AS1 Chemistry Papers: 2006 - 2014 & Mark …...10 The enthalpy of neutralisation when an acid reacts with an alkali is, by definition, the number of kilojoules released by A the formation

TIME

1 hour 30 minutes.

INSTRUCTIONS TO CANDIDATES

Write your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all sixteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all six questions in Section B. Write your answers in the spaces provided in this question paper.

INFORMATION FOR CANDIDATES

The total mark for this paper is 100.Quality of written communication will be assessed in question 14(d)(iii).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.

ADVANCED SUBSIDIARY (AS)General Certificate of Education

2006

Chemistry

Assessment Unit AS 1assessing

Module 1: General Chemistry

WEDNESDAY 7 JUNE, MORNING

BP

2

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6EA

AS1 Chemistry Papers: 2006 - 2014& Mark Schemes2006-09: Old-Old SpecJune 2009-14: Old Spec

395 Pages in total.

http://code-industry.net/




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Section A

For each of the questions only one of the lettered responses (A – D) is correct.

Select the correct response in each case and mark its code letter by connecting the dots as illustrated on the answer sheet.

1 20 cm3 of 0.3 mol dm–3 potassium hydroxide solution is exactly neutralised by

A 10 cm3 of 0.3 mol dm–3 sulphuric acid B 10 cm3 of 0.6 mol dm–3 sulphuric acid C 20 cm3 of 0.3 mol dm–3 sulphuric acid D 20 cm3 of 0.6 mol dm–3 sulphuric acid

2 Which one of the following electron configurations has two unpaired electrons?

A 1s2 2s1

B 1s2 2s2 2p3

C 1s2 2s2 2p4

D 1s2 2s2 2p6 3s2 3p5

3 Which one of the following shows the trend in electronegativity values of the elements in the Periodic Table?

Across a period Down a group

A decrease decrease

B decrease increase

C increase decrease

D increase increase








4 Which one of the following sodium compounds produces a gas when treated with dilute sulphuric acid?

A sodium carbonate B sodium chloride C sodium fluoride D sodium iodide

5 If the price of one tonne (1000 kg) of sulphur is £160, what is the cost (to the nearest pound) of the sulphur needed to make one tonne of sulphuric acid (H2SO4)?

A £52 B £98 C £160 D £490

6 A positively charged particle with the electron configuration 2.8 is

A an aluminium ion. B a fluoride ion. C an oxide ion. D a potassium ion.

7 The orbitals of a nitrogen atom may be represented as shown.

Which one of the following diagrams represents the arrangement of electrons in the ground state of the nitrogen atom?

A

B

C

D

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1s 2s 2p








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8 Which one of the following chloro-compