnesium half-cell.

D. Negative ions flow through the salt bridge from the magnesium half-cell to the iron half-cell. (Total 1 mark)

33. What process occurs at the cathode in a voltaic cell and at the anode in an electrolytic cell?

Cathode of voltaic cell Anode of Electrolytic cell A. Oxidation Reduction

B. Oxidation Oxidation

C. Reduction Oxidation

D. Reduction Reduction (Total 1 mark)

34. Consider the following reaction:

H2SO3(aq) + Sn4+(aq) + H2O(l) → Sn2+(aq) + HSO4–(aq) + 3H+(aq)


A. H2SO3 is the reducing agent because it undergoes reduction.

B. H2SO3 is the reducing agent because it undergoes oxidation.

C. Sn4+ is the oxidizing agent because it undergoes oxidation.

D. Sn4+ is the reducing agent because it undergoes oxidation. (Total 1 mark)

35. Which processes occur during the electrolysis of molten sodium chloride?

I. Sodium and chloride ions move through the electrolyte. II. Electrons move through the external circuit. III. Oxidation takes place at the anode.

A. I and II only

B. I and III only

C. II and III only


36. Tin(II) ions can be oxidized to tin(IV) ions by acidified potassium permanganate(VII) solution according to the following unbalanced equation.

__ Sn2+ + __MnO4– + __ H+ →__ Sn4+ + __Mn2+ + __ H2O

(a) Identify the oxidizing agent and the reducing agent.

Oxidizing agent .........................................................................................................

Reducing agent .......................................................................................................... (1)

(b) Balance the equation above.

....................................................................................................................................

.................................................................................................................................... (1)

(Total 2 marks)

37. Consider the following redox equation.

5Fe2+(aq) +MnO4–(aq) +8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2O(l)

(i) Determine the oxidation numbers for Fe and Mn in the reactants and in the products.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(ii) Based on your answer to (i), deduce which substance is oxidized.

…………………………………………………………………………………………… (1)

(iii) The compounds CH3OH and CH2O contain carbon atoms with different oxidation numbers. Deduce the oxidation numbers and state the kind of chemical change needed to make CH2O from CH3OH.

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 6 marks)

38. A part of the reactivity series of metals, in order of decreasing reactivity, is shown below.

magnesium zinc iron lead copper silver

If a piece of copper metal were placed in separate solutions of silver nitrate and zinc nitrate

(i) determine which solution would undergo reaction.

…………………………………………………………………………………………… (1)

(ii) identify the type of chemical change taking place in the copper and write the half-equation for this change.

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(iii) state, giving a reason, what visible change would take place in the solutions.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(Total 5 marks)

39. (i) Solid sodium chloride does not conduct electricity but molten sodium chloride does. Explain this difference, and outline what happens in an electrolytic cell during the electrolysis of molten sodium chloride using carbon electrodes.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(ii) State the products formed and give equations showing the reactions at each electrode.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(iii) State what practical use is made of this process.

……………………………………………………………………………………………

…………………………………………………………………………………………… (1)

(Total 9 marks)

40. Electrolysis can be used to obtain chlorine from molten sodium chloride. Write an equation for the reaction occurring at each electrode and describe the two different ways in which electricity is conducted when the cell is in operation.

(Total 4 marks)

41. Two reactions occurring in the manufacture of bromine from sea water are

I Cl2(g) + 2Br–(aq) → 2Cl–(aq) + Br2(g) II Br2(g) + SO2(g) + 2H2O(l) → 2HBr(g) + H2SO4(g)

(i) Explain, by reference to electrons, why reaction I is referred to as a redox reaction. (2)

(ii) State and explain whether SO2 is reduced or oxidized in reaction II by referring to the oxidation numbers of sulfur in this reaction.

(3) (Total 5 marks)

42. In terms of electron transfer define:

(i) oxidation

....................................................................................................................................

.................................................................................................................................... (1)

(ii) oxidizing agent

....................................................................................................................................

.................................................................................................................................... (1)

(Total 2 marks)

43. Deduce the change in oxidation number of chromium in the below reaction. State with a reason whether the chromium has been oxidized or reduced.

CrO72− + 14H+ + 6Fe2+ → 2Cr3+ + 6Fe3+ + 7H2O

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 2 marks)

44. (a) (i) Define oxidizing agent in terms of electron transfer.

.........................................................................................................................

......................................................................................................................... (1)

(ii) Deduce the change in oxidation number of chromium in the reaction below. State with a

reason whether the chromium has been oxidized or reduced.

Cr2O7 2− + 14H+ + 6Fe2+ → 2Cr3+ + 7H2O

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(Total 3 marks)

45. Iron in food, in the form of Fe3+, reacts with ascorbic acid (vitamin C), C6H8O6, to form dehydroascorbic acid, C6H6O6.

(i) Write an ionic half-equation to show the conversion of ascorbic acid to dehydroascorbic acid in aqueous solution.

.........................................................................................................................

......................................................................................................................... (1)

(ii) In the other ionic half-equation Fe3+ is converted to Fe2+. Deduce the overall equation for the reaction between C6H8O6 and Fe3+.

.........................................................................................................................

......................................................................................................................... (1)

(Total 2 marks)

46. (i) Draw a diagram of apparatus that could be used to electrolyse molten potassium bromide. Label the diagram to show the polarity of each electrode and the product formed.

(3)

(ii) Describe the two different ways in which electricity is conducted in the apparatus. (2)

(iii) Write an equation to show the formation of the product at each electrode. (2)

(Total 7 marks) 47. Iodide ions, I–(aq), react with iodate ions, IO3

–(aq), in an acidic solution to form molecular iodine and water.

(i) Determine the oxidation number of iodine in each iodine-containing species in the reaction.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(ii) Identify, with a reason, the species that undergoes:

oxidation ....................................................................................................................

....................................................................................................................................

reduction ....................................................................................................................

.................................................................................................................................... (2)

(Total 4 marks)

IB Chemistry – SL

Topic 9 Answers

1. D [1]

2. B [1]

3. D [1]

4. D [1]

5. D [1]

6. D [1]

7. D [1]

8. D [1]

9. B [1]

10. D [1]

11. B [1]

12. A [1]

13. A [1]

14. A [1]

15. B [1]

16. B [1]

17. A [1]

18. B [1]

19. A [1]

20. B [1]

21. D [1]

22. C [1]

23. B [1]

24. D [1]

25. C [1]

26. A [1]

27. A [1]

28. B [1]

29. D [1]

30. A [1]

31. A [1]

32. A [1]

33. C [1]

34. B [1]

35. D [1]

36. (a) oxidizing agent: (acidified) potassium permanganate(VII)/(H+) and Mn and reducing agent: Sn2+; 1 Both oxidizing agent and reducing agent required for [1].

(b) 5Sn2+ + 2MnO4– + 16H+ → 5Sn4+ + 2Mn2+ + 8H2O; 1

[2]

37. (i) Fe reactant +2 AND Fe product +3 AND Mn product +2; Mn reactant +7; 2 Do not accept Roman numerals.

(ii) Fe2+/iron(ii) ions/ferrous ions; 1 Do not accept “iron”.

(iii) CH3OH oxidation state –2; CH2O oxidation state 0; (change is) oxidation/dehydrogenation; 3 [6]

38. (i) silver nitrate; 1

(ii) oxidation; Cu → Cu2+ + 2e; 2

(iii) (silver nitrate) solution turns blue/grey or black or silver solid forms; copper ions form/Cu2+ ions form/silver deposited; 2

[5]

39. (i) sodium chloride crystals consist of ions in a rigid lattice/ions can not move about; when melted the ions are free to move or ions move when a voltage is applied; in electrolysis positive sodium ions or Na+ ions move to the negative electrode or cathode; and negative chloride ions or Cl– move to the positive electrode or anode; 4

(ii) sodium formed at cathode or negative electrode; Na+ + e → Na; chlorine formed at anode or positive electrode; 4 2Cl– → Cl2 + 2e;

1st and 3rd marks can be scored in (i).

(iii) manufacture of sodium and chlorine/one stated use of chlorine or sodium; 1 [9]

40. at negative electrode (cathode)

Na+ + e− → Na;

at positive electrode (anode)

2Cl− → Cl2 + 2e−; If both equations correct but electrodes incorrect or not stated, then deduct [1].

electrons flow through the external circuit or wires; ions gain/lose electrons at electrodes/ions move to electrodes. 4

[4]

41. (i) chlorine/Cl2 gains electrons and is reduced; bromide (ions)/Br− loses electrons and is oxidized; 2

Award [1] max if no mention of reduced and oxidized.

(ii) S in SO2

+4;

S in H2SO4

+6; Award only [1] for 4 + and 6 + or 4 and 6.

SO2 oxidized because oxidation number (of sulfur) increases; 3 [5]

42. (i) loss of electrons; 1

(ii) (a species that) gains electrons (from another species)/causes electron loss; 1 [2]

43. changes by 3; reduced because its oxidation number decreased / +6 → +3 / 6+ → 3+ / it has gained electrons;

[2]

44. (i) (a species that) gains electrons (from another species)/causes electron loss; 1

(ii) changes by 3; reduced because its oxidation number decreased / +6 → +3 / 6+ → 3+ / it has gained electrons; 2

[3]

45. (i) C6H8O6 → C6H6O6 +2H+ + 2e; 1

(ii) C6H8O6 + 2Fe3+ → C6H6O + 2H+ + 2Fe2+; 1 [2]

46. (i) (diagram showing)

container, liquid, electrodes and power supply; bromine formed at + electrode; potassium formed at − electrode; 3

Award [1] for both correct products shown at wrong electrodes, or if no polarity indicated.

(ii) electrons flow through connecting wires; ions move (through liquid) to electrodes (and lose/gain electrons); 2

(iii) K+ + e− → K; 2Br− → Br2 + 2e−; 2

No need to indicate polarity of electrodes.

Accept e instead of e−. [7]

47. (i) I− = −1 / 1−

IO3− = +5 / 5+

I2 = 0 2 Award [2] for all three correct, [1] for any two correct, Signs must be included Do not accept Roman numerals

(ii) oxidation

I− (to I2), increase in oxidation number/loss of electron(s);

reduction

IO3– (to I2), decrease in oxidation number/gain of electron(s); 2

[4]

Topic 7 Questions

1. I2(g) + 3Cl2(g) 2ICl3(g)

What is the equilibrium constant expression for the reaction above?

A. Kc =

B. Kc =

C. Kc =

D. Kc =

2. 2SO2(g) + O2 (g) 2SO3(g) ∆Hο = –200 kJ

According to the above information, what temperature and pressure conditions produce the greatest amount of SO3?

Temperature Pressure

A. low low

B. low high

C. high high

D. high low

3. Which statement(s) is/are true for a mixture of ice and water at equilibrium?

I. The rates of melting and freezing are equal.

II. The amounts of ice and water are equal.

III. The same position of equilibrium can be reached by cooling water and heating ice.

A. I only

B. I and III only

C. II only

D. III only

4. What will happen to the position of equilibrium and the value of the equilibrium constant when the temperature is increased in the following reaction?

Br2(g) + Cl2(g) 2BrCl(g) ∆H = +14 kJ

Position of equilibrium Value of equilibrium constant

A. Shifts towards the reactants Decreases

B. Shifts towards the reactants Increases

C. Shifts towards the products Decreases

D. Shifts towards the products Increases

5. Which statement concerning a chemical reaction at equilibrium is not correct?

A. The concentrations of reactants and products remain constant.

B. Equilibrium can be approached from both directions.

C. The rate of the forward reaction equals the rate of the reverse reaction.

D. All reaction stops.

6. In the reaction below

N2(g) + 3H2(g) 2NH3(g) ∆H = –92 kJ

which of the following changes will increase the amount of ammonia at equilibrium?

I. Increasing the pressure

II. Increasing the temperature

III. Adding a catalyst

A. I only

B. II only

C. I and II only

D. II and III only

7. In the Haber process for the synthesis of ammonia, what effects does the catalyst have?

Rate of formation of NH3(g) Amount of NH3(g) formed A. Increases Increases B. Increases Decreases C. Increases No change D. No change Increases

8. What will happen if CO2(g) is allowed to escape from the following reaction mixture at equilibrium?

CO2(g) + H2O(l) H+(aq) + HCO3–(aq)

A. The pH will decrease.

B. The pH will increase.

C. The pH will remain constant.

D. The pH will become zero.

9. Which statements are correct for a reaction at equilibrium?

I. The forward and reverse reactions both continue.

II. The rates of the forward and reverse reactions are equal.

III. The concentrations of reactants and products are equal.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

10. The manufacture of sulfur trioxide can be represented by the equation below.

2SO2(g) + O2(g) 2SO3(g) ∆Hο = –197 kJ mol–1

What happens when a catalyst is added to an equilibrium mixture from this reaction?

A. The rate of the forward reaction increases and that of the reverse reaction decreases.

B. The rates of both forward and reverse reactions increase.

C. The value of ∆Hο increases.

D. The yield of sulfur trioxide increases.

11. Which changes will shift the position of equilibrium to the right in the following reaction?

2CO2(g) 2CO(g) +O2(g)

I. adding a catalyst

II. decreasing the oxygen concentration

III. increasing the volume of the container

A. I and II only B. I and III only

C. II and III only D. I, II and III

12. Which statement is always true for a chemical reaction that has reached equilibrium?

A. The yield of product(s) is greater than 50%.

B. The rate of the forward reaction is greater than the rate of the reverse reaction.

C. The amounts of reactants and products do not change.

D. Both forward and reverse reactions have stopped.

13. The equation for a reversible reaction used in industry to convert methane to hydrogen is shown below.

CH4(g) + H2O(g) CO(g) + 3H2(g) ⊗HӨ = +210 kJ

Which statement is always correct about this reaction when equilibrium has been reached?

A. The concentrations of methane and carbon monoxide are equal.

B. The rate of the forward reaction is greater than the rate of the reverse reaction.

C. The amount of hydrogen is three times the amount of methane.

D. The value of ⊗HӨ for the reverse reaction is –210 kJ.

14. The equation for a reaction used in the manufacture of nitric acid is

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) ⊗HӨ = –900 kJ

Which changes occur when the temperature of the reaction is increased?

Position of equilibrium Value of Kc

A. shifts to the left increases

B. shifts to the left decreases

C. shifts to the right increases

D. shifts to the right decreases

15. Which changes cause an increase in the equilibrium yield of SO3(g) in this reaction?

2SO2(g) + O2(g) 2SO3(g) ⊗HӨ = –196 kJ

I. increasing the pressure

II. decreasing the temperature

III. adding oxygen

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

16. Iron(III) ions react with thiocyanate ions as follows.

Fe3+(aq) + CNS–(aq) Fe(CNS)2+(aq)

What are the units of the equilibrium constant, Kc, for the reaction?

A. mol dm–3

B. mol2 dm–6

C. mol–1 dm3

D. mol–2 dm6

17. Consider the following equilibrium reaction in a closed container at 350°C.

SO2(g) +�Cl2(g) SO2Cl2(g) ⊗HӨ = −85 kJ


A. Decreasing the temperature will increase the amount of SO2Cl2(g).

B. Increasing the volume of the container will increase the amount of SO2Cl2(g).

C. Increasing the temperature will increase the amount of SO2Cl2(g).

D. Adding a catalyst will increase the amount of SO2Cl2(g).

18. Which of the following equilibria would not be affected by pressure changes at constant temperature?

A. 4HCl(g) + O2(g) 2H2O(g) + 2Cl2(g)

B. CO(g) + H2O(g) H2(g) + CO2(g)

C. C2H4(g) + H2O(g) C2H5OH(g)

D. PF3Cl2(g) PF3(g) + Cl2(g)

19. Consider the following equilibrium reaction in a closed container at 350°C

SO2(g) + Cl2(g) SO2Cl2(g) ⊗HӨ = −85 kJ


A. Decreasing the temperature will increase the amount of SO2Cl2(g).

B. Increasing the volume of the container will increase the amount of SO2Cl2(g).

C. Increasing the temperature will increase the amount of SO2Cl2(g).

D. Adding a catalyst will increase the amount of SO2Cl2(g).

20. What is the equilibrium constant expression, Kc, for the reaction below?

N2(g) + 2O2(g) 2NO2(g)

A. Kc =

B. Kc =

C. Kc =

D. Kc =

21. Sulfur dioxide and oxygen react to form sulfur trioxide according to the equilibrium.

2SO2(g) + O2(g) 2SO3(g)

How is the amount of SO2 and the value of the equilibrium constant for the reaction affected by an increase in pressure?

A. The amount of SO3 and the value of the equilibrium constant both increase.

B. The amount of SO3 and the value of the equilibrium constant both decrease.

C. The amount of SO3 increases but the value of the equilibrium constant decreases.

D. The amount of SO3 increases but the value of the equilibrium constant does not change.

22. The equation for the Haber process is:

N2(g) + 3H2(g) 2NH3(g) ⊗HӨ = −92.2 kJ

Which conditions will favour the production of the greatest amount of ammonia at equilibrium?

A. High temperature and high pressure

B. High temperature and low pressure

C. Low temperature and high pressure

D. Low temperature and low pressure

23. The sequence of diagrams represents the system as time passes for a gas phase reaction in which reactant X is converted to product Y.


A. At t = 5 days the rate of the forward reaction is greater than the rate of the backward reaction.

B. At t = 7 seconds the reaction has reached completion.

C. At t = 10 minutes the system has reached a state of equilibrium.

D. At t = 5 days the rate of the forward reaction is less than the rate of the backward reaction.

24. What changes occur when the temperature is increased in the following reaction at equilibrium?

Br2(g) + Cl2(g) 2BrCl(g) ∆Hο = +14 kJ mol–1

Position of equilibrium Value of equilibrium constant A. Shifts towards the reactants Decreases B. Shifts towards the reactants Increases C. Shifts towards the products Decreases D. Shifts towards the products Increases

25. The table below gives information about the percentage yield of ammonia obtained in the Haber process under different conditions.

Pressure/ Temperature/°C

atmosphere 200 300 400 500

10 50.7 14.7 3.9 1.2 100 81.7 52.5 25.2 10.6 200 89.1 66.7 38.8 18.3 300 89.9 71.1 47.1 24.4 400 94.6 79.7 55.4 31.9 600 95.4 84.2 65.2 42.3

(a) From the table, identify which combination of temperature and pressure gives the highest yield of ammonia.

………………………………………………………………………………………. (1)

(b) The equation for the main reaction in the Haber process is

N2(g) + 3H2(g) 2NH3(g) ∆H is negative

Use this information to state and explain the effect on the yield of ammonia of increasing

(i) pressure: …………………………….………………………………………..

……………………………………………………………..………………….

……………………………………………………………………………….. (2)

(ii) temperature: ………………………………………………………………….

…………………………………………………………………………….….

………………………………………………………………………………..

……………………………………………………………………………….. (2)

(c) In practice, typical conditions used in the Haber process are a temperature of 500 °C and a pressure of 200 atmospheres. Explain why these conditions are used rather than those that give the highest yield.

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(d) Write the equilibrium constant expression, Kc, for the production of ammonia.

……………………………………………………………………………………….

………………………………………………………………………………………. (1)

(Total 8 marks)

26. Consider the following equilibrium reaction.

2SO2(g) + O2(g) 2SO3(g) ∆H = –198 kJ

Using Le Chatelier’s Principle, state and explain what will happen to the position of equilibrium if

(a) the temperature increases.

.....................................................................................................................................

..................................................................................................................................... (2)

(b) the pressure increases.

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 4 marks)

27. Ammonia is produced by the Haber process according to the following reaction.

N2(g) + 3H2(g) 2NH3(g) ⊗H is negative

(a) State the equilibrium constant expression for the above reaction.

....................................................................................................................................

.................................................................................................................................... (1)

(b) Predict, giving a reason, the effect on the position of equilibrium when the pressure in the reaction vessel is increased.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(c) State and explain the effect on the value of Kc when the temperature is increased.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(d) Explain why a catalyst has no effect on the position of equilibrium.

....................................................................................................................................

.................................................................................................................................... (1)

(Total 6 marks)

28. (a) The following equilibrium is established at 1700°C.

CO2(g) + H2(g) H2O(g) CO(g)

If only carbon dioxide gas and hydrogen gas are present initially, sketch on a graph a line representing rate against time for (i) the forward reaction and (ii) the reverse reaction until shortly after equilibrium is established. Explain the shape of each line.

(7)

(b) Kc for the equilibrium reaction is determined at two different temperatures. At 850°C, Kc = 1.1 whereas at 1700°C, Kc = 4.9.

On the basis of these Kc values explain whether the reaction is exothermic or endothermic. (3)

(Total 10 marks)

29. The equation for one reversible reaction involving oxides of nitrogen is shown below:

N2O4(g) 2NO2(g) ⊗HӨ = +58 kJ

Experimental data for this reaction can be represented on the following graph:

(i) Write an expression for the equilibrium constant, Kc, for the reaction. Explain the significance of the horizontal parts of the lines on the graph. State what can be deduced about the magnitude of Kc for the reaction, giving a reason.

(4)

(ii) Use Le Chatelier’s principle to predict and explain the effect of increasing the temperature on the position of equilibrium.

(2)

(iii) Use Le Chatelier’s principle to predict and explain the effect of increasing the pressure on the position of equilibrium.

(2)

(iv) State and explain the effects of a catalyst on the forward and reverse reactions, on the position of equilibrium and on the value of Kc.


30. Consider the following reaction in the Contact process for the production of sulfuric acid for parts (a) to (d) in this question.

2SO2 + O2 2SO3

(a) Write the equilibrium constant expression for the reaction. (1)

(b) (i) State the catalyst used in this reaction of the Contact process. (1)

(ii) State and explain the effect of the catalyst on the value of the equilibrium constant and on the rate of the reaction.

(4)

(c) Use the collision theory to explain why increasing the temperature increases the rate of the reaction between sulfur dioxide and oxygen.

(2)

(d) Using Le Chatelier’s principle state and explain the effect on the position of equilibrium of

(i) increasing the pressure at constant temperature. (2)

(ii) removing of sulfur trioxide. (2)

(iii) using a catalyst. (2)

(Total 14 marks)

31. Consider the following reaction in the Contact process for the production of sulfuric acid for parts (a) to (c) in this question.

2SO2 + O2 2SO3

(a) Write the equilibrium constant expression for the reaction. (1)

(b) (i) State the catalyst used in this reaction of the Contact process. (1)

(ii) State and explain the effect of the catalyst on the value of the equilibrium constant and on the rate of the reaction.

(4)

(c) Using Le Chatelier’s principle explain the effect on the position of equilibrium of

(i) increasing the pressure at constant temperature. (2)

(ii) removing sulfur trioxide. (2)

(Total 10 marks)

32. Many reversible reactions in industry use a catalyst. State and explain the effect of a catalyst on the position of equilibrium and on the value of Kc.

(Total 4 marks)

33. The equation for a reaction used in industry is

CH4(g) + H2O(g) 3H2(g) + CO(g) ⊗HӨ = +210 kJ

Deduce the equilibrium constant expression, Kc, for this reaction. (Total 1 mark)

34. Consider the following reaction where colourless bromide ions react with colourless hydrogen peroxide to form a red-brown bromine solution.

2Br–(aq) + H2O2(aq) + 2H+(aq) Br2(aq) + 2H2O(l) ⊗H = negative

(a) Predict and explain the effect on the position of equilibrium when

(i) a small amount of sodium bromide solution is added. (2)

(ii) a small amount of sodium hydroxide solution is added.

(2)

(iii) a catalyst is added. (2)

(b) State and explain the effect on the value of the equilibrium constant when the temperature of the reaction is increased.

(2)

(c) State and explain the colour change when hydrochloric acid is added to the reaction solution at equilibrium.


35. The equation for the exothermic reaction in the Contact process is given below:

2SO2(g) + O2(g) 2SO2(g)

(i) Write the equilibrium constant expression for the reaction. (1)

(ii) State and explain qualitatively the pressure and temperature conditions that will give the highest yield of sulfur trioxide.

(4)

(iii) In practice, conditions used commercially in the Contact process are 450°C and 2 atmospheres of pressure. Explain why these conditions are used rather than those that give the highest yield.

(3)

(iv) Name a catalyst used in the Contact process. State and explain its effect on the value of the equilibrium constant.


36. In the gaseous state, methane and steam react to form hydrogen and carbon dioxide.

(i) Write an equation for the endothermic equilibrium reaction. Deduce the equilibrium expression for the reaction and state its units.

(4)

(ii) Deduce and explain the conditions of temperature and pressure under which the forward reaction is favoured.

(4)

(iii) Explain, at the molecular level, why the reaction is carried out at high pressure in industry. (2)

(Total 10 marks)

37. The diagrams below represent equilibrium mixtures for the reaction Y + X2 XY + X at 350 K and 550 K respectively. Deduce and explain whether the reaction is exothermic or endothermic.

(Total 2 marks)

38. The equation for the main reaction in the Haber process is:

N2(g) + 3H2(g) 2NH3(g) ∆H is negative

(i) Determine the equilibrium constant expression for this reaction. (1)

(ii) State and explain the effect on the equilibrium yield of ammonia with increasing the pressure and the temperature.

(4)

(iii) In practice, typical conditions used in the Haber process involve a temperature of 500°C and a pressure of 200 atm. Explain why these conditions are used rather than those that give the highest yield.

(2)

(iv) At a certain temperature and pressure, 1.1 dm3 of N2(g) reacts with 3.3 dm3 of H2(g). Calculate the volume of NH3(g), that will be produced.

(1)

(v) Suggest why this reaction is important for humanity. (1)

(vi) A chemist claims to have developed a new catalyst for the Haber process, which increases the yield of ammonia. State the catalyst normally used for the Haber process, and comment on the claim made by this chemist.


IB Chemistry – SL

Topic 7 Answers

1. D

2. B

3. B

4. D

5. D

6. A

7. C

8. B

9. A

10. B

11. C

12. C

13. D

14. B

15. D

16. C

17. A

18. B

19. A

20. C

21. D

22. C

23. C

24. D

25. (a) 200°C 600 atm. (both for [1], units not needed); 1 allow the “highest pressure and the lowest temperature”

(b) (i) yield increases/equilibrium moves to the right/more ammonia; 2 4 (gas) molecules → 2/decrease in volume/fewer molecules on right hand side;

(ii) yield decreases/equilibrium moves to the left/less ammonia; exothermic reaction/OWTTE; 2

(c) high pressure expensive/greater cost of operating at high pressure/reinforced pipes etc. needed; lower temperature – greater yield, but lowers rate;2 Do not award a mark just for the word “compromise”.

(d) Kc = (ignore units); 1 [8]

26. (a) (position of) equilibrium shifts to the left/towards reactants; (forward) reaction is exothermic/∆H is negative/the reverse reaction is endothermic/OWTTE; 2

Do not accept “Le Chatelier’s Principle” without some additional explanation.

(b) (position of) equilibrium shifts to the right/towards products; fewer gas molecules on the right hand side/volume decreases in forward reaction/OWTTE; 2

Do not accept “Le Chatelier’s Principle” without some additional explanation.

[4]

27. (a) 1 Do not allow round brackets unless Kp is used.

(b) equilibrium shifts to the right/products; 4 mol → 2 mol of gas/fewer moles of gas on the right/products; 2

(c) Kc decreases; equilibrium position shifts to the left/reactants/forward reaction is exothermic /reverse reaction is

endothermic; 2

(d) catalyst increases the rate of the forward and backward reactions equally /lowers the activation energy of both forward and backward reaction equally /lowers Ea so rate of forward and backward reactions increase; 1

[6]

28. (a)

two curves – one labelled “forward” starting up high up y-axis and one labelled “reverse” starting from zero; curves merge and become horizontal; No penalty for failing to label axes.

forward reaction: highest concentration, thus rate high to begin with; as reaction proceeds, concentrations decrease, so does rate;

reverse reaction: zero rate initially/at t = 0 (since no products present); rate increases as concentration of products increases; equilibrium established when rate of forward reaction = rate of reverse reaction; 7

(b) (reaction is) endothermic; Kc increases with (increasing) temperature; forward reaction favoured/heat used up/OWTTE; 3

[10]

29. (i)

(horizontal line) concentration of reactant and product remains constant/equilibrium reached;

(magnitude of) Kc greater than 1; Accept 1.6.

product concentration greater than reactant concentration; 4

(ii) increased temperature shifts equilibrium position to right; (forward) reaction is endothermic/absorbs heat; 2

(iii) increased pressure shifts equilibrium to left; fewer (gas) moles/molecules on left; 2

(iv) both/forward and reverse rates increased/increase in forward reverse rates are equal; activation energy reduced;

position of equilibrium unchanged; concentration/amount of reactants and products remain constant; value of Kc unchanged; Kc only affected by changes in temperature; 6

[14]

30. (a) K/Kc = [SO3]2÷[SO2]2[O2]; 1 Accept correct Kp expression.

(b) (i) vanadium(V) oxide/(di)vanadium pentaoxide/V2O5; 1 Allow just vanadium oxide but not correct formula.

(ii) catalyst does not affect the value of Kc; forward and reverse rates increase equally/by the same factor; catalyst increases the rate of the reaction;

(by providing an alternative path for the reaction with) lower activation energy; 4

(c) more energetic collisions/more molecules have energy greater than activation energy; more frequent collisions; 2

(d) (i) shifts equilibrium position to the products/right; to the side with fewer gas molecules or moles/lower volume of gas; 2

(ii) shifts equilibrium position to the products/right; to compensate for loss of SO3/produce more SO3; 2

(iii) no effect; forward and backward rates increased equally/by the same factor; 2

[14]

31. (a) K / Kc = [SO3]2÷[SO2]2[O2]; 1 Exactly as written. Accept correct Kp expression.

(b) (i) vanadium(V) oxide/(di)vanadium pentaoxide/V2O5/Pt; 1 Allow just vanadium oxide but not incorrect formula.

(ii) catalyst does not affect the value of Kc; forward and reverse rate increase equally/by the same factor; catalyst increases the rate of the reaction;

(by providing an alternative path for the reaction with) lower activation energy; 4

(c) (i) shifts equilibrium position to the products/right; to the side with least gas molecules or moles/lower volume of gas; 2

(ii) shifts equilibrium position to the products/right; to compensate for loss of SO3/produce more SO3; 2

[10]

32. no effect on position of equilibrium;

forward and reverse reactions speeded up equally/affects the rate of reaction but not the extent of the reaction; no effect on value of Kc; no change in concentrations of reactants or products/Kc only changes if temperature alters;

[4]

33. [1]

34. (a) (i) shifts to the right/toward products/forward reaction favoured; to consume excess Br− added; 2

Do not accept “due to Le Chatelier’s principle”.

(ii) shifts to the left/toward reactants/reverse reaction favoured; NaOH reacts to consume H+/an increase in the amount of H2O resulting from

neutralization; 2

(iii) no effect; catalyst increases the rate of the forward and backward reactions equally/lowers the

activation energy of both forward and backward reaction equally/lowers EA so rate of forward and backward reactions increase equally; 2

(b) equilibrium constant decreases; forward reaction is exothermic/produces heat/reverse reaction is endothermic /absorbs heat; 2

(c) colour change from red-brown to darker red-brown of Br2/red-brown colour intensifies/OWTTE; equilibrium position shifts to the right/products; to consume H+; 3

[11]

35. (i) 1

(ii) pressure high pressure (will allow system to occupy smaller volume);

Vproduct <Vreactant/equilibrium moves to the right to reduce pressure /reaction proceeds to lower/lowest number of gaseous molecules /OWTTE;

Temperature low temperature; (exothermic reaction) forward reaction favoured to replace some of the heat removed/equilibrium moves to the right to produce heat /OWTTE; 4

No mark for just saying “due to Le Chatelier's principle”

(iii) rate is faster at 450°C (than at low temperatures); >95%/90 − 99% yield/(very) high conversion takes place; unnecessary to use expensive high pressure equipment/(to achieve) high pressure is very expensive; 3

(iv) vanadium pentoxide/vanadium(V) oxide/V2O5/finely divided platinum/Pt; no effect on Kc; forward and reverse rates speeded up (equally); 3

[11]

36. (i) CH4(g) + 2H2O(g) 4H2(g) + CO2(g); States not required. Award [1] for balanced equation and [1] for

equilibrium sign.

Kc = ;

units: mol2 dm−6/mol2 L−2/mol2 l−2; do not accept: M2 4

(ii) (endothermic reaction) increase in temperature (favours the forward reaction); absorbs (some of) the heat supplied/OWTTE;

Award no marks for saying: “because of Le Chatelier’s principle”.

low pressure (will allow system to occupy more volume); Vproduct > Vreactant/reaction proceeds to greater number of gaseous moles /molecules/more moles of gases on right/OWTTE;

ECF from (i) 4

(iii) at high pressure concentration increases/reaction rate faster; more frequent collisions; 2 [10]

37. less product is present at higher temperatures; Therefore the forward reaction is exothermic; 2 [2]

38. (i) (Kc =) (ignore units); 1

(ii) Increasing the pressure: Yield increases/equilibrium moves to the right/more ammonia; 4 gas molecules → 2/decrease in volume/fewer gas molecules on right hand side;

Increasing the temperature: Yield decreases/equilibrium moves to the left/less ammonia; Exothermic reaction/OWTTE; 4

(iii) Higher temperature increases rate; Lower pressure is less expensive/lower cost of operating at low pressure/reinforced pipes not needed; 2

Do not award a mark just for the word “compromise”.

(iv) 2.2 (dm3); 1 Penalize incorrect units.

(v) Fertilizers/increasing crop yields; Production of explosives for mining; 1 max

(vi) Fe/iron; Allow magnetite/iron oxide.

Claim is not valid since catalysts do not alter the yield/position of equilibrium/only increase the rate of reaction; 2

Topic 6 Questions

1. Which of the following is (are) important in determining whether a reaction occurs?

I. Energy of the molecules

II. Orientation of the molecules

A. I only

B. II only

C. Both I and II

D. Neither I nor II

2. Consider the reaction between solid CaCO3 and aqueous HCl. The reaction will be speeded up by an increase in which of the following conditions?

I. Concentration of the HCl

II. Size of the CaCO3 particles

III. Temperature

A. I only

B. I and III only

C. II and III only

D. I, II and III

3. Excess magnesium was added to a beaker of aqueous hydrochloric acid on a balance. A graph of the mass of the beaker and contents was plotted against time (line 1).

What change in the experiment could give line 2?

I. The same mass of magnesium but in smaller pieces

II. The same volume of a more concentrated solution of hydrochloric acid

III. A lower temperature

A. I only

B. II only

C. III only

D. None of the above

4. The rate of a reaction between two gases increases when the temperature is increased and a catalyst is added. Which statements are both correct for the effect of these changes on the reaction?

Increasing the temperature Adding a catalyst

A. Collision frequency increases Activation energy increases

B. Activation energy increases Activation energy does not change

C. Activation energy does not change Activation energy decreases

D. Activation energy increases Collision frequency increases

5. Which of the following is (are) altered when a liquid at its boiling point is converted to a gas at the same temperature?

I. The size of the molecules

II. The distance between the molecules

III. The average kinetic energy of the molecules

A. I only

B. II only

C. III only

D. I and II only

6. Based on the definition for rate of reaction, which units are used for a rate?

A. mol dm–3

B. mol time–1

C. dm time–1

D. mol dm–3 time–1

7. Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a catalyst?

Time

A. I only

B. III only

C. I and II only

D. II and III only

8. In the Haber process for the synthesis of ammonia, what effects does the catalyst have?

Rate of formation of NH3(g) Amount of NH3(g) formed A. Increases Increases B. Increases Decreases C. Increases No change D. No change Increases

9. Which statement is correct for a collision between reactant particles leading to a reaction?

A. Colliding particles must have different energy.

B. All reactant particles must have the same energy.

C. Colliding particles must have a kinetic energy higher than the activation energy.

D. Colliding particles must have the same velocity.

10. Which change of condition will decrease the rate of the reaction between excess zinc granules and dilute hydrochloric acid?

A. increasing the amount of zinc

B. increasing the concentration of the acid

C. pulverize the zinc granules into powder

D. decreasing the temperature

11. The table shows the concentrations of reactants and products during this reaction.

2A + B → C + 2D

[A] / mol dm–3 [B] / mol dm–3 [C] / mol dm–3 [D] / mol dm–3

at the start 6 3 0 0

after 1 min 4 2 1 2

The rate of reaction can be measured by reference to any reactant or product. Which rates are correct for this reaction?

I. rate = –2 mol dm–3 min–1 for A II. rate = –1 mol dm–3 min–1 for B III. rate = –1 mol dm–3 min–1 for C

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

12. In general, the rate of a reaction can be increased by all of the following except

A. increasing the temperature.

B. increasing the activation energy.

C. increasing the concentration of reactants.

D. increasing the surface area of the reactants.

13. At 25°C, 100 cm3 of 1.0 mol dm–3 hydrochloric acid is added to 3.5 g of magnesium carbonate. If the sample of magnesium carbonate is kept constant, which conditions will not increase the initial rate of reaction?

Volume of HCl / cm3 Concentration of HCl / mol dm–3 Temperature / °C

A. 200 1.0 25

B. 100 2.0 25

C. 100 1.0 35

D. 200 2.0 25

14. At 25°C, 100 cm3 of 1.0 mol dm–3 hydrochloric acid is added to 3.5 g of magnesium carbonate. If the sample of magnesium carbonate is kept constant, which conditions will not increase the initial rate of reaction?

Volume of HCl / cm3 Concentration of HCl / mol dm–3 Temperature / °C

A. 200 1.0 25

B. 100 2.0 25

C. 100 1.0 35

D. 200 2.0 25

15. Which statement is correct with regard to the catalysed and uncatalysed pathways for a given reaction?

A. The enthalpy change of the catalysed reaction is less than the enthalpy change for the uncatalysed reaction.

B. The enthalpy change of the catalysed reaction is greater than the enthalpy change for the uncatalysed reaction.

C. The enthalpy change of the catalysed reaction is equal to the enthalpy change for the uncatalysed reaction.

D. The activation energy of the catalysed reaction is greater than the activation energy for the uncatalysed reaction.

16. Which changes increase the rate of a chemical reaction?

I. Increase in the concentration of an aqueous solution

II. Increase in particle size of the same mass of a solid reactant

III. Increase in the temperature of the reaction mixture

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

17. Excess magnesium, was added to a beaker of aqueous hydrochloric acid. A graph of the mass of the beaker and contents was plotted against time (line 1).

What change in the experiment could give line 2?

A. The same mass of magnesium in smaller pieces

B. The same volume of a more concentrated solution of hydrochloric acid

C. A lower temperature

D. A more accurate instrument to measure the time

18. Which quantities in the enthalpy level diagram are altered by the use of a catalyst?

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

19. The graph below shows the volume of carbon dioxide gas produced against time when excess calcium carbonate is added to x cm3 of 2.0 mol dm–3 hydrochloric acid.

(i) Write a balanced equation for the reaction.

………………………………………………………………………………………….. (1)

(ii) State and explain the change in the rate of reaction with time. Outline how you would determine the rate of the reaction at a particular time.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (4)

(iii) Sketch the above graph on an answer sheet. On the same graph, draw the curves you would expect if:

I. the same volume (x cm3) of 1.0 mol dm–3 HCl is used.

II. double the volume (2x cm3) of 1.0 mol dm–3 HCl is used.

Label the curves and explain your answer in each case.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (5)

(Total 10 marks)

20. When excess lumps of magnesium carbonate are added to dilute hydrochloric acid the following reaction takes place.

MgCO3(s) + 2HCl(aq) → MgCl2(aq) + CO2(g) + H2O(l)

(a) Outline two ways in which the rate of this reaction could be studied. In each case sketch a graph to show how the value of the chosen variable would change with time.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(b) State and explain three ways in which the rate of this reaction could be increased.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (6)

(c) State and explain whether the total volume of carbon dioxide gas produced would increase, decrease or stay the same if

(i) more lumps of magnesium carbonate were used.

……………………………………………………………………………………

……………………………………………………………………………………

…………………………………………………………………………………… (2)

(ii) the experiments were carried out at a higher temperature.

……………………………………………………………………………………

……………………………………………………………………………………

…………………………………………………………………………………… (2)

(Total 14 marks)

21. Carbon dioxide gas in the atmosphere reacts slightly with rainwater as shown below.

CO2(g) + H2O(l) H+(aq) + HCO3–(aq)

(i) State the meaning of the sign.

…………………………………………………………………………………………… (1)

(ii) Predict the effect, if any, of the presence of a catalyst on the acidity of rainwater. Give a reason for your answer.

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(iii) Use Le Chatelier’s principle to predict the effect of the addition of a small quantity of an alkali on the acidity of rainwater. Explain what effect, if any, this would have on the equilibrium constant, Kc.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 6 marks)

22. Excess 0.100 mol dm–3 nitric acid is added to a certain mass of powdered calcium carbonate at 20°C. The rate of reaction is monitored by measuring the change in mass over time due to the loss of carbon dioxide.

(a) Define the term rate of reaction.

....................................................................................................................................

.................................................................................................................................... (1)

(b) Explain why the mass loss remains constant after a certain time.

....................................................................................................................................

.................................................................................................................................... (1)

(c) Draw a line on the graph above, to show what the graph would look like if the same mass of calcium carbonate in larger pieces were reacted with excess 0.100 mol dm–3 nitric acid.

(1)

(d) Explain in terms of the collision theory what would happen to the rate if the reaction was conducted at 50°C.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (3)

(e) Determine the rate of formation of carbon dioxide when the nitric acid reacts at a rate of 2.00⋅10–3 mol cm–3 s–1.

....................................................................................................................................

.................................................................................................................................... (1)

(Total 7 marks)

23. (i) Draw a graph that shows the distribution of molecular energies in a sample of a gas at two different temperatures, T1 and T2, such that T2 is greater than T1.

(2)

(ii) Define the term activation energy. (1)

(iii) State and explain the effect of a catalyst on the rate of an endothermic reaction. (2)

(Total 5 marks)

24. (i) Magnesium is added to a solution of hydrochloric acid. Sketch a graph of acid concentration on the y-axis against time on the x-axis to illustrate the progress of the reaction.

(1)

(ii) Describe how the slope of the line changes with time. (1)

(iii) Use the collision theory to state and explain the effect of decreasing concentration on the rate of the reaction.

(2) (Total 4 marks)

25. The reaction between ammonium chloride and sodium nitrite in aqueous solution can be represented by the following equation.

NH4Cl(aq) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl(aq)

The graph below shows the volume of nitrogen gas produced at 30 second intervals from a mixture of ammonium chloride and sodium nitrite in aqueous solution at 20°C.

(a) (i) State how the rate of formation of nitrogen changes with time. Explain your answer in terms of collision theory.

............................................................................................................................

............................................................................................................................

............................................................................................................................

............................................................................................................................

............................................................................................................................ (2)

(ii) Explain why the volume eventually remains constant.

............................................................................................................................

............................................................................................................................ (1)

(b) (i) State how the rate of formation of nitrogen would change if the temperature were increased from 20°C to 40°C.

............................................................................................................................

............................................................................................................................ (1)

(ii) State two reasons for the change described in (b)(i) and explain which of the two is more important in causing the change.

............................................................................................................................

............................................................................................................................

............................................................................................................................

............................................................................................................................

............................................................................................................................ (3)

(iii) The reaction between solid ammonium chloride and aqueous sodium nitrite can be represented by the following equation.

NH4Cl(s) + NaNO2(aq) → N2(g) + 2H2O(l) + NaCl(aq)

State and explain how the rate of formation of nitrogen would change if the same amount of ammonium chloride was used as large lumps instead of as a fine powder.

............................................................................................................................

............................................................................................................................

............................................................................................................................ (2)

(Total 9 marks)

26. (a) Define the term average bond enthalpy, illustrating your answer with an equation for methane, CH4.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(b) The equation for the reaction between methane and chlorine is

CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)

Use the values from Table 10 of the Data Booklet to calculate the enthalpy change for this reaction.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(c) Explain why no reaction takes place between methane and chlorine at room temperature unless the reactants are sparked, exposed to UV light or heated.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(d) Draw an enthalpy level diagram for this reaction.

(2)

(Total 10 marks)

27. (a) Identify two features of colliding molecules that react together in the gas phase.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(b) For many reactions, the rate approximately doubles for a 10°C rise in temperature. State two reasons for this increase and identify which of the two is the more important.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(Total 5 marks)

28. (a) Define the term rate of reaction.

...................................................................................................................................

................................................................................................................................... (1)

(b) The reaction between gases C and D is slow at room temperature.

(i) Suggest two reasons why the reaction is slow at room temperature.

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(ii) A relatively small increase in temperature causes a relatively large increase in the rate of this reaction. State two reasons for this.

.........................................................................................................................

......................................................................................................................... (2)

(iii) Suggest two ways of increasing the rate of reaction between C and D other than increasing temperature.

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(Total 7 marks) 29. The reaction between two substances A and B

A + B → C + D

has the following rate expression:

rate = k [B]

Draw the graphical representation of:

(Total 3 marks

IB Chemistry – SL Topic 6 Answers

1. C

2. B

3. B

4. C

5. B

6. D

7. C

8. C

9. C

10. D

11. A

12. B

13. A

14. A

15. C

16. B

17. B

18. A

19. (i) CaCO3(s) + 2H+(aq) → Ca2+(aq) + H2O(l) + CO2(g) 1 States not required, accept molecular equation.

(ii) rate decreases with time; as concentration decreases so fewer (successful) collisions; draw tangent to the curve at time t; rate = slope or gradient; 4

(iii)

I. (less CO2 because) amount of HCl is limiting and half the orginal/OWTTE;

II. (same amount of CO2 because) amount of HCl is the same; curve less steep because less frequent (accept fewer) collisions 5

20. (a) measure volume of carbon dioxide/CO2/gas produced/measure pH; 4

starts at origin and levels off

measure mass of chemicals/apparatus;

starts high and decreases Graph should show increase as reaction progresses (as HCl is consumed).

(b) Method 1 use powdered MgCO3/OWTTE; particles collide more frequently/increased surface area/OWTTE;

Method 2 increase (reaction) temperature/heat/warm; more of the collisions are successful/more

particles with E > Ea/OWTTE;

Method 3 increase acid concentration; more frequent (reactant) collisions;

Method 4 add catalyst; lowers activation energy/Ea/OWTTE; 6 max Award [2] each for any three methods

(c) (i) stays the same; MgCO3 was already in excess; 2

(ii) stays the same; same quantities of reactants used; 2 [14]

21. (i) reversible reaction/reaction may proceed in either direction (depending on reaction conditions) equilibrium/dynamic equilibrium; 1

(ii) no effect; catalyst will speed up both forward and reverse reactions (equally)/ increase the rate at which equilibrium is achieved; 2

(iii) acidity: no effect; equilibrium shifts to the right; Kc: no change; 3 [6]

22. (a) change of concentration/mass/amount/volume/of a reactant/product with time; 1 Do not accept “substance”.

(b) all the CaCO3(s) has been consumed/no further CO2(g) is produced/reaction is complete; 1 Do not accept reaction has stopped or all reactants used up.

(c) line on graph should be initially less steep/a smaller gradient and should plateau at the same mass loss; 1

(d) there are more particles with KE greater than or equal to Ea; collisions more frequent/more collisions per unit time/more successful/forceful collisions per unit

time; the rate increases; 3

(e) 1.00×10−3 (mol cm−3 s−1) 1 Ignore units even if wrong. Apply −1(sf).

[7]

23. (i)

T2 peak lower/T1 higher; T2 peak at higher energies/T1 curve at lower energies; 2

Maximum [1] if axes not labeled correctly

(ii) minimum energy required to react/energy difference between reactants and transition state; 1

(iii) makes the reaction go faster; because it lowers the activation energy/Ea; 2 [5]

24. (i) a curve showing concentration decreases with time; 1

No penalty if curve reaches x axis Do not accept a straight line

(ii) slope decreases; 1

(iii) rate decreases; fewer collisions per unit time; 2 [4]

25. (a) (i) it is decreasing; less frequent collisions/fewer collisions per second or (unit) time; 2

(ii) reactant(s) used up/reaction is complete; 1 Do not accept reaction reaches equilibrium.

(b) (i) it would increase; 1 Accept a quantitative answer such as “doubles”.

(ii) more frequent collisions; collisions or molecules have more energy (OWTTE); more molecules with energy ≥ Ea; 3

(iii) rate would be lower; smaller surface area; 2 [9]

26. (a) energy for the conversion of a gaseous molecule into (gaseous) atoms; (average values) obtained from a number of similar bonds/compounds/OWTTE; CH4(g) → C(g) + 4H(g); 3

State symbols needed.

(b) (bond breaking) = 1890/654; (bond formation) = 2005/769; enthalpy = –115(kJ mol–1) 3 Allow ECF from bond breaking and forming. Award [3] for correct final answer. Penalize [1] for correct answer with wrong sign.

(c) molecules have insufficient energy to react (at room temperature)/ wrong collision geometry/unsuccessful collisions; extra energy needed to overcome the activation energy/Ea for

the reaction; 2

(d)

exothermic shown;

activation energy/Ea shown; 2 Allow ECF from (b).

[10]

27. (a) molecules must have sufficient/minimum energy/energy ≥ activation energy; appropriate collision geometry/correct orientation; 2

(b) increased frequency of collisions/collisions more likely; Not just “more collisions”, there must be a reference to time.

increased proportion of molecules with sufficient energy to react/E ≥ Ea; Not “activation energy is reduced”.

Proportion of molecules with E ≥ Ea is more important; (dependent on correct second marking point); 3

28. (a) increase in product concentration per unit time/decrease in reactant concentration per unit time; 1 Accept change instead of increase or decrease.

(b) (i) high activation energy/not enough molecules have Ea/OWTTE; incorrect collision geometry/OWTTE; infrequent collisions; 2

Award [1] for any two reasons.

(ii) more energetic collisions/more molecules have (energy ≥) Ea; more frequent collisions/collide more often; 2

(iii) add a catalyst; increase the (total) pressure/decrease the volume of the container; increase the concentration of C (or D); 2

Do not accept surface area. Award [1] for any two.

[7] 29. ([A] against time) - straight line with negative gradient;

Accept any decreasing curve ([B] against time) - decreasing curve;

Award [1] unless half - lives clearly not constant (rate against [A]) - any horizontal straight line; (rate against [B]) - straight line through origin;

Award [3] for all four correct, award [2] for any three correct and [1] for any two correct.

IB Chemistry – SL

Topic 4 Questions

1. What is the formula for the compound formed by calcium and nitrogen?

A. CaN

B. Ca2N

C. Ca2N3

D. Ca3N2 (Total 1 mark)

2. Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be present in the compound formed when X and Y react together?

A. X+ and Y–

B. X 2+ and Y–

C. X+ and Y2–

D. X2– and Y+ (Total 1 mark)

3. Based on electronegativity values, which bond is the most polar?

A. B―C

B. C―O

C. N―O

D. O―F (Total 1 mark)

4. What is the Lewis (electron dot) structure for sulfur dioxide?

A.

B.

C.

D. (Total 1 mark

5. Which substance is most soluble in water (in mol dm–3) at 298 K?

A. CH3CH3

B. CH3OCH3

C. CH3CH2OH

D. CH3CH2CH2CH2OH (Total 1 mark)

6. According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in the order

A. lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

B. bond pair-bond pair > lone pair-bond pair > lone pair-lone pair.

C. lone pair-lone pair > bond pair-bond pair > bond pair-lone pair.

D. bond pair-bond pair > lone pair-lone pair > lone pair-bond pair. (Total 1 mark)

7. Which molecule is linear?

A. SO2

B. CO2

C. H2S

D. Cl2O (Total 1 mark)

8. Why is the boiling point of PH3 lower than that of NH3?

A. PH3 is non-polar whereas NH3 is polar.

B. PH3 is not hydrogen bonded whereas NH3 is hydrogen bonded.

C. Van der Waals’ forces are weaker in PH3 than in NH3.

D. The molar mass of PH3 is greater than that of NH3. (Total 1 mark)

9. Which molecule is non-polar?

A. H2CO

B. SO3

C. NF3

D. CHCl3

(Total 1 mark)

10. What happens when sodium and oxygen combine together?

A. Each sodium atom gains one electron.

B. Each sodium atom loses one electron.

C. Each oxygen atom gains one electron.

D. Each oxygen atom loses one electron. (Total 1 mark)

11. Which statement is correct about two elements whose atoms form a covalent bond with each other?

A. The elements are metals.

B. The elements are non-metals.

C. The elements have very low electronegativity values.

D. The elements have very different electronegativity values. (Total 1 mark)

12. Which substance has the lowest electrical conductivity?

A. Cu(s)

B. Hg(l)

C. H2(g)

D. LiOH(aq) (Total 1 mark)

13. When the following bond types are listed in decreasing order of strength (strongest first), what is the correct order?

A. covalent > hydrogen > van der Waals’

B. covalent > van der Waals’ > hydrogen

C. hydrogen > covalent > van der Waals’

D. van der Waals’ > hydrogen > covalent (Total 1 mark)

14. Which statement is true for most ionic compounds?

A. They contain elements of similar electronegativity.

B. They conduct electricity in the solid state.

C. They are coloured.

D. They have high melting and boiling points. (Total 1 mark)

15. What is the valence shell electron pair repulsion (VSEPR) theory used to predict?

A. The energy levels in an atom

B. The shapes of molecules and ions

C. The electronegativities of elements

D. The type of bonding in compounds (Total 1 mark)

16. Which fluoride is the most ionic?

A. NaF

B. CsF

C. MgF2

D. BaF2 (Total 1 mark)

17. Which substance is most similar in shape to NH3?

A. GaI3

B. BF3

C. FeCl3

D. PBr3 (Total 1 mark)

18. Which statement is a correct description of electron loss in this reaction?

2Al + 3S → Al2S3

A. Each aluminium atom loses two electrons.

B. Each aluminium atom loses three electrons.

C. Each sulfur atom loses two electrons.

D. Each sulfur atom loses three electrons. (Total 1 mark)

19. Which molecule has the smallest bond angle?

A. CO2

B. NH3

C. CH4

D. C2H4 (Total 1 mark)

20. In which substance is hydrogen bonding present?

A. CH4

B. CH2F2

C. CH3CHO

D. CH3OH (Total 1 mark)

21. Which is a correct description of metallic bonding?

A. Positively charged metal ions are attracted to negatively charged ions.

B. Negatively charged metal ions are attracted to positively charged metal ions.

C. Positively charged metal ions are attracted to delocalized electrons.

D. Negatively charged metal ions are attracted to delocalized electrons. (Total 1 mark)

22. What intermolecular forces are present in gaseous hydrogen?

A. Hydrogen bonds

B. Covalent bonds

C. Dipole-dipole attractions

D. Van der Waals’ forces (Total 1 mark)

23. Which molecule is polar?

A. CO2

B. PF3

C. CH4

D. BF3 (Total 1 mark)

24. What are responsible for the high electrical conductivity of metals?

A. Delocalized positive ions

B. Delocalized valence electrons

C. Delocalized atoms

D. Delocalized negative ions (Total 1 mark)

25. Which compound has the least covalent character?

A. SiO2

B. Na2O

C. MgCl2

D. CsF (Total 1 mark)

26. When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the correct order?

A. C2H6, C2H2, C2H4

B. C2H4, C2H2, C2H6

C. C2H2, C2H4, C2H6

D. C2H4, C2H6, C2H2 (Total 1 mark)

27. Which compound contains both ionic and covalent bonds?

A. MgCl2

B. HCl

C. H2CO

D. NH4Cl (Total 1 mark)

28. When the species BF2+, BF3 and BF4

– are arranged in order of increasing F−B−F bond angle, what is the correct order?

A. BF3, BF4–, BF2

+

B. BF4–, BF3, BF2

+

C. BF2+, BF4

–, BF3

D. BF2+, BF3, BF4

– (Total 1 mark)

29. Which species has a trigonal planar shape?

A. CO32–

B. SO32–

C. NF3

D. PCl3

30. When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the correct order?

A. C2H6, C2H2, C2H4

B. C2H4, C2H2, C2H6

C. C2H2, C2H4, C2H6

D. C2H4, C2H6, C2H2 (Total 1 mark)

31. What is the formula for an ionic compound formed between an element, X, from group 2 and an element, Y, from group 6?

A. XY

B. X2Y

C. XY2

D. X2Y6

(Total 1 mark)

32. In the molecules N2H4, N2H2, and N2, the nitrogen atoms are linked by single, double and triple bonds, respectively. When these molecules are arranged in increasing order of the lengths of their nitrogen to nitrogen bonds (shortest bond first) which order is correct?

A. N2H4, N2, N2H2

B. N2H4, N2H2, N2

C. N2H2, N2, N2H4

D. N2, N2H2, N2H4 (Total 1 mark)

33. The compounds listed have very similar molar masses. Which has the strongest intermolecular forces?

A. CH3CHO

B. CH3CH2OH

C. CH3CH2F

D. CH3CH2CH3 (Total 1 mark)

34. What is the shape of the CO32– ion and the approximate O–C–O bond angle?

A. Linear, 180°

B. Trigonal planar, 90°

C. Trigonal planar, 120°

D. Pyramidal, 109° (Total 1 mark)

35. Which combination of ⊗Hvaporization and boiling point is the result of strong intermolecular forces?

⊗Hvaporization Boiling Point

A. large high

B. large low

C. small low

D. small high (Total 1 mark)

36. What is the formula of the compound formed when aluminium reacts with oxygen?

A. Al3O2

B. Al2O3

C. AlO2

D. AlO3 (Total 1 mark)

37. Which statement is true for compounds containing only covalent bonds?

A. They are held together by electrostatic forces of attraction between oppositely charged ions.

B. They are made up of metal elements only.

C. They are made up of a metal from the far left of the periodic table and a non-metal from the far right of the periodic table.

D. They are made up of non-metal elements only. (Total 1 mark)

38. How many electrons are used in the carbon-carbon bond in C2H2?

A. 4

B. 6

C. 10

D. 12 (Total 1 mark)

39. Which compound has the highest boiling point?

A. CH3CH2CH3

B. CH3CH2OH

C. CH3OCH3

D. CH3CHO (Total 1 mark)

40. What type of solid materials are typically hard, have high melting points and poor electrical conductivities?

I. Ionic II. Metallic III. Covalent-network

A. I and II only

B. I and III only

C. II and III only


41. The boiling points of the hydrides of the group 6 elements are shown below.

(i) Explain the trend in boiling points from H2S to H2Te.

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(ii) Explain why the boiling point of water is higher than would be expected from the group trend.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(Total 4 marks)

42. (i) State the shape of the electron distribution around the oxygen atom in the water molecule and state the shape of the molecule.

………………………………………………………………………………………….

…………………………………………………………………………………………. (2)

(ii) State and explain the value of the HOH bond angle.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(Total 4 marks)

43. Explain why the bonds in silicon tetrachloride, SiCl4, are polar, but the molecule is not.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (Total 2 marks)

44. The diagrams below represent the structures of iodine, sodium and sodium iodide.

A B C

(a) (i) Identify which of the structures (A, B and C) correspond to iodine, sodium and sodium iodide.

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (1)

(ii) State the type of bonding in each structure.

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (3)

(b) (i) Sodium and sodium iodide can both conduct electricity when molten, but only sodium can conduct electricity when solid. Explain this difference in conductivity in terms of the structures of sodium and sodium iodide.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (4)

(ii) Explain the high volatility of iodine compared to sodium and sodium iodide.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(Total 10 marks)

45. (i) Draw Lewis (electron dot) structures for CO2 and H2S showing all valence electrons.

(2)

(ii) State the shape of each molecule and explain your answer in terms of VSEPR theory.

CO2 .............................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

H2S .............................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (4)

(iii) State and explain whether each molecule is polar or non-polar.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 8 marks)

46. Identify the strongest type of intermolecular force in each of the following compounds.

CH3Cl ...................................................................................................................................

CH4 .......................................................................................................................................

CH3OH ................................................................................................................................. (Total 3 marks)

47. (a) An important compound of nitrogen is ammonia, NH3. The chemistry of ammonia is influenced by its polarity and its ability to form hydrogen bonds. Polarity can be explained in terms of electronegativity.

(i) Explain the term electronegativity.

……………………………………………………………………………………

…………………………………………………………………………………… (2)

(ii) Draw a diagram to show hydrogen bonding between two molecules of NH3. The diagram should include any dipoles and/or lone pairs of electrons

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

…………………………………………………………………………………… (3)

(iii) State the H–N–H bond angle in an ammonia molecule.

……………………………………………………………………………………… (1)

(iv) Explain why the ammonia molecule is polar.

……………………………………………………………………………………

……………………………………………………………………………………

…………………………………………………………………………………… (1)

(b) Ammonia reacts with hydrogen ions forming ammonium ions, NH4+.

(i) State the H–N–H bond angle in an ammonium ion.

…………………………………………………………………………………… (1)

(ii) Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond angle of NH4

+; referring to both species in your answer.

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

…………………………………………………………………………………… (3)

(Total 11 marks)

48. State the type of bonding in the compound SiCl4. Draw the Lewis structure for this compound.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

(Total 3 marks)

49. Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

(Total 3 marks)

50. (i) Use the VSEPR theory to predict and explain the shape and the bond angle of each of the molecules SCl2 and C2Cl2

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (6)

(ii) Deduce whether or not each molecule is polar, giving a reason for your answer.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 9 marks)

51. Draw a Lewis structure of a water molecule, name the shape of the molecule and state and explain why the bond angle is less than the bond angle in a tetrahedral molecule such as methane.

(Total 4 marks)

52. Predict and explain the order of the melting point for propanol, butane and propanone with reference to their intermolecular forces.

(Total 4 marks)

53. The elements sodium, aluminium, silicon, phosphorus and sulfur are in period 3 of the periodic table.

Describe the metallic bonding present in aluminium and explain why aluminium has a higher melting point than sodium.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 3 marks)

54. Draw the Lewis structure of NCl3. Predict, giving a reason, the Cl – N – Cl bond angle in NCl3. (Total 3 marks)

55. Arrange the following in decreasing order of bond angle (largest one first), and explain your reasoning.

NH2–, NH3, NH4

+ (Total 5 marks)

56. (i) Outline the principles of the valence shell electron pair repulsion (VSEPR) theory. (3)

(ii) Use the VSEPR theory to deduce the shape of H3O+ and C2H4. For each species, draw the Lewis structure, name the shape, and state the value of the bond angle(s).

(6)

(iii) Predict and explain whether each species is polar. (2)

(iv) Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or N-H would be most polar.


57. Predict and explain which of the following compounds consist of molecules: NaCl, BF3, CaCl2, N2O, P4O6, FeS and CBr4.

(Total 2 marks)

58. Diamond, graphite and C60 fullerene are three allotropes of carbon.

(i) Describe the structure of each allotrope. (3)

(ii) Compare the bonding in diamond and graphite. (2)

(Total 5 marks)

59. State two physical properties associated with metals and explain them at the atomic level. (Total 4 marks)

60. (a) Draw the Lewis structure of methanoic acid, HCOOH.

(1)

(b) In methanoic acid, predict the bond angle around the (2)

(i) carbon atom. .....................................................................................................

(ii) oxygen atom bonded to the hydrogen atom. ...................................................

(c) State and explain the relationship between the length and strength of the bonds between the carbon atom and the two oxygen atoms in methanoic acid.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (3)

(Total 6 marks)

Topic 4 Answers

1. D [1]

2. B [1]

3. B

[1] 4. D

[1] 5. C

[1]

6. A [1]

7. B [1]

8. B [1]

9. B [1]

10. B

[1] 11. B

[1] 12. C

[1] 13. A

[1] 14. D

[1] 15. B

[1] 16. B

[1] 17. D

[1]

18. B [1]

19. B [1]

20. D [1]

21. C [1]

22. D [1]

23. B [1]

24. B

[1]

25. D [1]

26. C [1]

27. D [1]

28. B [1]

29. A [1]

30. C [1]

31. A [1]

32. D [1]

33. B [1]

34. C [1]

35. A [1]

36. B [1]

37. D [1]

38. A [1]

39. B [1]

40. B [1]

41. (i) as molecules become larger/heavier/have higher Mr values/ number of electrons increases; van der Waals’/London/ dispersion forces increase; 2

(ii) hydrogen bonding between molecules in H2O; this bonding is stronger (than van der Waals’ forces); 2

Must be an implied comparison with (i) [4]

42. (i) tetrahedral (accept correct 3-D diagram); bent/V-shape/angular (accept suitable diagram); 2

(ii) 105° (accept 103 – 106°); lone pairs repel each other more than bonding pairs; 2 Do not accept repulsion of atoms.

[4]

43. bonds are polar as Cl more electronegative than Si; Allow “electronegativities are different”

molecule is symmetrical, hence polar effects cancel out/OWTTE; 2 [2]

44. (a) (i) A – sodium iodide, B – sodium, C – iodine (three correct [1]); 1 Accept correct formulas.

(ii) A – ionic bonding; B – metallic bonding; C – van der Waals’ forces (and covalent bonding); 3

(b) (i) (for Na) (lattice of) positive ions/atoms; delocalized/free electrons/sea of electrons; (for NaI) oppositely charged ions/positive and negative ions; free to move (only) in molten state; 4

(ii) forces between I2 molecules are weak; ionic/metallic bonding strong(er); 2 [10]

45. (i) 2 Accept dots, crosses, a combination of dots and crosses or a line to represent a pair of electrons.

(ii) CO2 is linear; two charge centres or bonds and no lone pairs (around C); H2S is bent/v-shaped/angular; two bond pairs, two lone pairs (around S); 4

(iii) CO2 is non-polar, H2S is polar; bond polarities cancel CO2 but not in H2S; 2 [8]

46. CH3Cl – dipole-dipole attractions; CH4 – van der Waals’/dispersion/London forces; CH3OH – hydrogen bond; 3

[3]

47. (a) (i) (relative) measure of an atom’s attraction for electrons; in a bond; 2

(ii)

Suitable diagram indicating dipoles; lone pairs of electrons; hydrogen bonding; 3

(iii) 107°; 1

Accept answer in range 107 to 109° .

(iv) molecule is asymmetrical/OWTTE; 1

(b) (i) 109.5°; 1

(ii) NH4+ has four bonding pairs (around central atom so is a regular tetrahedron); NH3 has

three bonding pairs (of electrons) and one non-bonding pair; non-bonding pairs (of electrons) exert a greater repulsive force; 3

Accept suitable diagrams. [11]

48. Si—Cl bonds are covalent; 3

Accept lines for electron pairs. Award [1] for covalent bonds and [1] for lone pairs.

[3] 49. find number of electron pairs/charge centres in (valence shell of) central atom; electron pairs/charge centres (in valence shell) of central atom repel each other; to positions of minimum energy/repulsion/maximum stability; pairs forming a double or triple bond act as a single bond; non-bonding pairs repel more than bonding pairs/OWTTE; 3 max

Do not accept repulsion between bonds or atoms. Award [1] each for any three points.

[3]

50. (i) SCl2 two bonding pairs, two non-bonding pairs; angular/bent/non-linear/V-shaped; Both these marks can be scored from a diagram. 90° < angle < 107°;

C2Cl2 two charge centres around each C; linear; Both these marks can be scored from a diagram.

angle = 180°; 6

(ii) SCl2 is polar; C2Cl2 is non-polar; 3 No net dipole movement for C2Cl2 but angular SCl2 has a resultant dipole / OWTTE;

Mark can be scored from a diagram. Allow ECF based on the answers given to (i).

[9]

51.

Allow a combination of dots, crosses or lines.

bent/V shaped/angular 104.5°;

Accept answers in range 104° to 106°.

repulsion of the two non-bonding pairs of electrons forces bond angle to be smaller/non-bonding pairs repel more than bonding pairs; 4

[4]

52. butane < propanone < propanol;

butane has van der Waals’ forces; Accept vdW, dispersion or London forces or attractions between temporary dipoles.

propanone has dipole-dipole attractions;

propanol has (the stronger) H-bonding; 4 [4]

53. delocalized electrons; (attracted) to positive ions; more delocalized/mobile/outer shell electrons/higher ionic charge; 3

[3]

54. All electrons must be shown. Accept molecular structures using lines to represent bonding and lone electron pairs.

bond angle: 107°−109°

greater repulsion between lone pair and bonding pairs/OWTTE; 3 NOT between electron pairs and atoms. Award [1 max] if lone pair missed on nitrogen, ECF for bond angle of 120°.

[3]

55. NH4+ > NH3 > NH2

–; NH4

+ has four bonded electron pairs (and no lone electron pairs); NH3 has three bonded electron pairs and one electron lone pair; NH2

– has two bonded electron pairs and two electron lone pairs; Accept correct Lewis structures with lone electron pairs clearly shown.

lone pair-lone pair > lone pair-bonded pair > bonded pair-bonded pair/ lone pairs of electrons repel more than bonding pairs of electrons/OWTTE; 5

Do not accept repulsion between atoms. [5]

56. (i) Find number of electron pairs/charge centres in (valence shell of) central atom; electron pairs/charge centres (in valence shell) of central atom repel each other;

Any one of the following: to positions of minimum energy/repulsion/maximum stability; pairs forming a double or triple bond act as a single bond; non-bonding pairs repel more than bonding pairs/OWTTE; 3 max

Do not accept repulsion between bonds or atoms.

(ii) 6

Species Lewis (electron-dot) structure Shape Bond angle(s)

H3O+ Trigonal/triangular pyramidal; Allow values in the range 106° to 109.5°;

C2H4 Trigonal/triangular planar; Allow values of approximately 120°;

Accept crosses and dots for electrons in Lewis structures also. As the Lewis structures were asked for, and not 3D representations, do not penalize incorrectly drawn geometries. Do not accept structure of hydronium cation without lone pair on oxygen. No penalty for missing charge.

(iii) H3O+: is polar and explanation either using a diagram or in words, involving the net dipole moment;

e.g. the three individual O-H bond dipole moments add as vectors to give a net dipole moment.

C2H4: is non-polar and explanation either using a diagram or in words, involving no net dipole moment; 2

e.g. the vector sum of the individual bond dipole moments is zero. For simple answers such as bond polarities do not cancel for H3O+ and do cancel for C2H4, Award [1], only for the last two marking points.

(iv) O-H is most polar; O-H has greatest difference between electronegativities/calculation showing values of 1.4, 0.5 and 0.9 respectively; 2

[13]

57. BF3, N2O, P4O6 and CBr4; Non-metals only/small difference in electronegativity values of the elements; 2 [2]

58. (i) 3 Allotrope Structure

Diamond 3D array/network involving tetrahedral carbons/each carbon atom joined to four others;

Graphite layer structure involving trigonal (triangular) planar carbons/with each carbon atom joined to three others/with hexagonal (six-membered) rings of carbon atoms;

C60 fullerene truncated icosahedrons; Accept carbon atoms form a ‘ball’ with 32 faces, of which 12 are pentagons and 20 are hexagons, exactly like a soccer ball. Do not accept soccer ball alone.

(ii) Diamond: covalent bonds (only); Graphite: covalent bonds and the separated layers held together by (weak) London/van der Waals’/dispersion forces; 2

[5]

59. Electrical conductivity: Bonding electrons are delocalised; Current flow occurs without displacement of

atoms within the metal/ able to flow within the metal;

Malleability: Can be hammered into thin sheets; atoms capable of slipping with respect to one another; 4 [4]

60. (a)

1 No mark without lone electron pairs. Correct shape not necessary. Do not award mark if dots/crosses and bond lines are shown. Accept lone pairs represented as straight lines.

(b) O − C − O = 120°/H − C − O = 120°; C − O − H = 109°/<109°; 2 No mark for 109.5° Accept answer in range 100–109°

(c) length: C = O < C − O; strength: C = O > C – O; greater number of electrons between nuclei pull atoms together and require greater energy to break;

Or

double bonds are shorter/single bonds are longer; double bonds are stronger/single bonds are weaker; 3

Accept stronger attraction between nuclei and (bonding) electrons. [6]

Topic 3 Questions

1. Which pair of elements reacts most readily?

A. Li + Br2

B. Li + Cl2

C. K + Br2

D. K + Cl2 (Total 1 mark)

2. Which of the following properties of the halogens increase from F to I?

I. Atomic radius

II. Melting point

III. Electronegativity

A. I only

B. I and II only

C. I and III only

D. I, II and III

(Total 1 mark)

3. Which pair would react together most vigorously?

A. Li and Cl2

B. Li and Br2

C. K and Cl2

D. K and Br2 (Total 1 mark)

4. For which element are the group number and the period number the same?

A. Li

B. Be

C. B

D. Mg (Total 1 mark)

5. Which of the physical properties below decrease with increasing atomic number for both the alkali metals and the halogens?

I. Atomic radius

II. Ionization energy

III. Melting point

A. I only

B. II only

C. III only

D. I and III only (Total 1 mark)

6. Rubidium is an element in the same group of the periodic table as lithium and sodium. It is likely to be a metal which has a

A. high melting point and reacts slowly with water.

B. high melting point and reacts vigorously with water.

C. low melting point and reacts vigorously with water.

D. low melting point and reacts slowly with water.

(Total 1 mark)

7. When the following species are arranged in order of increasing radius, what is the correct order?

A. Cl–, Ar, K+

B. K+, Ar , Cl–

C. Cl–, K+, Ar

D. Ar, Cl–, K+ (Total 1 mark)

8. What increases in equal steps of one from left to right in the periodic table for the elements lithium to neon?

A. the number of occupied electron energy levels

B. the number of neutrons in the most common isotope

C. the number of electrons in the atom

D. the atomic mass (Total 1 mark)

9. Which property decreases down group 7 in the periodic table?

A. atomic radius

B. electronegativity

C. ionic radius

D. melting point (Total 1 mark)

10. Which properties are typical of most non-metals in period 3 (Na to Ar)?

I. They form ions by gaining one or more electrons.

II. They are poor conductors of heat and electricity.

III. They have high melting points.

A. I and II only

B. I and III only

C. II and III only


11. A potassium atom has a larger atomic radius than a sodium atom. Which statement about potassium correctly explains this difference?

A. It has a larger nuclear charge.

B. It has a lower electronegativity.

C. It has more energy levels occupied by electrons.

D. It has a lower ionization energy. (Total 1 mark)

12. Which factors lead to an element having a low value of first ionization energy?

I. large atomic radius

II. high number of occupied energy levels

III. high nuclear charge

A. I and II only

B. I and III only

C. II and III only


13. Which statement about electronegativity is correct?

A. Electronegativity decreases across a period.

B. Electronegativity increases down a group.

C. Metals generally have lower electronegativity values than non-metals.

D. Noble gases have the highest electronegativity values. (Total 1 mark)

14. Which statement is correct for a periodic trend?

A. Ionization energy increases from Li to Cs.

B. Melting point increases from Li to Cs.

C. Ionization energy increases from F to I.

D. Melting point increases from F to I. (Total 1 mark)

15. Which compound of an element in period 3 reacts with water to form a solution with a pH greater than 7?

A. SiO2

B. SiCl4

C. NaCl

D. Na2O (Total 1 mark)

16. Which equation represents the first ionization energy of fluorine?

A. F(g) + e– → F–(g)

B. F–(g) → F(g) + e–

C. F+(g) → F(g) + e–

D. F(g) → F+(g) + e– (Total 1 mark)

17. Which statement is correct for the halogen group?

A. Halide ions are all reducing agents, with iodide ions being the weakest.

B. Halogens are all oxidizing agents, with chlorine being the strongest.

C. Chloride ions can be oxidized to chlorine by bromine.

D. Iodide ions can be oxidized to iodine by chlorine. (Total 1 mark)

18. Which of the following statements are correct?

I. The melting points decrease from Li → Cs for the alkali metals. II. The melting points increase from F → I for the halogens. III. The melting points decrease from Na → Ar for the period 3 elements.

A. I and II only

B. I and III only

C. II and III only


19. Which element is a transition metal?

A. Ca

B. Cr

C. Ge

D. Se (Total 1 mark)

20. When Na, K, and Mg are arranged in increasing order of atomic radius (smallest first), which order is correct?

A. Na, K, Mg

B. Na, Mg, K

C. K, Mg, Na

D. Mg, Na, K (Total 1 mark)

21. Which oxides produce an acidic solution when added to water?

I. SiO2

II. P4O6

III. SO2

A. I and II only

B. I and III only

C. II and III only


22. Which series is arranged in order of increasing radius?

A. Ca2+ < Cl– < K+

B. K+ < Ca2+ < Cl–

C. Ca2+ < K+ < Cl–

D. Cl– < K+ < Ca2+ (Total 1 mark)

23. Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character.

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................ (Total 4 marks)

24. Table 6 of the Data Booklet lists melting points of the elements. Explain the trend in the melting points of the alkali metals, halogens and period 3 elements.

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................ (Total 8 marks)

25. (i) Explain how the first ionization energy of K compares with that of Na and Ar.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(ii) Explain the difference between the first ionization energies of Na and Mg.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (4)

(iii) Suggest why much more energy is needed to remove an electron from Na+ than from Mg+.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (1)

(Total 8 marks)

26. Nitrogen is found in period 2 and group 5 of the periodic table.

(i) Distinguish between the terms period and group.

……………………………………………………………………………………………

…………………………………………………………………………………………… (1)

(ii) State the electron arrangement of nitrogen and explain why it is found in period 2 and group 5 of the periodic table.

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 4 marks)

27. Table 8 of the Data Booklet gives the atomic and ionic radii of elements. State and explain the difference between

(i) the atomic radius of nitrogen and oxygen.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(ii) the atomic radius of nitrogen and phosphorus.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (1)

(iii) the atomic and ionic radius of nitrogen.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(Total 5 marks)

28. State and explain the trends in the atomic radius and the ionization energy

(i) for the alkali metals Li to Cs.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(ii) for the period 3 elements Na to Cl.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(Total 8 marks)

29. (i) Describe three similarities and one difference in the reactions of lithium and potassium with water.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (4)

(ii) Give an equation for one of these reactions. Suggest a pH value for the resulting solution, and give a reason for your answer.

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (3)

(Total 7 marks)

30. (a) Classify each of the following oxides as acidic, basic or amphoteric.

(i) aluminium oxide

…………………………………………………………………………………… (1)

(ii) sodium oxide

…………………………………………………………………………………… (1)

(iii) sulfur dioxide

…………………………………………………………………………………… (1)

(b) Write an equation for each reaction between water and

(i) sodium oxide

……………………………………………………………………………………

…………………………………………………………………………………… (1)

(ii) sulfur dioxide.

……………………………………………………………………………………

…………………………………………………………………………………… (1)

(Total 5 marks)

31. This question is about Period 3 elements and their compounds.

(a) Explain, in terms of their structure and bonding, why the element sulfur is a non-conductor of electricity and aluminium is a good conductor of electricity.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (4)

(b) Explain, in terms of its structure and bonding, why silicon dioxide, SiO2, has a high melting point.

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(Total 6 marks)

32. Explain why

(i) the first ionization energy of magnesium is lower than that of fluorine. (2)

(ii) magnesium has a higher melting point than sodium. (3)

(Total 5 marks)

33. Discuss the acid-base nature of the period 3 oxides. Write an equation to illustrate the reaction of one of these oxides to produce an acid, and another equation of another of these oxides to produce a hydroxide.

(Total 5 marks)

34. Information about the halogens appears in the Data Booklet.

(i) Explain why the ionic radius of chlorine is less than that of sulfur. (2)

(ii) Explain what is meant by the term electronegativity and explain why the electronegativity of chlorine is greater than that of bromine.

(3) (Total 5 marks)

35. (a) (i) State the meaning of the term electronegativity and explain why the noble gases are not assigned electronegativity values.

(2)

(ii) State and explain the trend in electronegativity across period 3 from Na to Cl. (2)

(iii) Explain why Cl2 rather than Br2 would react more vigorously with a solution of I–. (2)

(b) State the acid-base properties of the following period 3 oxides.

MgO Al2O3 P4O6

Write equations to demonstrate the acid-base properties of each compound. (7)

(Total 13 marks)

36. (i) Define the term ionization energy. (1)

(ii) Write an equation for the reaction of lithium with water. (1)

(iii) State and explain the trend in the ionization energy of alkali metals down the group. (3)

(iv) Explain why the electronegativity of phosphorus is greater than that of aluminium.

(2)

(v) Table 8 in the Data Booklet contains two values for the ionic radius of silicon. Explain, by reference to atomic structure and electron arrangements, why the two values are very different.


37. Explain why sulfur has a lower first ionization energy than oxygen, and also a lower first ionization energy than phosphorus.

(Total 4 marks)

38. With reference to the types of bonding present in period 3 elements:

(i) explain why Mg has a higher melting point than Na. (2)

(ii) explain why Si has a very high melting point. (2)

(iii) explain why the other non-metal elements of period 3 have low melting points. (2)

(Total 6 marks)

39. Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character.

(Total 3 marks)

40. Explain the following statements.

(a) The first ionization energy of sodium is

(i) less than that of magnesium.

…………………………………………………………………………………

…………………………………………………………………………………

…………………………………………………………………………………

………………………………………………………………………………… (2)

(ii) greater than that of potassium.

…………………………………………………………………………………

………………………………………………………………………………… (1)

(b) The electronegativity of chlorine is higher than that of sulfur.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(Total 5 marks)

41. (a) (i) Define the term ionization energy.

.........................................................................................................................

......................................................................................................................... (2)

(ii) Write an equation, including state symbols, for the process occurring when measuring the first ionization energy of aluminium.

......................................................................................................................... (1)

(b) The first ionization energies of the elements are shown in Table 7 of the Data Booklet. Explain why the first ionization energy of magnesium is greater than that of sodium.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(c) Lithium reacts with water. Write an equation for the reaction and state two observations that could be made during the reaction.

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (3)

(Total 8 marks) 42. (a) State the meaning of the term electronegativity.

....................................................................................................................................

.................................................................................................................................... (1)

(b) State and explain the trend in electronegativity across period 3 from Na to Cl.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(c) Explain why Cl2 rather than Br2 would react more vigorously with a solution of I–.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 5 marks)

Topic 3 Answers

1. D [1]

2. B [1]

3. C [1]

4. B [1]

5. B [1]

6. C [1]

7. B [1]

8. C [1]

9. B [1]

10. A [1]

11. C [1]

12. A [1]

13. C [1]

14. D [1]

15. D [1]

16. D [1]

17. D [1]

18. A [1]

19. B [1]

20. D [1]

21. C [1]

22. C [1]

23. oxides of: Na, Mg: basic; Al: amphoteric; Si to Cl: acidic; Ar: no oxide;

All four correct [2], two or three correct [1].

Na2O + H2O → 2NaOH; SO3 + H2O → H2SO4; 4 Must be balanced for marks. Award marks for alternative correct equations such as SO3 with NaOH.

[4]

24. alkali metals: metallic bonding/a bed of cations in a sea of electrons; as radius increases down the group, valence electrons are further away from nucleus (and strength of metallic bonding decreases);

halogens: non-polar/van der Waals’ forces between molecules; as size increases van der Waals’ forces increase (and melting point increases);

period 3 elements: increase in melting points of metals (Na, Mg, Al) due to increase in number of valence electrons and decrease in size/the way atoms are packed as solids; Award mark just for “increased number of delocalized or valence electrons”.

silicon: network covalent solid (with very high melting point); Award mark also for “many or strong covalent bonds”.

P → Ar: simple molecular (atomic in case of Ar) substances with weak van der Waals’ forces (and lower melting points); trend in P4, S8, Cl2, Ar due to size/mass of particles; 8

Award mark for “decreasing mass or size”. Molecular formulae not necessary.

[8]

25. (i) and (ii) marked together.

K less than Na because electron removed (from K) is from higher energy level/further from nucleus/in n = 4 compared to n = 3; this is more important than the extra 8 protons in K/OWTTE; increase repulsion by extra shell of electrons/greater shielding effect; so less strongly attracted by nucleus;

K less than Ar because electron removed (from K) is from higher energy level/further from nucleus/ in n = 4 compare to n = 3; and has only one more proton; increase repulsion by extra shell of electrons/greater shielding effect; so less strongly attracted by nucleus;

Mg greater than Na because (Mg has) greater nuclear charge/one more proton/12 protons compare to 11; electron removed is in same (main) higher energy level/shell; smaller (atomic) radius; so more strongly attracted by nucleus; 7

Accept opposite worded arguments, i.e. why Na is greater than K. Award [7] for any seven correct but accept less/more strongly attracted to nucleus once only.

(iii)" second"electron"in"Na"removed"from"n"="2,"whereas"second"electron"in"Mg"removed"from"n"="3

" 1 [8]

26. (i) period is a horizontal row in the periodic table and a group is a vertical column/OWTTE;1

(ii)" 2,5;"electrons"in"two"energy"levels/shells;"five"outer/valence"electrons;" 3 [4]

27. (i) atomic radius of N > O because O has greater nuclear charge; greater attraction for the outer electrons/OWTTE; 2

(ii)" atomic"radius"of"P">"N"because"P"has"outer"electrons"in"an"energy""level"further"from"the"

nucleus/OWTTE;" 1"

(iii)" N3–">N/ionic"radius">"atomic"radius"because"N

3–"has"more"electrons""than"protons;"so"the"

electrons"are"held"less"tightly/OWTTE;" 2 Award [1] for greater repulsion in N3– due to more electrons (no reference to protons).

[5]

28. (i) Li to Cs atomic radius increases; because more full energy levels are used or occupied/outer electrons further from nucleus/outer electrons in a higher shell; ionization energy decreases; because the electron removed is further from the nucleus/increased repulsion by inner-shell electrons; 4

Accept increased shielding effect.

(ii) Na to Cl atomic radius decreases; because nuclear charge increases and electrons are added to same main (outer) energy level; ionization energy increases; because nuclear charge increases

and the electron removed is closer to the nucleus/is in the same energy level; 4 Accept “core charge” for “nuclear charge”. In (i) and (ii) explanation mark dependent on correct trend.

[8]

29. (i) similarities [3 max] the metal floats/moves on the surface; fizzing/effervescence/bubbles; (accept sound is produced) solution gets hot; solution becomes alkaline/basic; they react to form the metal hydroxide; hydrogen is evolved;

differences [1 max] flame/hydrogen burns with potassium (and not with lithium) /reaction faster/more vigorous with potassium/slower or less vigorous with lithium; 4 max

(ii) 2Li + 2H2O → 2Li+ + 2OH– + H2 / 2K + 2H2O → 2K+ + 2OH– + H2; Accept LiOH/KOH.

pH ≥ 11; LiOH/KOH is a strong base/strong alkali/high concentration of OH–; 3 [7]

30. (a) (i) aluminium oxide amphoteric;

(ii) sodium oxide basic;

(iii) sulfur dioxide acidic; 3

(b) (i) Na2O + H2O → 2Na+ + 2OH–;

(ii) SO2 + H2O → H2SO3; 2 Accept NaOH and H+ + HSO3

–/2H+ + SO32–.

[5]

31. (a) sulfur is (simple) molecular; (contains) covalent bonds/no delocalized electrons/all (outer) electrons used in bonding;

aluminium contains positive ions and delocalized electrons; (delocalized) electrons move (when voltage applied or current flows); 4

(b) silicon dioxide is macromolecular/giant covalent; many/strong covalent bonds must be broken; 2

Award max [1] if no mention of covalent. Do not accept weakened instead of broken.

[6]

32. (i) electron removed from higher energy level/further from nucleus/ greater atomic radius;

increased repulsion by extra inner shell electrons/increased shielding effect; 2

(ii) Mg has twice as many/more delocalized electrons (compared to Na); the ionic charge is twice as big/greater in Mg (than Na); (electrostatic) attraction between ions and electrons is much greater; 3

[5]

33. oxides of Na, Mg are basic Al is amphoteric Si, P, S and Cl are acidic

Award 7 correct [3], 6/5 correct [2] and 4/3 correct [1].

SO2 + H2O → H2SO3/SO3 + H2O → H2SO4/

P4O10 + 6H2O → 4H3PO4/P4O6 + 6H2O → 4H3PO3;

Na2O + H2O → 2NaOH/MgO + H2O → Mg(OH)2; 5 Accept equation using P2O3 or P2O5.

[5]

34. (i) (chlorine has) an extra proton/more protons/greater nuclear charge/ 17+ compared to 16+; outer electrons attracted more strongly; 2

(ii) ability of atom to attract bonding pair of electrons/electrons in a covalent bond;

chlorine has a smaller radius/(electrons) closer to nucleus/in lower energy level; repelled by fewer inner electrons/decreased shielding effect; 3

[5]

35. (a) (i) the ability of an atom to attract a bonding pair of electrons; inert/do not react/do not attract electrons/stable electron configuration/full outer electron

shell/do not form bonds; 2

(ii) electronegativity increases (along period 3 from Na to Cl); number of protons increases/nuclear charge increase/core charge increase/size of atom

decreases; 2 Do not accept “greater nuclear attraction”.

(iii) Cl2 stronger oxidising agent; Cl2 has greater attraction for electrons/has a higher electron affinity; 2

Accept converse statements for Br2.

(b) MgO − basic oxide/alkali; MgO + 2HCl → MgCl2 + H2O/MgO + H2O → Mg(OH)2;

Al2O3 − amphoteric oxide/acidic and basic oxide; Al2O3 + 6HCl → 2AlCl3 + 3H2O; Al2O3 + 2OH− + 3H2O → 2Al(OH)4

–/Al2O3 + 2OH− → 2AlO2– + H2O;

P4O6 − acidic oxide; P4O6 + 6H2O → 4H3PO3; 7

All equations must be balanced. [13]

36. (i) minimum energy required to remove one (mole of) electron(s) from (one mole of) (a) gaseous

atom(s)/OWTTE; 1

(ii) 2Li(s) + 2H2O(1) → 2LiOH(aq) + H2(g)/Li(s) + H2O(1) → LiOH(aq) + 1/2H2(g); 1 State symbols not required

(iii) (ionization energy) decreases; radius increases/valence electrons further away from nucleus/ electron removed from higher shell; (nuclear charge increases but) shielding/screening effect increases/ more electrons between

nucleus and valence electron/lower effective nuclear charge/Zeff; 3

(iv) phosphorus has a higher (effective) nuclear charge/Zeff; radius of P is smaller; electron pair/bonding electrons attracted more strongly; 2

(v) both have same number of protons/14 protons/nuclear charge/core charge; Si4+ formed by electron loss, Si4− formed by electron gain; Si4+ : 2.8 arrangement/2 (complete) energy levels/electrons in n = 2; Si4− : 2.8.8 arrangement/3 (complete) energy levels/electrons in n = 3; explanation of proton : electron ratio; higher effective nuclear charge/Zeff in Si4+; 4

[11]

37. IES < IEO:

valence electron in S in n = 3, in O in n = 2/e− further away/S has another electron shell/atomic radius of S greater than that of O; less attracted to nucleus/experiences greater screening from inner electrons; IES < IEP:

electron removed from S is paired;

greater repulsion due to two electrons in the same (p) orbital/paired electrons in S; 4 [4]

38. (i) Mg has greater nuclear charge/greater charge on cation/more valence e−/greater number of delocalized electrons/Na has lesser nuclear charge/lesser charge on cation/less valence e−/lesser number of delocalized electrons; stronger attraction between cation and delocalized/ free/valence electrons; 2

If neither mark scored, accept stronger metallic bonding in Mg for [1 max].

(ii) giant/network/lattice/macromolecular structure; many/strong covalent bonds (need to be broken); 2

(iii) (simple) molecular substances; weak van der Waals’/dispersion/London forces between molecules; 2

“Weak intermolecular forces” not sufficient for second mark [6]

39. Oxides of: Na and Mg are basic; Al is amphoteric; Si to Cl are acidic; Ar has no oxide; All four correct award [2], two or three correct award [1].

Na2O + H2O → 2NaOH and SO3 + H2O → H2SO4; 3

Must be balanced for mark. Award marks for alternative correct equations such as SO3 with NaOH.

[3]

40. (a) (i) Na has lower nuclear charge/number of protons; electrons being removed are from same energy level/shell; or Na has larger radius/electron further from nucleus; 2 max

Award this mark if both electron arrangements are given.

(ii) Na electron closer to nucleus/in lower energy level/Na has less 1 shielding effect; Allow counter arguments for Mg in (i) and K in (ii).

(b) chlorine has a higher nuclear charge; attracts the electron pair/electrons in bond more strongly; 2

[5]

41. (a) (i) the (minimum) energy required/needed for the removal of one electron; from a gaseous/isolated atom; 2

(ii) Al(g) → Al+(g) + e; 1 Do not penalize the answer if (g) is after e.

(b) greater nuclear charge/greater number of protons/atom radius g is smaller; stronger attraction (for electron); 2

(c) 2Li + 2H2O → 2LiOH + H2; Ignore state symbols.

effervescence/fizzing/bubbles/OWTTE; lithium moves around/decrease in size of piece;

Accept dissolves or disappears.

heat produced; 3 Award [1] each for any two of last three observations.

[8] 42. (a) the ability of an element/atom/nucleus to attract a bonding pair of electrons; 1

(b) electronegativity increases (along period 3 from Na to Cl); number of protons increases/nuclear charge increases/core charge increases /size of atoms

decreases; 2 Do not accept greater nuclear attraction.

(c) Cl2 is a stronger oxidizing agent/Chlorine’s outer shell closer to nucleus; Cl2 has greater attraction for electrons/has a higher electron affinity; 2

Accept converse argument for Br2. [5]

Topic 2 Questions

1. Consider the composition of the species W, X, Y and Z below. Which species is an anion?

Species Number of protons Number of neutrons Number of electrons

W 9 10 10

X 11 12 11

Y 12 12 12

Z 13 14 10

A. W

B. X

C. Y

D. Z (Total 1 mark)

2. Energy levels for an electron in a hydrogen atom are

A. evenly spaced.

B. farther apart near the nucleus.

C. closer together near the nucleus.

D. arranged randomly. (Total 1 mark)

3. Which is related to the number of electrons in the outer main energy level of the elements from the alkali metals to the halogens?

I. Group number

II. Period number

A. I only

B. II only

C. Both I and II

D. Neither I nor II (Total 1 mark)

4. How do bond length and bond strength change as the number of bonds between two atoms increases?

Bond length Bond strength

A. Increases increases

B. Increases decreases

C. Decreases increases

D. Decreases decreases (Total 1 mark)

5. Which of the following is true for CO2?

CO bond CO2 molecule

A. Polar non-polar

B. non-polar polar

C. Polar polar

D. non-polar non-polar (Total 1 mark)

6. The molar masses of C2H6, CH3OH and CH3F are very similar. How do their boiling points compare?

A. C2H6 < CH3OH < CH3F

B. CH3F < CH3OH < C2H6

C. CH3OH < CH3F < C2H6

D. C2H6 < CH3F < CH3OH (Total 1 mark)

7. What is the correct number of each particle in a fluoride ion, 19F–?

protons neutrons electrons

A. 9 10 8

B. 9 10 9

C. 9 10 10

D. 9 19 10 (Total 1 mark)

8. Which statement is correct for the emission spectrum of the hydrogen atom?

A. The lines converge at lower energies.

B. The lines are produced when electrons move from lower to higher energy levels.

C. The lines in the visible region involve electron transitions into the energy level closest to the

nucleus.

D. The line corresponding to the greatest emission of energy is in the ultraviolet region.

9. Which is the correct description of polarity in F2 and HF molecules?

A. Both molecules contain a polar bond.

B. Neither molecule contains a polar bond.

C. Both molecules are polar.

D. Only one of the molecules is polar. (Total 1 mark)

10. Which types of bonding are present in CH3CHO in the liquid state?

I. Single covalent bonding

II. Double covalent bonding

III. Hydrogen bonding

A. I and II only

B. I and III only

C. II and III only


11. Which statement(s) is/are generally true about the melting points of substances?

I. Melting points are higher for compounds containing ions than for compounds containing molecules.

II. A compound with a low melting point is less volatile than a compound with a high melting point.

III. The melting point of a compound is decreased by the presence of impurities.

A. I only

B. I and III only

C. II and III only


12. How many valence electrons are present in an atom of an element with atomic number 16?

A. 2

B. 4

C. 6

D. 8 (Total 1 mark)

13. A certain sample of element Z contains 60% of 69Z and 40% of 71Z. What is the relative atomic mass of element Z in this sample?

A. 69.2

B. 69.8

C. 70.0

D. 70.2 (Total 1 mark)

14. What is the difference between two neutral atoms represented by the symbols Co and Ni?

A. The number of neutrons only.

B. The number of protons and electrons only.

C. The number of protons and neutrons only.

D. The number of protons, neutrons and electrons. (Total 1 mark)

15. How many electrons are there in one ion?

A. 10

B. 12

C. 14


16. The electron arrangement of sodium is 2.8.1. How many occupied main electron energy levels are there in an atom of sodium?

A. 1

B. 3

C. 10


17. Information is given about four different atoms:

Atom neutrons protons

W 22 18

X 18 20

Y 22 16

Z 20 18

Which two atoms are isotopes?

A. W and Y

B. W and Z

C. X and Z

D. X and Y (Total 1 mark)

18. Which statement is correct about a line emission spectrum?

A. Electrons absorb energy as they move from low to high energy levels.

B. Electrons absorb energy as they move from high to low energy levels.

C. Electrons release energy as they move from low to high energy levels.

D. Electrons release energy as they move from high to low energy levels. (Total 1 mark)

19. How many neutrons are there in the ion 18O2–?

A. 8

B. 10

C. 16


20. What is the electron arrangement of silicon?

A. 2.4

B. 2.8

C. 2.8.4

D. 2.8.8 (Total 1 mark)

21. Which statement is correct about the isotopes of an element?

A. They have the same mass number

B. They have the same electron arrangement

C. They have more protons than neutrons

D. They have the same numbers of protons and neutrons (Total 1 mark)

22. What is the difference between two neutral atoms represented by the symbols Po and At?

A. The number of neutrons only.

B. The number of protons and electrons only.

C. The number of protons and neutrons only.

D. The number of protons, neutrons and electrons. (Total 1 mark)

23. Which statements are correct for the emission spectrum of the hydrogen atom?

I. The lines converge at lower energies. II. Electron transition to n =1 are responsible for lines in the UV region. III. Lines are produced when electrons move from higher to lower energy levels.

A. I and II only

B. I and III only

C. II and III only


24. What is the symbol for a species that contains 15 protons, 16 neutrons and 18 electrons?

A.

B.

C.

D. (Total 1 mark)

25. What is the electron arrangement of an Al3+ ion?

A. 2, 8

B. 2, 3

C. 2, 8, 3

D. 2, 8, 8 (Total 1 mark)

26. What will happen to the volume of a fixed mass of gas if the pressure and the Kelvin temperature are both doubled?

A. It will remain the same.

B. It will be double its initial volume.

C. It will be one-half its initial volume.

D. It will be four times its initial volume. (Total 1 mark)

27. Which species has 54 electrons and 52 protons?

A.

B.

C.

D. (Total 1 mark)

28. What is the correct sequence for the processes occurring in a mass spectrometer?

A. vaporization, ionization, acceleration, deflection

B. vaporization, acceleration, ionization, deflection

C. ionization, vaporization, acceleration, deflection

D. ionization, vaporization, deflection, acceleration (Total 1 mark)

29. The percentage composition by mass of a hydrocarbon is C = 85.6% and H = 14.4%.

(a) Calculate the empirical formula of the hydrocarbon.

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.....................................................................................................................................

..................................................................................................................................... (2)

(b) A 1.00 g sample of the hydrocarbon at a temperature of 273 K and a pressure of 1.01×105 Pa (1.00 atm) has a volume of 0.399 dm3.

(i) Calculate the molar mass of the hydrocarbon.

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..........................................................................................................................

..........................................................................................................................

.......................................................................................................................... (2)

(ii) Deduce the molecular formula of the hydrocarbon. (1)

(Total 5 marks)

30. State the number of protons, electrons and neutrons in the ion N3–.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

(Total 2 marks)

31. A sample of germanium is analysed in a mass spectrometer. The first and last processes in mass spectrometry are vaporization and detection.

(a) (i) State the names of the other three processes in the order in which they occur in a mass spectrometer.

.........................................................................................................................

......................................................................................................................... (2)

(ii) For each of the processes named in (a) (i), outline how the process occurs.

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.........................................................................................................................

......................................................................................................................... (3)

(b) The sample of germanium is found to have the following composition:

Isotope 70Ge 72Ge 74Ge 76Ge

Relative abundance / % 22.60 25.45 36.73 15.22

(i) Define the term relative atomic mass.

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(ii) Calculate the relative atomic mass of this sample of germanium, giving your answer to two decimal places.

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(Total 9 marks)

32. Define the following terms.

(i) atomic number

....................................................................................................................................

.................................................................................................................................... (1)

(ii) mass number

....................................................................................................................................

.................................................................................................................................... (1)

(Total 2 marks)

33. State the electron arrangements of the following species:

Si ........................................................................................................................................

P3– ........................................................................................................................................ (Total 2 marks)

34. Identify the numbers of protons, neutrons and electrons in the species 33S2–.

..............................................................................................................................................

.............................................................................................................................................. (Total 1 mark)

35. State the electron arrangement for atoms of aluminium, nitrogen and fluorine. (Total 2 marks)

36. The relative atomic mass of chlorine is 35.45. Calculate the percentage abundance of the two isotopes of chlorine, 35Cl and 37Cl in a sample of chlorine gas.

(Total 2 marks)

37. (a) Describe the following stages in the operation of the mass spectrometer.

(i) ionization (2)

(ii) deflection (2)

(iii) acceleration (1)

(b) (i) State the meaning of the term isotopes of an element. (1)

(ii) Calculate the percentage abundance of the two isotopes of rubidium 85Rb and 87Rb. (2)

(iii) State two physical properties that would differ for each of the rubidium isotopes.

(1)

(iv) Determine the full electron configuration of an atom of Si, an Fe3+ ion and a P3– ion. (3)

(Total 12 marks)

38. Naturally occurring copper has a relative atomic mass, (Ar), of 63.55 and consists of two isotopes 63Cu and 65Cu.

(i) Define the term relative atomic mass, Ar.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (1)

(ii) State and explain which is the more abundant isotope.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (1)

(Total 2 marks)

39. The element vanadium has two isotopes, and and a relative atomic mass of 50.94.

(a) Define the term isotope.

……………………………………………………………………………………….

………………………………………………………………………………………. (1)

(b) State the number of protons, electrons and neutrons in

……………………………………………………………………………………….

………………………………………………………………………………………. (2)

(c) State and explain which is the more abundant isotope.

……………………………………………………………………………………….

………………………………………………………………………………………. (1)

(d) State the name and the mass number of the isotope relative to which all atomic masses are measured.

………………………………………………………………………………………. (1)

(Total 5 marks)

40. (a) State a physical property that is different for isotopes of an element.

..................................................................................................................................... (1)

(b) Chlorine exists as two isotopes, 35Cl and 37Cl. The relative atomic mass of chlorine is 35.45. Calculate the percentage abundance of each isotope.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 3 marks)

41. (a) Define the term isotope.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(b) A sample of argon exists as a mixture of three isotopes.

mass number 36, relative abundance 0.337% mass number 38, relative abundance 0.0630% mass number 40, relative abundance 99.6%

Calculate the relative atomic mass of argon.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(c) State the number of electrons, protons and neutrons in the ion 56Fe3+.

electrons: ............................. protons: ............................. neutrons: ........................... (2)

(Total 6 marks)

42. The element bromine exists as the isotopes 79Br and 81Br, and has a relative atomic mass of 79.90.

(a) Complete the following table to show the numbers of sub-atomic particles in the species shown.

an atom of 79Br an ion of 81Br–

protons

neutrons

electrons (3)

(b) State and explain which of the two isotopes 79Br and 81Br is more common in the element bromine.

...................................................................................................................................

................................................................................................................................... (1)

(c) The element calcium is in the same period of the Periodic Table as bromine.

(i) Write the electron arrangement for an atom of calcium.

......................................................................................................................... (1)

(ii) Deduce the formula of the compound calcium bromide.

......................................................................................................................... (1)

(Total 6 marks)

43. Some vaporized magnesium is introduced into a mass spectrometer. One of the ions that reaches the detector is 25Mg+.

(a) Identify the number of protons, neutrons and electrons in the 25Mg+ ion.

....................................................................................................................................

.................................................................................................................................... (1)

(b) State how this ion is accelerated in the mass spectrometer.

....................................................................................................................................

.................................................................................................................................... (1)

(c) The 25Mg2+ ion is also detected in this mass spectrometer by changing the magnetic field. Deduce and explain, by reference to the m/z values of these two ions of magnesium, which of the ions 25Mg2+ and 25Mg+ is detected using a stronger magnetic field.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 4 marks) 44. (a) List the following types of electromagnetic radiation in order of increasing wavelength (shortest

first).

I. Yellow light

II. Red light

III. Infrared radiation

IV. Ultraviolet radiation

..................................................................................................................................... (1)

(b) Distinguish between a continuous spectrum and a line spectrum.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (1)

(c) The thinning of the ozone layer increases the amount of UV-B radiation that reaches the Earth’s surface.

Type of Radiation Wavelength / nm

UV-A 320–380

UV-B 290–320

Based on the information in the table above explain why UV-B rays are more dangerous than UV-A.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(Total 5 marks)

Topic 2 Answers

1. A [1]

2. B [1]

3. A [1]

4. C [1]

5. A [1]

6. D [1]

7. C [1]

8. D [1]

9. D [1]

10. A [1]

11. B [1]

12. C [1]

13. B [1]

14. D

[1]

15. A [1]

16. B [1]

17. B [1]

18. D [1]

19. B [1]

20. C [1]

21. B [1]

22. D [1]

23. C [1]

24. D [1]

25. A [1]

26. A [1]

27. A [1]

28. A [1]

29. (a) mole ratio C:H = = 7.13:4.3; No penalty for using integer atomic masses.

empirical formula is 2

(b) (i) number of moles of gas n = ;

= 56.3 (g mol–1) 2

OR

molar mass is the at STP; = = 56.1 (g mol–1)

Accept answers in range 56.0 to 56.3. Accept two, three or four significant figures.

(ii) C4H8; 1 No ECF.

[5]

30. 7 protons, 8 neutrons, 10 electrons; 2 Award [2] for three correct and [1] for two correct.

[2]

31. (a) (i) ionization, acceleration, deflection/separation; 2 Award [1] for all three names and [1] for correct order. Award [1] for two names in correct order.

(ii) ionization: sample bombarded with high-energy or high-speed electrons/OWTTE; acceleration: electric field/oppositely charged plates; deflection: (electro)magnet/magnetic field; 3

(b) (i) average or (weighted) mean of masses of all isotopes of an element; relative to (one atom of) 12C; 2

Both marks available from a suitable expression.

(ii) Ar = (70⋅0.2260) + (72⋅0.2545) + (74⋅0.3673) + (76⋅0.1522);

= 72.89; 2 No other final answer acceptable. Award [2] for correct final answer.

[9]

32. (i) number of protons in the nucleus/atom; 1 Do not accept protons and electrons.

(ii) number of protons and neutrons in the nucleus/atom; 1 [2]

33. Si 2.8.4/2,8,4; P3− 2.8.8/2,8,8; 2 [2]

34. 16 protons and 17 neutrons and 18 electrons; 1 [1]

35. Al − 2,8,3; N − 2,5; F − 2,7; 2 Award [2] for three correct, [1] for two or one correct. Accept correct configuration using s,p,d notation.

[2]

36. Ar(Cl) = 35.45 = ;

35Cl = 77.5% and 37Cl = 22.5%; 2 [2]

37. (a) (i) to produce positively charged ions; by the bombardment of fast moving electrons; 2

(ii) magnetic field at right angles to path of ions/accept suitably labelled diagram; moves ions in curve path/deflects ions; dependent on mass/charge ratio; 2

Award [1] each for any 2 points.

(iii) acceleration of the ions by electric field/towards negative plate/cathode; 1

(b) (i) atoms with the same number of protons/positive charges/atomic number but different number of neutrons/mass number; 1

(ii) Ar(Rb) = 85.47 = Accept other valid mathematical alternatives.

85Rb = 76.5 and 87Rb = 23.5%; 2

(iii) mass; density; boiling point; melting point; rate of diffusion in the gas phase; enthalpy of vaporization; enthalpy of fusion; rate of reaction in the gas/liquid phase; 1

Any two for one mark

(iv) Si: 1s22s22p63s23p2;

Fe3+: 1s22s22p63s23p63d5;

P3− : 1s22s22p63s23p6; 3 Allow [1 max] for 3 correct abbreviated structures using noble gas symbols.

38. (i) ratio of average mass of an atom to the mass of C-12 isotope/ average mass of an atom on a scale where one atom of C-12 has a mass of 12/sum of the weighted average mass of isotopes of an element compared to C-12/OWTTE; 1

Award no mark if ‘element’ is used in place of ‘atom’

(ii) 63Cu (more abundant) since Ar (Cu) is closer in mass to 63; 1 Explanation needed for mark

[2]

39. (a) atom of same element/same number of protons but with different mass number/number of neutrons; 1

(b) protons 23 electrons 23 neutrons 27

2 Three correct [2], two correct [1].

(c) /51 nearer to Ar value of 50.94; 1

(d) carbon, 12/12C; 1 [5]

40. (a) mass/density/for gases: rate of effusion or diffusion/melting point/ boiling point 1 Do not accept mass number.

(b) if 35Cl = x, then (x = 35.00) + (1 – x) 37.00 = 35.45 Award [1] for set up.

therefore, x = 0.775; 2 35Cl = 77.5% and 37Cl = 22.5%; (need both for mark); [3]

41. (a) atoms of the same element/same number of protons/same atomic number; having different numbers of neutrons/different (mass number); 2

Award only [1] max if reference made to elements but not atoms.

(b) relative atomic mass = ; 2

(c) 23 electrons; 26 protons; 30 neutrons; 2 Award [2] for three correct, [1] for two correct.

[6]

42. (a)

an atom of 79Br an ion of 81Br–

protons 35 35 ;

neutrons 44 46 ;

electrons 35 36 ;

3

(b) 79Br because Ar is closer to 79/OWTTE; 1

(c) (i) 2,8,8,2/2.8.8.2; 1

(ii) CaBr2; 1 [6]

43. (a) 12 protons and 13 neutrons and 11 electrons; 1

(b) electric field/oppositely charged plates/potential difference/OWTTE; 1

(c) 25Mg+; greater m/z value/less highly charged ions need stronger fields to deflect them/OWTTE; 2

Do not accept greater mass with no reference to charge, or greater mass and smaller charge.

[4]

44. (a) IV < I < II < III/ ultra violet radiation < yellow light < red light < infrared radiation; 1

(b) A continuous spectrum has all colours/wavelengths/frequencies whereas a line spectrum has only (lines of) sharp/discrete/specific colours/ wavelengths/frequencies; 1

(c) UV-B radiation has shorter wavelength; hence, has higher energy; increases risk of damage to skin cells/OWTTE/causes cancer; 3

[5] Topic13: Human Biochemistry Option B – Questions

1. Polypeptides and proteins are formed by the condensation reactions of amino acids.

(a) Give the general structural formula of a 2-amino acid.

(1)

(b) Give the structural formula of the dipeptide formed by the reaction of alanine and glycine.

State the other substance formed during this reaction.

………………………………………………………………………………………… (2)

(c) State two functions of proteins in the body.

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(Total 5 marks)

2. Electrophoresis can be used to identify the amino acids present in a given protein. The protein must first be hydrolyzed.

(i) State the reagent and conditions needed to hydrolyze the protein, and identify the bond that is broken during hydrolysis.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (4)

(ii) Explain how the amino acids could be identified using electrophoresis.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (4)

(Total 8 marks)

3. (a) Draw the straight chain structure of glucose.

(1)

(b) The structure of α-glucose is shown below.

Outline the structural difference between α-glucose and β-glucose.

......................................................................................................................................

...................................................................................................................................... (1)

(c) Glucose molecules can condense to form starch which can exist in two forms, amylose and amylopectin. Describe the structural differences between the two forms.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (2)

(d) 1.00 g of sucrose, C12H22O11, was completely combusted in a food calorimeter. The heat evolved was equivalent to increasing the temperature of 631 g of water from 18.36°C to 24.58 °C. Calculate the calorific value of sucrose (in kJ mol–1) given the specific heat capacity of water in Table 2 of the Data Booklet.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (3)

(Total 7 marks)

4. (a) The structures of three important vitamins are shown in Table 22 of the Data Booklet. State the name of each one and deduce whether each is water-soluble or fat-soluble, explaining your choices by reference to their structures.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (5)

(b) Identify the metal ion needed for the maintenance of healthy bones and state the name of the vitamin needed for its uptake.

.....................................................................................................................................

..................................................................................................................................... (2)

(c) State the name of the vitamin responsible for maintaining healthy eyesight and the name of the functional group which is most common in this vitamin.

.....................................................................................................................................

..................................................................................................................................... (2)

(d) Identify one major function of vitamin C in the human body and state the name of the most common disease caused by deficiency of this vitamin.

.....................................................................................................................................

..................................................................................................................................... (2)

(e) Fresh fruits and vegetables are good sources of vitamin C. Explain why some meals made from these foods may contain little vitamin C.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 13 marks)

5. The structures of two sex hormones, progesterone and testosterone, are shown in Table 22 of the Data Booklet.

(a) State the names of two functional groups that are present in both hormones.

.....................................................................................................................................

..................................................................................................................................... (2)

(b) Identify which of the two hormones is the female sex hormone and where in the human body it is produced.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(c) Outline the mode of action of oral contraceptives.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(Total 7 marks)

6. By referring to Table 22 of the Data Booklet, identify one vitamin that is water soluble and one vitamin that is fat soluble. Explain the differences in solubility in terms of their structures and intermolecular forces.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................ (Total 4 marks)

7. Discuss two benefits of using genetically modified foods.

...............................................................................................................................................

...............................................................................................................................................

...............................................................................................................................................

............................................................................................................................................... (Total 2 marks)

8. The structure of lactose, a disaccharide formed from glucose and galactose, is shown in the Data Booklet. Draw the ring structure of galactose and state whether it is an 〈 or ® isomer.

(Total 2 marks)

9. (a) The general formula for saturated fatty acids is CnC2nO2. The molecular formula of linoleic acid is C18H32O2.

(i) Determine the number of carbon to carbon double bonds in linoleic acid.

........................................................................................................................

........................................................................................................................ (1)

(ii) Iodine number is defined as the number of grams of iodine that adds to 100 g of a fat or an oil in an addition reaction. Determine the iodine number of linoleic acid.

........................................................................................................................

........................................................................................................................

........................................................................................................................ (2)

(b) (i) State one structural similarity between fats and oils.

........................................................................................................................

........................................................................................................................ (1)

(ii) Explain, by referring to their structures, why fats are solid at room temperature, but oils are liquid.

........................................................................................................................

........................................................................................................................

........................................................................................................................

........................................................................................................................

........................................................................................................................

........................................................................................................................ (3)

(Total 7 marks)

10. (a) State the name of a disease which results from the deficiency of each of the following vitamins.

vitamin A .................................................................................................................

vitamin B .................................................................................................................

vitamin C ................................................................................................................. (2)

(b) A person consumes an excess of both vitamin A and C. State, with a reason, which one is more likely to be stored in the body and which is more likely to be excreted.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(Total 4 marks)

11. Identify three types of interactions responsible for the tertiary structure of proteins.

.............................................................................................................................................

.............................................................................................................................................

.............................................................................................................................................

............................................................................................................................................. (Total 2 marks)

12. The structures of vitamins A and C are shown in Table 22 of the Data Booklet. State, with a reason, whether each is fat soluble or water soluble.

.............................................................................................................................................

.............................................................................................................................................

.............................................................................................................................................

.............................................................................................................................................

.............................................................................................................................................

............................................................................................................................................. (Total 3 marks

13. (a) (i) Deduce the structure of one of the dipeptides that can be formed when the two aminoacids below react together.

(2)

(ii) State the name given to this type of reaction and identify the other product of the reaction.

.........................................................................................................................

......................................................................................................................... (2)

(b) Describe how a mixture of aminoacids can be analysed using electrophoresis.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (4)

(c) (i) Explain what is meant by the primary structure of proteins.

.........................................................................................................................

......................................................................................................................... (1)

(ii) Explain, with reference to hydrogen bonding, why the 〈-helix and ®-sheet secondary structures of proteins are different.

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (2)

(Total 11 marks)

14. (a) The equilibria, which exist in an aqueous solution of glycine, are shown in the structures below.

H2N( CH2 ( COO− H3N+( CH2( COO− H3N+( CH2( COOH A B C

State which of the forms A, B or C occurs in the greatest concentration at:

low pH: ...............................

high pH: .............................. (2)

(b) A mixture of amino acids with different isoelectric points can be separated using electrophoresis.

(i) Outline the essential features of electrophoresis.

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

.........................................................................................................................

......................................................................................................................... (3)

(ii) Arginine, glutamic acid and glycine undergo electrophoresis at pH 6.0. Using table 20 of the Data Booklet identify the amino acid that moves towards:

the positive electrode: .....................................................................................

the negative electrode: .................................................................................... (2)

(Total 7 marks)

15. Iodine number is defined as the number of grams of iodine that reacts with 100 g of a triglyceride in an addition reaction. The iodine number of palmitic acid (Mr = 256) is 0 and linolenic acid (Mr = 278) is 274.

Determine the number of double bonds in linolenic acid, showing your working.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 3 marks)

16. The structures of two synthetic hormones are shown below:

Hormone A is similar in structure to testosterone and hormone B is similar in structure to progesterone.

(a) Explain why hormone A is prescribed to some patients.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Suggest why hormone A is banned for participants in major sporting events.

....................................................................................................................................

.................................................................................................................................... (1)

(c) Describe how hormone B functions as an oral contraceptive.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 5 marks)

17. The structures of the amino acids glycine and serine are shown in Table 20 of the Data Booklet.

(i) Draw the structure of one of the dipeptides formed when one molecule of glycine and one molecule of serine react together. Show all the bonds in the link between the two molecules.

(2)

(ii) State the type of reaction occurring and identify the other product of the reaction.

....................................................................................................................................

.................................................................................................................................... (1)

(Total 3 marks)

18. State the general role of hormones in the body and identify the gland that controls their production.

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 2 marks)

19. The structures of the amino acids glycine and serine are shown in Table 20 of the Data Booklet. Draw the structure of one of the dipeptides formed when one molecule of glycine and one molecule of serine react together. Show all the bonds in the link between the two molecules.

(Total 2 marks)

20. The structure of a protein can be analysed using paper chromatography.

(i) Describe the process that the protein must undergo before chromatography is used and explain why it is necessary.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(ii) Explain how paper chromatography is used to identify the individual amino acids.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (4)

(Total 6 marks)

21. Fats and oils are formed when fatty acids react with glycerol.

(a) Outline two structural differences between saturated and unsaturated fats.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Explain why saturated fats have higher melting points than unsaturated fats with similar relative molecular masses.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 4 marks)

22. The structures of some hormones are shown in Table 22 of the Data Booklet.

(a) Identify one hormone with a steroid backbone, state where it is produced and outline its specific

role in the body.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Identify one hormone with a non-steroid backbone, state where it is produced and outline its specific role in the body.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 4 marks)

23. The structure of sucrose is shown in Table 22 of the Data Booklet.

(a) State the name of the oxygen-containing link between the two rings in the structure.

.................................................................................................................................... (1)

(b) Deduce the ring structures of the two monosaccharides that condense to form a molecule of sucrose.

(2)

(c) Identify the other product formed during this condensation reaction.

.................................................................................................................................... (1)

(d) State the empirical formula of a monosaccharide.

.................................................................................................................................... (1)

(Total 5 marks)

24. The structures of three vitamins are shown in Table 22 of the Data Booklet.

(a) Predict which of the three vitamins is most soluble in water, giving a reason for your choice.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Name two functional groups present in all three vitamins.

....................................................................................................................................

.................................................................................................................................... (1)

(c) State the function of vitamin D in the human body and describe one effect of vitamin D deficiency.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 5 marks)

25. The structure of sucrose is shown in Table 22 of the Data Booklet.

(a) State the name of the oxygen-containing link between the two rings in the structure.

.................................................................................................................................... (1)

(b) Deduce the ring structures of the two monosaccharides that condense to form a molecule of sucrose.

(2)

(c) State the empirical formula of a monosaccharide.

....................................................................................................................................

.................................................................................................................................... (1)

(Total 4 marks)

26. The structures of three vitamins are shown in Table 22 of the Data Booklet.

(a) Predict which of the three vitamins is most soluble in water, giving a reason for your choice.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) State which two vitamins can be classified as primary alcohols.

....................................................................................................................................

.................................................................................................................................... (1)

(c) State the function of vitamin D in the human body and describe one effect of vitamin D deficiency.

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 5 marks)

27. When many 2-amino acid molecules react together a protein is formed. These proteins have primary, secondary and tertiary structures.

(a) State the type of intermolecular force responsible for maintaining the secondary structure.

..................................................................................................................................... (1)

(b) Describe two other ways in which the tertiary structure of the protein is maintained.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 3 marks)

28. (a) For each of the following vitamins describe its function in a diet and one effect of its deficiency.

Vitamin C ...................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

Vitamin D ...................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (4)

(b) Discuss two solutions for the prevention of nutrient deficiencies.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 6 marks)

29. (a) State what is meant by dietary fibre.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(b) Give two examples of dietary fibre.

.....................................................................................................................................

..................................................................................................................................... (2)

(c) Describe two reasons for the inclusion of dietary fibre in a healthy diet.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 6 marks) 30. (a) Compare the structural properties of starch and cellulose.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (4)

(b) Explain why humans cannot digest cellulose.

.....................................................................................................................................

..................................................................................................................................... (1)

(Total 5 marks) Topic13: Human Biochemistry Option B – Answers

1. (a) RCH(NH2)COOH; 1

(b) H2NCH(CH3)CONHCH2COOH / H2NCH2CONHCH(CH3)COOH water/H2O; 2

(c) structure/catalysis or enzymes/energy source/oxygen transport; 2 Any two, [1] each. Accept specific structures, e.g. hair, muscle.

[5]

2. (i) acid/hydrochloric acid/HCl (accept H2SO4) Accept base/NaOH. concentrated/heat or high temperature or boil/ time (any two, [1] each);

4

(ii) mixture/amino acids spotted on paper/gel; apply voltage; develop/ninhydrin/organic dye; measure distances moved/compare with known samples/measure isoelectric points and compare with data; 4

Marks may be given for a suitable diagram. [8]

3. (a)

1 No penalty for “sticks” or for OH groups written back-to-front, eg. OH- instead of HO-.

(b) the –OH group on the first carbon atom is inverted in β-glucose 1

(c) one (amylose) is a straight chain polymer whereas the other (amylopectin) is branched; one

(amylose) has only 1,4 bonds (between the monomers) whereas the other (amylopectin) has 1,4 and 1,6 bonds; 2

(d) Mr for sucrose =342;

heat evolved = 0.631(kg)×4.18 (kJ kg–1K–1 )×6.22(K) = 16.4 kJ;

calorific value = = 5.61×103 kJ mol–1; 3 Allow answers in range 5610 to 5620. Penalize for more than 5 sig. figs. ECF from incorrect Mr.

[7]

4. (a) vitamin A retinol is fat-soluble; vitamin C ascorbic acid is water-soluble; vitamin D calciferol is fat-soluble;

fat-soluble because mainly composed of hydrocarbon chain/non polar groups; water-soluble because of presence of several/many hydroxyl/ OH/polar groups; 5

(b) Ca2+/calcium;

vitamin D/calciferol; 2

(c) vitamin A/retinol; alkene; 2

(d) helps to form collagen/connective tissue/acts as antioxidant; 2 scurvy/scorbutus;

(e) dissolves in water; oxidized/destroyed by heating/boiling; 2 [13]

5. (a) carbonyl/ketone Accept alkanone but not aldehyde.

alkene; 2

(b) progesterone; ovaries; 2

(c) change release of hormones/FHS/LH (from hypothalamus/pituitary gland); prevent ovulation/egg release; prevent attachment of egg to uterus; prevent sperm from reaching egg; 3

Award [1] each for any three. [7]

6. vitamin C is water soluble and vitamin A/D is fat soluble; vitamin C has 4/several OH groups/vitamin A/D has only 1/fewer OH groups; vitamin A/D has large non-polar/hydrocarbon part/chain/ring; vitamin C has hydrogen bonding and vitamin A/D has van der Waals’ forces; 4

[4]

7. Benefits improve food productivity/provide more food; (food) crops are more resistant to disease/more resistant to insect attack /more tolerant to toxins; improve aesthetics/composition of some foods; improved flavour; improved texture; improved nutritional value; improved shelf life; incorporation of anti-cancer substances/vaccines/reduce exposure to less healthy fats; 2 max

Award [1] each for any two. [2]

8.

[2]

9. (a) (i) 2 double bonds; 1

(ii) 280.5 g linoleic acid adds to 507.6 g I2;

2 Allow ECF from (i). Do not penalize for use of whole number atomic masses.

(b) (i) both are esters/tri-glycerides/tri-esters; 1

(ii) Fats Oil saturated/no C=C bonds or unsaturated/1 or more C=C

bonds; (saturated) chains pack closely or (unsaturated) chains pack less

closely; van der Waals’ forces are stronger or van der Waals’ forces are weaker; 3

Accept intermolecular forces for van der Waals’ forces. [7]

10. (a) vitamin A night blindness/xerophthalmia;

vitamin C scurvy/scorbutus;

vitamin D rickets; 2

Award [2] for 3 correct, [1] for 2 correct.

(b) vitamin A is stored (in the body) because it is fat-soluble;

vitamin C is excreted because it is water-soluble; 2

[4]

11. hydrogen bonding; disulphide links/bonds/bridges; van der Waals’ forces; ionic/ion-dipole/dipole-dipole;

Award [2] for any three. Award [1] for any two.

[2]

12. A is fat soluble and C is water soluble; A has only one OH group/A is mostly hydrocarbon; C has many OH groups which can form hydrogen bonds with water;

Do not penalise if OH is stated with a minus sign. [3]

13. (a) (i) 2 Award [1] for the correct peptide bond and an additional [1] if the rest of the structure is correct.

(ii) condensation; H2O/water; 2

(b) mixture placed on gel/paper; use of buffer solution; potential difference applied; amino acids move differently (depending on pH/isoelectric point); develop/spray with ninhydrin;

compare distances travelled with standards (OWTTE)/compare the isoelectric points; 4 Award [1] each for any four.

(c) (i) sequence/chain of amino acids; 1

(ii) 〈 - helix = intramolecular/spiral/OWTTE; ® - sheet = attraction between chains (accept intermolecular)/OWTTE; 2

Accept suitable diagrams. [11]

14. (a) low pH C; high pH A; 2

(b) (i) place sample on gel; with (buffer) solution of known pH; apply voltage/potential difference;

Do not accept current applied. develop/spray with ninhydrin;

measure distance moved/compare with known iso-electric point/ compare with standards; 3 Award [1] each for any three.

(ii) positive electrode glutamic acid;

negative electrode arginine; 2

[7]

15. 761.7 g of I2/274÷253.8 = 1.08 mol; 3 mol of I2/6 mol of I atoms/100÷278 = 0.360 mol; (1.08÷0.360 = ) 3 double bonds;

Some correct working must be shown. Allow ECF if Mr of iodine used as 126.9 instead of 253.8. Accept correct alternative methods.

[3]

16. (a) recovery from injury/surgery/starvation/illness/disease; increased rate of protein synthesis/tissue/muscle building/increase in muscle mass; 2

(b) enhances performance/strength unfairly; 1

(c) mimics the action of progesterone in pregnancy; prevents release of the egg/no ovulation; prevents release of FSH and LH by the pituitary gland; 2

Award [1] each for any two. [5]

17. (i) 2

If peptide bond abbreviated, eg −CO−NH− but structure otherwise correct, award [1].

(ii) condensation and water/H2O; 1 [3]

18. chemical messengers; pituitary (gland)/hypothalamus;

[2]

19.

Award [1] for peptide bonds correctly shown in full and a further [1] if rest of structure correct. If peptide bond abbreviated, eg −CO−NH− but structure otherwise correct, award [1].

[2]

20. (i) (warm with dilute) hydrochloric acid; to hydrolyse protein/to break it down into amino acids/to break the peptide bonds; 2

(ii) (mixture of) amino acids spotted on paper (and known amino acids spotted on paper); water/solvent/eluent flows up/down paper;

amino acids separate because they have different solubilities in water/ solvent/eluent and/or different adsorption on paper; amino acid positions identified/sprayed with ninhydrin/locating agent;

locations compared with known amino acids/Rf values compared; 4 [6]

21. (a) saturated have only single carbon to carbon/C−C bonds/unsaturated have double carbon to carbon/C=C bonds;

Do not award mark if no reference to carbon-carbon bonds. saturated have a straight hydrocarbon chain/unsaturated have a kinked hydrocarbon

chain/OWTTE; 2 Accept bond angle of 109(.5)° in saturated and 120° in unsaturated.

(b) chains pack closer together; stronger intermolecular forces/van der Waals' forces; 2

Accept London forces and dispersion forces in place of van der Waals' forces. Do not accept stronger hydrogen bonding. Award [0] if any reference to breaking carbon-carbon bonds.

[4]

22. (a) (testosterone) testes;

development of male sex organs/characteristics/tissue/muscle/bone growth/anabolic effect;

OR

(oestradiol) ovaries; ovulation/development of female sexual characteristics;

OR

(progesterone) ovaries; prepares uterus for fertilized egg; 2

(b) (adrenaline/epinephrine) adrenal glands; regulates body's preparation for stress/OWTTE;

OR

(thyroxine) thyroid gland; regulates body's metabolism;

OR (Insulin) pancreas/Islets of Langerhans;

Regulation of glucose concentration in bloodstream/regulates blood sugar levels; 2 [4]

23. (a) glycosidic/glucoside/ether; 1

(b)

2 Do not penalize candidates who draw bonds connected to incorrect atoms e.g. −HO instead of −OH.

(c) water/H2O; 1

(d) CH2O; 1 [5]

24. (a) vitamin C/ascorbic acid; four/many OH groups/small proportion of hydrocarbon/can form hydrogen bonds with

water/OWTTE; 2

(b) alkene and hydroxyl/alcohol; 1

(c) needed for uptake of calcium/phosphate; bone problem such as softening/weakness/malformation/rickets; 2

[5]

25. (a) glycosidic/glucoside/ether; 1

(b)

2 Do not penalize candidates who draw bonds connected to the incorrect atom − e.g. −HO instead of −OH

(c) CH2O; 1 [4]

26. (a) vitamin C/ascorbic acid; four/many OH groups/small proportion of hydrocarbon/can form hydrogen bonds with

water/OWTTE; 2

(b) vitamins A/retinol and C/ascorbic acid; 1

(c) needed for uptake of calcium/phosphate; bone problem such as softening/weakness/malformation/rickets; 2

[5]

27. (a) hydrogen bonding; 1

(b) van der Waals’ forces/hydrophobic interactions/London/dispersion forces; ionic bonding/(formation of) salt bridges/electrostatic attractions; covalent bonding/(formation of) disulfide bridges; 2 max

Award [1] each for any two. Do not accept sulfur bridges or hydrogen bonding.

[3]

28. (a) vitamin C function collagen formation/production of connective tissue/enhances absorption of iron (from food)/helps healing of wounds/can prevent bacterial infection/antioxidant/bone or teeth formation; effects of deficiency scorbutus/scurvy; vitamin D function uptake of calcium/phosphorus/bone or teeth formation; effects of deficiency rickets; 4

(b) Any two of the following: providing food rations that are composed of fresh and vitamin- and mineral-rich foods; adding nutrients missing in commonly consumed foods; genetic modification of foods; providing nutritional supplements; 2 max

[6]

29. (a) plant material that is not hydrolysed by enzymes (secreted by the human digestive tract); may be digested by microflora in the gut; 2

(b) Any two of the following: cellulose; hemicellulose; lignin; pectin; 2 max

(c) (may be helpful in the prevention of conditions/health problems such as) Any two of the following: diverticulosis; irritable bowel syndrome; constipation; obesity; Crohn’s disease; haemorrhoids; diabetes mellitus; 2 max

[6] 30. (a) both are polymers of glucose; starch has two forms: amylose a straight chain polymer with 〈 – 1, 4

linkages; and amylopectin a branched polymer with 〈 – 1, 4 and 〈 – 1, 6 linkages; cellulose has ® – 1, 4 linkages; 4

(b) absence of cellulase enzyme; 1 [5]

Topic 8 Questions

1. An aqueous solution of which of the following reacts with magnesium metal?

A. Ammonia

B. Hydrogen chloride

C. Potassium hydroxide

D. Sodium hydrogencarbonate

2. Which of the following is/are formed when a metal oxide reacts with a dilute acid?

I. A metal salt

II. Water

III. Hydrogen gas

A. I only B. I and II only


3. Four aqueous solutions, I, II, III and IV, are listed below.

I. 0.100 mol dm–3 HCl

II. 0.010 mol dm–3 HCl

III. 0.100 mol dm–3 NaOH

IV. 0.010 mol dm–3 NaOH

What is the correct order of increasing pH of these solutions?

A. I, II, III, IV

B. I, II, IV, III

C. II, I, III, IV

D. II, I, IV, III

4. Which substance can be dissolved in water to give a 0.1 mol dm–3 solution with a high pH and a high electrical conductivity?

A. HCl

B. NaCl

C. NH3

D. NaOH

5. The pH of a solution is 2. If its pH is increased to 6, how many times greater is the [H+] of the original solution?

A. 3

B. 4

C. 1000

D. 10 000

6. The pH of solution X is 1 and that of Y is 2. Which statement is correct about the hydrogen ion concentrations in the two solutions?

A. [H+] in X is half that in Y.

B. [H+] in X is twice that in Y.

C. [H+] in X is one tenth of that in Y.

D. [H+] in X is ten times that in Y.

7. Lime was added to a sample of soil and the pH changed from 4 to 6. What was the corresponding change in the hydrogen ion concentration?

A. increased by a factor of 2

B. increased by a factor of 100

C. decreased by a factor of 2

D. decreased by a factor of 100

8. When the following 1.0 mol dm–3 solutions are listed in increasing order of pH (lowest first), what is the correct order?

A. HNO3 < H2 CO3 < NH3 < Ba(OH)2

B. NH3 < Ba (OH)2 < H2 CO3 < HNO3

C. Ba (OH)2 < H2 CO3 < NH3 < HNO3

D. HNO3 < H2 CO3 < Ba (OH)2 < NH3

9. Which change in [H+] causes the biggest increase in pH?

A. A change in [H+(aq)] from 1×10–3 to 1×10–2 mol dm–3

B. A change in [H+(aq)] from 1×10–3 to 1×10–4 mol dm–3

C. A change in [H+(aq)] from 1×10–4 to 1×10–2 mol dm–3

D. A change in [H+(aq)] from 1×10–4 to 1×10–6 mol dm–3

10. Which methods can distinguish between solutions of a strong monoprotic acid and a weak monoprotic acid of the same concentration?

I. Add magnesium to each solution and measure the rate of the formation of gas bubbles. II. Add aqueous sodium hydroxide to each solution and measure the temperature change. III. Use each solution in a circuit with a battery and lamp and see how bright the lamp glows.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

11. Which species are a conjugate pair according to the Brønsted-Lowry theory?

A. CH3COOH and CH3CHO

B. NH3 and BF3

C. H2NO3+ and NO3

–

D. H2SO4 and HSO4–

12. Which is not a strong acid?

A. Nitric acid B. Sulfuric acid

C. Carbonic acid D. Hydrochloric acid

13. Lime is added to a lake to neutralize the effects of acid rain. The pH value of the lake water rises from 4 to 7. What is the change in concentration of H+ ions in the lake water?

A. An increase by a factor of 3

B. An increase by a factor of 1000

C. A decrease by a factor of 3

D. A decrease by a factor of 1000

14. Which is a Brønsted-Lowry acid-base pair?

A. H2O and O2–

B. CH3COOH and CH3COO–

C. NH4+ and NH2

–

D. H2SO4 and SO42–

15. Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH3COOH(aq)) of the same concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which statement is correct?

A. CH3COOH(aq) reacted slower because it has a lower pH than HCl(aq).

B. A smaller volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).

C. A greater volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).

D. The same volume of CO2(g) was produced with both CH3COOH(aq) and HCl(aq).

16. Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH3COOH(aq)) of the same concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which statement is correct?

A. CH3COOH(aq) reacted slower because it has a lower pH than HCl(aq).

B. A smaller volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).

C. A greater volume of CO2(g) was produced with CH3COOH(aq) than with HCl(aq).

D. The same volume of CO2(g) was produced with both CH3COOH(aq) and HCl(aq).

17. Which acids are strong?

I. HCl(aq)

II. HNO3(aq)

III. H2SO4(aq)

A. I and II only B. I and III only


18. The pH of a solution changes from pH = 1 to pH = 3. What happens to the [H+] during this pH change?

A. It increases by a factor of 100.

B. It decreases by a factor of 100.

C. It increases by a factor of 1000.

D. It decreases by a factor of 1000.

19. What is the conjugate base of the HSO4–(aq) ion?

A. H2SO4(aq)

B. SO42–(aq)

C. H2O(l)

D. H3O+(aq)

20. Which species can act as a Lewis acid?

A. BF3

B. OH–

C. H2O

D. NH3

21. Which substance, when dissolved in water, to give a 0.1 mol dm–3 solution, has the highest pH?

A. HCl

B. NaCl

C. NH3

D. NaOH

22. Which methods will distinguish between equimolar solutions of a strong base and a strong acid?

I. Add magnesium to each solution and look for the formation of gas bubbles.

II. Add aqueous sodium hydroxide to each solution and measure the temperature change.

III. Use each solution in a circuit with a battery and lamp and see how bright the lamp glows.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

23. (a) Aqueous XO43– ions form a precipitate with aqueous silver ions, Ag+. Write a balanced equation

for the reaction, including state symbols.

........................................................................................................................................... (2)

(b) When 41.18 cm3 of a solution of aqueous silver ions with a concentration of 0.2040 mol dm–3 is added to a solution of XO4

3– ions, 1.172 g of the precipitate is formed.

(i) Calculate the amount (in moles) of Ag+ ions used in the reaction.

(1)

............................................................................................................................

(ii) Calculate the amount (in moles) of the precipitate formed. (1)

............................................................................................................................

(iii) Calculate the molar mass of the precipitate.

............................................................................................................................

............................................................................................................................

............................................................................................................................

............................................................................................................................ (2)

(iv) Determine the relative atomic mass of X and identify the element.

............................................................................................................................

............................................................................................................................ (2)

(Total 8 marks)

24. (a) (i) A solution of hydrochloric acid has a concentration of 0.10 mol dm–3 and a pH value of 1. The solution is diluted by a factor of 100. Determine the concentration of the acid and the pH value in the diluted solution.

..........................................................................................................................

.......................................................................................................................... (2)

(ii) Explain why 0.10 mol dm–3 ethanoic acid solution and the diluted solution in (a) (i) have similar [H+] values.

..........................................................................................................................

..........................................................................................................................

..........................................................................................................................

.......................................................................................................................... (3)

(b) Suggest one method, other than measuring pH, which could be used to distinguish between solutions of a strong acid and a weak acid of the same concentration. State the expected results.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 7 marks)

25. Define the terms strong acid and weak acid. Using hydrochloric and ethanoic acid as examples, write equations to show the dissociation of each acid in aqueous solution.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (Total 4 marks)

26. (i) Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid of the same concentration. State one similarity and one difference in the observations you could make.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(ii) Write an equation for the reaction between hydrochloric acid and calcium carbonate.

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(iii) Determine the volume of 1.50 mol dm–3 hydrochloric acid that would react with exactly 1.25 g of calcium carbonate.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(iv) Calculate the volume of carbon dioxide, measured at 273 K and 1.01×105 Pa, which would be

produced when 1.25 g of calcium carbonate reacts completely with the hydrochloric acid.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(Total 9 marks)

27. The pH values of solutions of three organic acids of the same concentration were measured.

acid X pH = 5 acid Y pH = 2 acid Z pH = 3

(i) Identify which solution is the least acidic. (1)

(ii) Deduce how the [H+] values compare in solutions of acids Y and Z. (2)

(iii) Arrange the solutions of the three acids in decreasing order of electrical conductivity, starting with the greatest conductivity, giving a reason for your choice.

(2) (Total 5 marks)

28. The equilibrium reached when ethanoic acid is added to water can be represented by the following equation:

CH3COOH(l) + H2O(l) CH3COO–(aq)+H3O+(aq)

Define the terms Brønsted-Lowry acid and Lewis base, and identify two examples of each of these species in the equation.

(Total 4 marks)

29. Identify one example of a strong acid and one example of a weak acid. Outline three different methods to distinguish between equimolar solutions of these acids in the laboratory. State how the results would differ for each acid.

(Total 5 marks)

30. Vinegar has a pH of approximately 3 and some detergents have a pH of approximately 8. State and explain which of these has the higher concentration of H+ and by what factor.

(Total 1 mark)

31. Define the terms Brønsted-Lowry acid and Lewis acid. For each type of acid, identify one example other than water and write an equation to illustrate the definition.

(Total 5 marks)

32. The pH values of three acidic solutions, X, Y and Z, are shown in the following table:

Solution Acid pH

X HCl(aq) 2

Y HCl(aq) 4

Z CH3COOH(aq) 4

(i) Solutions X and Z have the same acid concentration. Explain, by reference to both acids, why they have different pH values.

(2)

(ii) Deduce by what factor the values of [H+] in solutions X and Y differ. (1)

(Total 3 marks)

33. State and explain two methods, other than measuring pH, which could be used to distinguish between 1.0 mol dm–3 solutions of nitric acid and ethanoic acid.

(Total 4 marks)

34. Propanoic acid is classified as a weak acid.

(a) State the meaning of the term weak acid.

....................................................................................................................................

.................................................................................................................................... (1)

(b) State, giving a reason in each case, two methods other than measuring pH, that could be used to distinguish between 0.100 mol dm–3 propanoic acid and 0.100 mol dm–3 nitric acid.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 3 marks) 35. State an equation for the reaction of propanoic acid with water. Identify one conjugate Brønsted-Lowry

pair.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 2 marks)

Topic 8 Answers

1. B

2. B

3. B

4. D

5. D

6. D

7. D

8. A

9. D

10. D

11. D

12. C

13. D

14. B

15. D

16. D

17. D

18. B

19. B

20. A

21. D

22. A

23. (a) 3Ag+(aq) + XO43–(aq) → Ag3XO4(s); states; 2

[1] for balanced equation and [1] for states.

(b) (i) nAg+ = cV = 0.2040 mol dm–3×0.04118 dm3 = 0.008401/8.401×10–3 mol (–1 SF) 1 Ignore units even if wrong, do not award mark unless 4 sig fig.

(ii) = nAg+ = ×0.008401 mol = 0.002800/2.800×10–3 mol 1

ECF from (a) and (b)(i)

(iii) 0.002800 mol weighs 1.172 g 1 mol weighs = 418.6 g mol–1 2 418.6; Accept answer in range 418 to 419. No penalty for too many sig figs. ECF from (b) (ii) g mol–1 Do not accept g.

(iv) (3×107.87) + x + 4 (16.0) = 418.6 therefore, x = 30.99 (accept 31.0/31); P/phosphorous; 2

[8]

24. (a) (i) 0.0010 / 1.0×10−3 (mol dm−3);

pH = 3; 2

(ii) HCl: strong acid/fully dissociated; CH3COOH : weak acid/partially dissociated; HCl less concentrated/CH3COOH more concentrated;

only one molecule in 100 dissociates in ethanoic acid so [H+] 1/100/OWTTE 3

(b) measure electrical conductivity; strong acids are good conductors/weak acids are poor conductors;

OR

react with magnesium or a named active metal/(metal) carbonate; hydrogen carbonate/bicarbonate;

strong acids have a faster reaction/more gas bubbles (per unit time) /more heat produced/weak acids have a slower reaction/less gas bubbles (per unit time)/less heat produced; 2

titration curves: namely strong acid and strong base will have an equivalence point pH of 7 and a weak acid and strong base will have an equivalence point pH of >7. OR

temperature change: on neutralization for temperature change: namely, neutralization (H+ + OH−) is exothermic, weak acid is partially dissociated so some energy used up in dissociation of weak acid − net result, weak acid would produce less energy/less temperature increase compared to neutralization of strong acid.

[7]

25. strong acid completely dissociated/ionized; weak acid only partially dissociated/ionized;

HCl(aq) → H+(aq) + Cl+(aq); + CH3COOH(aq) CH3COO–(aq) + H+(aq); 4 Insist on both arrows as shown, state symbols not needed. Also accept H2O(1) and H3O+(aq) in equations

[4]

26. (i) bubbling/effervescence/dissolving of CaCO3/gas given off (do not accept CO2 produced); more vigorous reaction with HCl/OWTTE; 2

(ii) 2HCl(aq) + CaCO3(s) → CaCl2(aq) + CO2(g) + H2O(1); 2 [1] for correct formulas, [1] for balanced, state symbols not essential.

(iii) amount of CaCO3 = (no penalty for use of 100); amount of HCl = 2×0.0125 = 0.0250 mol (allow ECF); volume of HCl = 0.0167 dm3/16.7 cm3 (allow ECF); 3

(iv) 1:1 ratio of CaCO3 to CO2 /use 0.0125 moles CO2 (allow ECF); (0.0125×22.4) = 0.28 dm3/280 cm3/2.8×10–4 m3 (allow ECF); 1

Accept calculation using pV=nRT. [9]

27. (i) X; 1

(ii) greater in Y/smaller in Z; by a factor of 10; 2

(iii) Y > Z > X; most ions/greatest concentration of ions in Y/OWTTE; 2

[5]

28. Brønsted-Lowry acid proton donor/OWTTE; CH3COOH and H3O+;

Lewis base electron pair donor/OWTTE; H2O and CH3COO−; 4

[4]

29. HCl/H2SO4/HNO3/any strong acid; CH3COOH/H2CO3/any weak acid; Measure pH − the strong acid has the lower pH;

Accept universal indicator and two correct colours. Measure (electrical) conductivity − this is greater for the stronger acid; Add magnesium/carbonate − more gas bubbles with the stronger acid/Mg or carbonate would disappear faster with stronger acid;

[5]

30. vinegar and factor of 105; [1]

31. Brønsted-Lowry acid a proton donor;

Lewis acid electron pair acceptor;

Brønsted-Lowry acid Any suitable equation;

Lewis acid − BF3/AlCl3/transition metal ions that form complex ion with ligands;

For example BF3 + NH3 → BF3NH3/Cu2+ + 4NH3 → [Cu(NH3)4]2+/AlCl3 + Cl− → ; 5 Or any suitable equation.

[5]

32. (i) HCl/X is strong and CH3COOH/Z is weak; HCl/X is fully dissociated and CH3COOH is slightly dissociated; [H+ ] is greater in HCl/X than in CH3COOH/Z; 2

Any two for [1] each.

(ii) a factor of 100; 1 [3]

33. conductivity; nitric acid will contain more ions and have a higher conductivity/ethanoic acid will have fewer ions and have a lower conductivity;

rate of reaction with metal/carbonate/hydrogencarbonate; nitric acid will react more rapidly/produce bubbles faster/ethanoic acid will react less rapidly/produce bubbles more slowly;

reaction with alkali; temperature change will be less for ethanoic acid; 4

Accept any two methods and explanations from above. [4]

34. (a) an acid that partially dissociates/ionizes/doesn’t fully dissociate/ionize; 1

(b) conductivity - propanoic acid will be lower because lower ion concentration /less dissociated; reaction with metal/metal carbonate/metal hydrogencarbonate - propanoic acid will react

slower/less vigorously because lower [H+]/less dissociated; reaction with alkali - temperature change will be less for propanoic acid because lower [H+]/less

dissociated; 2 Award [1] mark each for two.

[3] 35. CH3CH2COOH + H2O CH3CH2COO− + H3O+/CH3CH2COOH CH3CH2COO− + H+;

required for mark. CH3CH2COOH and CH3CH2COO−/H3O+ and H2O; 2

[2] Topic1 Questions and Answers

1. What amount of oxygen, O2, (in moles) contains 1.8×1022 molecules?

A. 0.0030

B. 0.030

C. 0.30


2. Which compound has the empirical formula with the greatest mass?

A. C2H6

B. C4H10

C. C5H10


3. __C2H2(g) + __O2(g) → __ CO2(g) + __ H2O(g)

When the equation above is balanced, what is the coefficient for oxygen?

A. 2

B. 3

C. 4

D. 5 (Total 1 mark)

4. 3.0 dm3 of sulfur dioxide is reacted with 2.0 dm3 of oxygen according to the equation below.

2SO2(g) + O2(g) → 2SO3(g)

What volume of sulfur trioxide (in dm3) is formed? (Assume the reaction goes to completion and all gases are measured at the same temperature and pressure.)

A. 5.0

B. 4.0

C. 3.0


5. What will happen to the volume of a fixed mass of gas when its pressure and temperature (in Kelvin) are both doubled?

A. It will not change.

B. It will increase.

C. It will decrease.

D. The change cannot be predicted. (Total 1 mark)

6. What amount (in moles) is present in 2.0 g of sodium hydroxide, NaOH?

A. 0.050

B. 0.10

C. 20


7. A hydrocarbon contains 90% by mass of carbon. What is its empirical formula?

A. CH2

B. C3H4

C. C7H10


8. Copper can react with nitric acid as follows.

3Cu +_HNO3 → _Cu(NO3)2 +_H2O + _NO

What is the coefficient for HNO3 when the equation is balanced?

A. 4

B. 6

C. 8


9. Lithium hydroxide reacts with carbon dioxide as follows.

2LiOH + CO2 → Li2 CO3 + H2O

What mass (in grams) of lithium hydroxide is needed to react with 11 g of carbon dioxide?

A. 6

B. 12

C. 24


10. Which change in conditions would increase the volume of a fixed mass of gas?

Pressure /kPa Temperature /K

A. Doubled Doubled

B. Halved Halved

C. Doubled Halved

D. Halved Doubled (Total 1 mark)

11. How many hydrogen atoms are contained in one mole of ethanol, C2H5OH?

A. 5

B. 6

C. 1.0×1023

D. 3.6×1024

(Total 1 mark)

12. The percentage by mass of the elements in a compound is

C = 72%, H = 12%, O = 16%.

What is the mole ratio of C:H in the empirical formula of this compound?

A. 1 : 1

B. 1 : 2

C. 1 : 6

D. 6 : 1 (Total 1 mark)

13. What is the coefficient for O2(g) when the equation below is balanced?

__C3H8(g) + __O2(g) → __CO2(g) + __H2O(g)

A. 2

B. 3

C. 5

D. 7 (Total 1 mark)

14. What amount of NaCl (in moles) is required to prepare 250 cm3 of a 0.200 mol dm–3 solution?

A. 50.0

B. 1.25

C. 0.800


15. For which set of conditions does a fixed mass of an ideal gas have the greatest volume?

Temperature Pressure

A. low low

B. low high

C. high high

D. high low (Total 1 mark)

16. Which of the following contains the greatest number of molecules?

A. 1 g of CH3Cl

B. 1 g of CH2Cl2

C. 1 g of CHCl3

D. 1 g of CCl4 (Total 1 mark

17. Which of the following compounds has/have the empirical formula CH2O?

I. CH3COOH

II. C6H12O6

III. C12H22O11

A. II only

B. III only

C. I and II only


18. Assuming complete reaction, what volume of 0.200 mol dm–3 HCl(aq) is required to neutralize 25.0 cm3 of 0.200 mol dm–3 Ba(OH)2(aq)?

A. 12.5 cm3

B. 25.0 cm3

C. 50.0 cm3

D. 75.0 cm3 (Total 1 mark)

19. Under what conditions would one mole of methane gas, CH4, occupy the smallest volume?

A. 273 K and 1.01×105 Pa

B. 273 K and 2.02×105 Pa

C. 546 K and 1.01×105 Pa

D. 546 K and 2.02×105 Pa (Total 1 mark)

20. The temperature in Kelvin of 2.0 dm3 of an ideal gas is doubled and its pressure is increased by a factor of four. What is the final volume of the gas?

A. 1.0 dm3

B. 2.0 dm3

C. 3.0 dm3

D. 4.0 dm3 (Total 1 mark)

21. Which is a correct definition of the term empirical formula?

A. formula showing the numbers of atoms present in a compound

B. formula showing the numbers of elements present in a compound

C. formula showing the actual numbers of atoms of each element in a compound

D. formula showing the simplest ratio of numbers of atoms of each element in a compound (Total 1 mark)

22. The reaction of ethanal and oxygen can be represented by the unbalanced equation below.

__ CH3CHO + __ O2 → __ CO2 + __ H2O

When the equation is balanced using the smallest possible integers, what is the coefficient for O2?

A. 3

B. 4

C. 5

D. 6 (Total 1 mark)

23. The equation for the complete combustion of butane is

2C4H10 + 13O2 → 8CO2 + 10H2O

What is the amount (in mol) of carbon dioxide formed by the complete combustion of three moles of butane?

A. 4

B. 8

C. 12


24. Which solution contains the greatest amount (in mol) of solute?

A. 10.0 cm3 of 0.500 mol dm–3 NaCl

B. 20.0 cm3 of 0.400 mol dm–3 NaCl

C. 30.0 cm3 of 0.300 mol dm–3 NaCl

D. 40.0 cm3 of 0.200 mol dm–3 NaCl (Total 1 mark)

25. A fixed mass of an ideal gas has a volume of 800 cm3 under certain conditions. The pressure (in kPa) and temperature (in K) are both doubled. What is the volume of the gas after these changes with other conditions remaining the same?

A. 200 cm3

B. 800 cm3

C. 1600 cm3

D. 3200 cm3

(Total 1 mark)

26. The complete oxidation of propane produces carbon dioxide and water as shown below.

C3H8 + __O2 →__CO2 + __H2O

What is the total of the coefficients for the products in the balanced equation for 1 mole of propane?

A. 6

B. 7

C. 12


27. The relative molecular mass (Mr) of a compound is 60. Which formulas are possible for this compound?

I. CH3CH2CH2NH2

II. CH3CH2CH2OH

III. CH3CH(OH)CH3

A. I and II only

B. I and III only

C. II and III only


28. Which sample has the least number of atoms?

A. 1 mol of H2SO4

B. 1 mol of CH3COOH

C. 2 mol of H2O2

D. 2 mol of NH3 (Total 1 mark)

29. Avogadro’s constant has the same value as the number of

A. molecules in 1 mol of solid iodine.

B. atoms in 1 mol of chlorine gas.

C. ions in 1 mol of solid potassium bromide.

D. protons in 1 mol of helium gas. (Total 1 mark)

30. What is the total of all the coefficients in the balanced equation for the reduction of 1 mol of MnO4–?

__ MnO4– +__H+ + __ e– →__Mn2+ + __H2O

A. 5

B. 9

C. 17


31. Which contains the same number of ions as the value of Avogadro’s constant?

A. 0.5 mol NaCl

B. 0.5 mol MgCl2

C. 1.0 mol Na2O

D. 1.0 mol MgO (Total 1 mark)

32. A reaction occurring in the extraction of lead from its ore can be represented by this unbalanced equation:

__ PbS + __O2 → __ PbO + __ SO2

When the equation is balanced using the smallest possible whole numbers, what is the coefficient for O2?

A. 1

B. 2

C. 3

D. 4 (Total 1 mark

33. The equation for a reaction occurring in the synthesis of methanol is

CO2 + 3H2 → CH3OH + H2O

What is the maximum amount of methanol that can be formed from 2 mol of carbon dioxide and 3 mol of hydrogen?

A. 1 mol

B. 2 mol

C. 3 mol

D. 5 mol (Total 1 mark)

34. Which solution contains 0.1 mol of sodium hydroxide?

A. 1 cm3 of 0.1 mol dm–3 NaOH

B. 10 cm3 of 0.1 mol dm–3 NaOH

C. 100 cm3 of 1.0 mol dm–3 NaOH

D. 1000 cm3 of 1.0 mol dm–3 NaOH (Total 1 mark)

35. A cylinder of gas is at a pressure of 40 kPa. The volume and temperature (in K) are both doubled. What is the pressure of the gas after these changes?

A. 10 kPa

B. 20 kPa

C. 40 kPa

D. 80 kPa (Total 1 mark)

36. Which of the following quantities has units?

A. Relative atomic mass

B. Relative molecular mass

C. Molar mass

D. Mass number (Total 1 mark)

37. The empirical formula of a compound is C2H4O. Which molecular formulas are possible for this compound?

I. CH3COOH

II. CH3CH2CH2COOH

III. CH3COOCH2CH3

A. I and II only

B. I and III only

C. II and III only


38. Calcium carbonate decomposes on heating as shown below.

CaCO3 → CaO + CO2

When 50 g of calcium carbonate are decomposed, 7 g of calcium oxide are formed. What is the percentage yield of calcium oxide?

A. 7%

B. 25%

C. 50%

D. 75% (Total 1 mark)

39. Sodium reacts with water as shown below.

__ Na + __ H2O → __ NaOH + __ H2

What is the total of all the coefficients when the equation is balanced using the smallest possible whole numbers?

A. 3

B. 4

C. 6

D. 7

(Total 1 mark)

40. What is the total number of ions present in the formula, Al2(SO4)3?

A. 2

B. 3

C. 5

D. 6 (Total 1 mark)

41. A 4 g sample of sodium hydroxide, NaOH, is dissolved in water and made up to 500 cm3 of aqueous solution. What is the concentration of the resulting solution?

A. 0.1 mol dm–3

B. 0.2 mol dm–3

C. 0.5 mol dm–3

D. 1.0 mol dm–3 (Total 1 mark)

42. Methane, CH4, burns in oxygen gas to form carbon dioxide and water. How many moles of carbon dioxide will be formed from 8.0 g of methane?

A. 0.25

B. 0.50

C. 1.0


43. What is the empirical formula of a compound containing 50% by mass of element X (Ar = 20) and 50% by mass of element Y (Ar = 25)?

A. XY

B. X3Y2

C. X4Y5

D. X5Y4 (Total 1 mark)

44. Assuming complete reaction, what volume of 0.200 mol dm–3 potassium hydroxide solution (KOH(aq)), is required to neutralize 25.0 cm3 of 0.200 mol dm–3 aqueous sulfuric acid, (H2SO4(aq))?

A. 12.5 cm3

B. 25.0 cm3

C. 50.0 cm3

D. 75.0 cm3 (Total 1 mark)

45. Consider the following reaction.

N2(g) + 3H2(g) 2NH3(g)

If the reaction is made to go to completion, what volume of ammonia (in dm3) can be prepared from 25 dm3 of nitrogen and 60 dm3 of hydrogen? All volumes are measured at the same temperature and pressure.

A. 40

B. 50

C. 85


46. The temperature in Kelvin of 1.0 dm3 of an ideal gas is doubled and its pressure is tripled. What is the final volume of the gas in dm3?

A.

B.

C.

D. (Total 1 mark)

47. On complete combustion, a sample of a hydrocarbon compound produces 1.5 mol of carbon dioxide and 2.0 mol of water. What is the molecular formula of this hydrocarbon?

A. C2H2

B. C2H4

C. C3H4


48. When excess BaCl2(aq) was added to a sample of Fe(NH4)2(SO4)2(aq) to determine the amount in moles of sulfate present, 5.02×10–3 mol of BaSO4 was obtained. How many moles of sulfate ions and iron ions were in the sample of Fe(NH4)2(SO4)2?

Amount of sulfate ions / moles Amount of iron ions / moles

A. 5.02×10–3 2.51×10–3

B. 10.04×10–3 5.02×10–3

C. 2.51×10–3 5.02×10–3

D. 10.04×10–3 2.51×10–3 (Total 1 mark)

49. What volume of 0.500 mol dm–3 sulfuric acid solution is required to react completely with 10.0 g of calcium carbonate according to the equation below?

CaCO3(s) + H2SO4(aq) → CaSO4(aq) + H2O(l) + CO2(g)

A. 100 cm3

B. 200 cm3

C. 300 cm3

D. 400 cm3 (Total 1 mark)

50. Which expression gives the amount (in mol) of a substance, if the mass is given in grams?

A.

B.

C.

D. mass × molar mass (Total 1 mark)

51. What is the total number of atoms in 0.20 mol of propanone, CH3COCH3?

A. 1.2×1022

B. 6.0×1023

C. 1.2×1024

D. 6.0×1024 (Total 1 mark)

52. When the equation below is balanced for 1 mol of C3H4, what is the coefficient for O2?

C3C4 + _O2 → _CO2 + _H2O

A. 2

B. 3

C. 4

D. 5 (Total 1 mark)

53. Ethyne, C2H2, reacts with oxygen according to the equation below. What volume of oxygen (in dm3) reacts with 0.40 dm3 of C2H2?

2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g)

A. 0.40

B. 0.80

C. 1.0


54. Ethyne, C2H2, reacts with oxygen according to the equation below. What volume of oxygen (in dm3) reacts with 0.40 dm3 of C2H2?

2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g)

A. 0.40

B. 0.80

C. 1.0


55. What is the coefficient for H+ when the redox equation below is balanced?

__Ag(s) + __NO3–(aq) +__H+(aq) → __Ag+(aq) + __NO(g) + __H2O(l)

A. 1

B. 2

C. 3

D. 4 (Total 1 mark)

56. How many hydrogen atoms are in one mole of ethanol, C2H5OH?

A. 1.00×1023

B. 3.61×1024

C. 5.00


57. What is the coefficient for H2SO4(aq) when the following equation is balanced, using the smallest possible integers?

__Mg3N2(s) + __H2SO4(aq) → __MgSO4(aq) + __(NH4)2SO4(aq)

A. 1

B. 3

C. 4

D. 7 (Total 1 mark)

58. Air bags in cars inflate when sodium azide decomposes to form sodium and nitrogen:

2NaN3(s) → 2Na(s) + 3N2(g)

Calculate the amount, in moles, of nitrogen gas produced by the decomposition of 2.52 mol of NaN3(s).

A. 1.68

B. 2.52

C. 3.78


59. What volume, in cm3, of 0.200 mol dm–3 HCl(aq) is required to neutralize 25.0 cm3 of 0.200 mol dm–3 Ba(OH)2(aq)?

A. 12.5

B. 25.0

C. 50.0


60. The relative molecular mass of aluminium chloride is 267 and its composition by mass is 20.3% Al and 79.7% chlorine. Determine the empirical and molecular formulas of aluminium chloride.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (Total 4 marks)

61. Sodium reacts with water as follows.

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

1.15 g of sodium is allowed to react completely with water. The resulting solution is diluted to 250 cm3. Calculate the concentration, in mol dm–3, of the resulting sodium hydroxide solution.

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

………………………………………………………………………………………………… (Total 3 marks)

62. (i) Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid of the same concentration. State one similarity and one difference in the observations you could make.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(ii) Write an equation for the reaction between hydrochloric acid and calcium carbonate.

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

(iii) Determine the volume of 1.50 mol dm–3 hydrochloric acid that would react with exactly 1.25 g of calcium carbonate.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

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………………………………………………………………………………………….. (3)

(iv) Calculate the volume of carbon dioxide, measured at 273 K and 1.01×105 Pa, which would be produced when 1.25 g of calcium carbonate reacts completely with the hydrochloric acid.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

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………………………………………………………………………………………….. (2)

(Total 9 marks)

63. An organic compound, A, containing only the elements carbon, hydrogen and oxygen was analysed.

(a) A was found to contain 54.5% C and 9.1% H by mass, the remainder being oxygen. Determine the empirical formula of the compound.

(3)

(b) A 0.230 g sample of A, when vaporized, had a volume of 0.0785 dm3 at 95°C and 102 kPa. Determine the relative molecular mass of A.

(3)

(c) Determine the molecular formula of A using your answers from parts (a) and (b).

...................................................................................................................................

................................................................................................................................... (1)

(Total 7 marks)

64. An organic compound A contains 62.0% by mass of carbon, 24.1% by mass of nitrogen, the remainder being hydrogen.

(i) Determine the percentage by mass of hydrogen and the empirical formula of A.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (3)

(ii) Define the term relative molecular mass.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(iii) The relative molecular mass of A is 116. Determine the molecular formula of A.

...................................................................................................................................

................................................................................................................................... (1)

(Total 6 marks)

65. An organic compound A contains 62.0% by mass of carbon, 24.1% by mass of nitrogen, the remainder being hydrogen.

(i) Determine the percentage by mass of hydrogen and the empirical formula of A.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (3)

(ii) Define the term relative molecular mass.

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(iii) The relative molecular mass of A is 116. Determine the molecular formula of A.

...................................................................................................................................

................................................................................................................................... (1)

(Total 6 marks)

66. Propane and oxygen react according to the following equation.

C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g)

Calculate the volume of carbon dioxide and water vapour produced and the volume of oxygen remaining, when 20.0 dm3 of propane reacts with 120.0 dm3 of oxygen. All gas volumes are measured at the same temperature and pressure.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 3 marks)

67. State and explain what would happen to the pressure of a given mass of gas when its absolute temperature and volume are both doubled.

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

..............................................................................................................................................

.............................................................................................................................................. (Total 3 marks)

68. (i) Crocetin consists of the elements carbon, hydrogen and oxygen. Determine the empirical formula of crocetin, if 1.00 g of crocetin forms 2.68 g of carbon dioxide and 0.657 g of water when it undergoes complete combustion.

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...................................................................................................................................

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...................................................................................................................................

................................................................................................................................... (6)

(ii) Determine the molecular formula of crocetin given that 0.300 mole of crocetin has a mass of 98.5 g

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...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... (2)

(Total 8 marks)

69. A solution containing ammonia requires 25.0 cm3 of 0.100 mol dm–3 hydrochloric acid to reach the equivalence point of a titration.

(i) Write an equation for the reaction of ammonia with hydrochloric acid (1)

(ii) Calculate the amount (in mol) of hydrochloric acid and ammonia that react. (2)

(iii) Calculate the mass of ammonia in the solution. (2)

(Total 5 marks)

70. A toxic gas, A, consists of 53.8% nitrogen and 46.2% carbon by mass. At 273 K and 1.01×105 Pa, 1.048 g of A occupies 462 cm3. Determine the empirical formula of A. Calculate the molar mass of the compound and determine its molecular structure.

(Total 3 marks)

71. An oxide of copper was reduced in a stream of hydrogen as shown below.

After heating, the stream of hydrogen gas was maintained until the apparatus had cooled.

The following results were obtained.

Mass of empty dish = 13.80 g Mass of dish and contents before heating = 21.75 g Mass of dish and contents after heating and leaving to cool = 20.15 g

(a) Explain why the stream of hydrogen gas was maintained until the apparatus cooled.

.....................................................................................................................................

..................................................................................................................................... (1)

(b) Calculate the empirical formula of the oxide of copper using the data above, assuming complete reduction of the oxide.

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (3)

(c) Write an equation for the reaction that occurred.

..................................................................................................................................... (1)

(d) State two changes that would be observed inside the tube as it was heated.

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.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... (2)

(Total 7 marks)

72. Copper metal may be produced by the reaction of copper(I) oxide and copper(I) sulfide according to the below equation.

2Cu2O + Cu2S → 6Cu + SO2

A mixture of 10.0 kg of copper(I) oxide and 5.00 kg of copper(I) sulfide was heated until no further reaction occurred.

(a) Determine the limiting reagent in this reaction, showing your working.

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....................................................................................................................................

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....................................................................................................................................

.................................................................................................................................... (3)

(b) Calculate the maximum mass of copper that could be obtained from these masses of reactants.

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....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(Total 5 marks)

73. The reaction below represents the reduction of iron ore to produce iron.

2Fe2O3 + 3C → 4Fe + 3CO2

A mixture of 30 kg of Fe2O3 and 5.0 kg of C was heated until no further reaction occurred. Calculate the maximum mass of iron that can be obtained from these masses of reactants.

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.............................................................................................................................................. (Total 5 marks)

74. 0.502 g of an alkali metal sulfate is dissolved in water and excess barium chloride solution, BaCl2(aq) is added to precipitate all the sulfate ions as barium sulfate, BaSO4(s). The precipitate is filtered and dried and weighs 0.672 g.

(a) Calculate the amount (in mol) of barium sulfate formed.

....................................................................................................................................

....................................................................................................................................

....................................................................................................................................

.................................................................................................................................... (2)

(b) Determine the amount (in mol) of the alkali metal sulfate present.

.................................................................................................................................... (1)

(c) Determine the molar mass of the alkali metal sulfate and state its units.

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....................................................................................................................................

.................................................................................................................................... (2)

(d) Deduce the identity of the alkali metal, showing your workings.

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....................................................................................................................................

.................................................................................................................................... (2)

(e) Write an equation for the precipitation reaction, including state symbols.

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(2) (Total 9 marks)

75. 0.600 mol of aluminium hydroxide is mixed with 0.600 mol of sulfuric acid, and the following reaction

occurs:

2Al(OH)3(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 6H2O(l)

(a) Determine the limiting reactant.

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..................................................................................................................................... (2)

(b) Calculate the mass of Al2(SO4)3 produced.

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.....................................................................................................................................

..................................................................................................................................... (2)

(c) Determine the amount (in mol) of excess reactant that remains.

.....................................................................................................................................

..................................................................................................................................... (1)

(d) Define a Brønsted-Lowry acid and a Lewis base.

Brønsted-Lowry acid

.....................................................................................................................................

Lewis base

.....................................................................................................................................

(2)

(e) H2SO4(aq) is a strong acid. State the name and the formula of any weak acid.

.....................................................................................................................................

..................................................................................................................................... (1)

(Total 8 mark

1. B

2. B

3. D

4. C

5. A

6. A

7. B

8. C

9. B

10. D 11. D

12. B

13. C

14. D

15. D"

16. A

17. C

18. C

19. B

20. A 21. D

22. C

23. C

24. C

25. B

26. B

27. C

28. A

29. A

30. D

31. A

32. C

33. A

34. C

35. C

36. C

37. C

38. B

39. D

40. C

41. B

42. B

43. D

44. C

45. A

46. B

47. D

48. A

49. B

50. A

51. C

52. C

53. C

54. C

55. D

56. B

57. C

58. C

59. C

60. AlCl or similar working (no penalty for use of 27 or 35.5);

empirical formula AlCl3;

molecular formula: n = = 2;

Al2Cl6; Full credit can be obtained if the calculations are carried out by another valid method. Two correct formulas but no valid method scores [2 max].

[4]

61. moles of Na = = 0.05;

moles of NaOH = 0.05; Accept “same as moles of Na”

concentration = = 0.20 (mol dm–3) 3 Allow ECF from moles of NaOH

[3]

62. (i) bubbling/effervescence/dissolving of CaCO3/gas given off (do not accept CO2 produced); more vigorous reaction with HCl/OWTTE; 2

(ii) 2HCl(aq) + CaCO3(s) → CaCl2(aq) + CO2(g) + H2O(1); 2 [1] for correct formulas, [1] for balanced, state symbols not

(iii) amount of CaCO3 = (no penalty for use of 100); amount of HCl = 2×0.0125 = 0.0250 mol (allow ECF); volume of HCl = 0.0167 dm3/16.7 cm3 (allow ECF); 3

(iv) 1:1 ratio of CaCO3 to CO2 /use 0.0125 moles CO2 (allow ECF); (0.0125×22.4) = 0.28 dm3/280

cm3/2.8×10–4 m3 (allow ECF); 1 Accept calculation using pV=nRT.

[9]

63. (a) % of oxygen = 36.4;

Do not penalize if 12, 1 and 16 are used.

C2H4O; 3 If atomic numbers or incorrect Ar values used, only first mark can be scored. Award [3] for correct formula without working.

(b) pV = nRT/pV = /correct rearrangement;

Award [1] for 368 even if incorrect expression given.

Mr = 87.8; 3 Accept answer in range 87.8 to 88. Do not allow ECF. Award [3] for correct final answer

(c) C4H8O2; 1 Answer does not need to show working to receive the mark. Do not allow ECF.

[7]

64. (i) C N H

Award [2] for above. No penalty for use of whole number atomic masses. If atomic numbers used then only mark for % of H can be awarded. If H % and calculation missing, award [1], and last mark cannot be scored. If H % calculation incorrect apply ECF.

C3NH8; 3 Correct empirical formula scores [3].

(ii) the average mass of a molecule;

compared to 1/12 of (the mass of) one atom of 12C/compared to C-12 taken as 12;

OR

2

Award [2] for the equation above.

(iii) C6N2H16; 1 [6]

65. (i) C N H

Award [2] for above. No penalty for use of whole number atomic masses. If atomic numbers used then only mark for % of H can be awarded. If H % and calculation missing, award [1], and last mark cannot be scored. If H % calculation incorrect apply ECF.

C3NH8; 3 Correct empirical formula scores [3].

(ii) the average mass of a molecule; compared to 1/12 of (the mass of) one atom of 12C/compared to C-12 taken as 12;

OR

2 Award [2] for the equation above.

(iii) C6N2H16; 1 [6]

66. 60.0 dm3 CO2;

80.0 dm3 H2O;

20.0 dm3 O2; 3 Apply −1(U).

[3]

67. overall there will be no change to the pressure; double absolute temperature and the pressure doubles; double volume and the pressure halves;

Apply ECF if points 2 and 3 are incorrect.

OR Use PV = nRT, Since n and R are constant; V and T are both doubled; P will remain unchanged;

OR OWTTE for mathematical interpretation

e.g. T 〈 P, therefore 2P; V 〈 1/P, therefore ½P; No change to P, ½P×2P = P; 3

68. (i) n(C)(= n(CO2) = 2.68 g÷44.01 g mol−1) = 0.0609 mol; n(H)(= 2×n(H2O) = 0.657 g÷18.02 g mol−1) = 0.0729 mol; m(C) = 0.0609 mol×12.01 g mol−1 = 0.731 g and m(H) = 0.0729 mol×1.01 g mol−1 = 0.0736 g; m(O) = (1.00 − 0.731 − 0.0736)g = 0.195g;

n(C) n(H) n(O) 0.0609 0.0730 0.195 16.00 0.0609 0.0730 0.0122 0.0609 0.0730 0.0122 0.0122 0.0122 0.0122

4.99 5.98 1.00;

empirical formula: C5H6O; 6 For C5H6 award [4 max]. Steps used to arrive at the correct amounts (in moles) are required for full marks.

(ii) M(crocetin) = 98.5 g÷0.300 mol = 328 (g mol−1);

molecular formula: C20H24O4; 2 ECF from (i).

[8]

69. (i) NH3(aq) + HCl(aq) → NH4Cl(aq)+; 1 States not required for mark

(ii) n(HCl) = cV = 0.100 mol dm−3×0.0250 dm3 = 0.00250 mol; n(NH3) = n(HCl) = 0.00250 mol; 2 ECF

(iii) (M (NH3) = 14.01 + 3(1.01) =) 17.04/17.0 (g mol−1); m(NH3) = 0.00250 mol×17.04g mol −1 = 0.0426g/0.0425g; 2

ECF [5]

70. empirical formula = CN; Working must be shown to get point.

Mr = 51.9 (g mol–1); :NC(CN:; 3 [3]

71. (a) to prevent (re)oxidation of the copper/OWTTE; 1

(b) number of moles of oxygen = = 0.10; number of moles of copper = = 0.10; empirical formula = Cu (0.10) : O (0.10) = CuO; 3

Allow ECF. Award [1] for CuO with no working. Alternate solution

= 79.8% = 20.2%

= 1.25 = 1.29

(c) H2 + CuO → Cu + H2O; 1 Allow ECF.

(d) (black copper oxide) solid turns red/brown; condensation/water vapour (on sides of test tube); 2

Accept change colour. Do not accept reduction of sample size. [7]

72. (a) n(Cu2O) = 10.0×103÷143.1 = 69.9 mol; n(Cu2S) = 5.00×103÷159.16 = 31.4 mol;

Penalise failure to convert kg → g once only.

Cu2S is the limiting reagent; 3 ECF from above answers.

(b) n(Cu) = 6×n(Cu2S) = 6×31.4 = 188 mol; m(Cu) = 188×63.55 = 11900 − 12000 g/11.9 − 12.0 kg; 2

If Cu2O given in (a), allow 3×n(Cu2O) and 3×n(Cu2O)×63.55. Allow ECF from (a).

[5]

73. n(Fe2O3) = 30×103÷159.7/n(Fe2O3) = 188 mol;

n(C) = 5.0×103÷12.01/n(C) = 416 mol;

Fe2O3 is the limiting reagent or implicit in calculation;

n(Fe) = 2×n(Fe2O3) = 2×188 = 376 mol;

m(Fe) = 376×55.85 = 21 kg; Accept 2 sig. fig. or 3 sig. fig., otherwise use − 1(SF). Correct final answers score [5]. Allow ECF.

[5]

74. (a) M(BaSO4) (= 137.34 + 32.06 + 4(16.00)) = 233.40 (g mol−1);

Accept 233.4 but not 233

n(BaSO4) = 0.00288 / 2.88×10−3(mol); 2 ECF from M value

(b) n (alkali metal sulfate) = 0.00288 / 2.88×10−3(mol); 1 ECF

(c) 174.31 / 174.3 / 174; ECF

units: g mol−1; 2

(d) (2(Ar) + 32 + 4(16) = 174, thus) Ar = 39 / Ar = =39; Accept answer between 38.9 and 39.2 ECF potassium/K; ECF from Ar value

2

(e) K2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2KCl(aq) 2 Award [1] for balanced equation and [1] for state symbols ECF if another alkali metal arrived at in (d) Accept net ionic equation If no answer arrived at in (d), but correct equation given involving any alkali metal, then award [1 max]

[9] 75. (a) 0.600 mol Al(OH)3 ≡ (1.5)(0.600) mol H2SO4/0.900 mol H2SO4 needed, but only 0.600 mol

used; H2SO4 limiting reactant; 2 Some working must be shown in order to score the second point.

(b) 0.200 mol Al2(SO4)3; 68.4(g); 2 Penalize incorrect units.

(c) 0.200 mol; 1 Use ECF from (a).

(d) A Brønsted-Lowry acid is a proton/H+ donor; A Lewis base is an electron-pair donor; 2

(e) H2CO3 and carbonic acid/CH3COOH and ethanoic acid; 1 Accept any other weak acid and correct formula.

[8] !

chemistry sl c. ii and iii only topic 5 questions - ib relics · chemistry sl topic 5 questions 1....

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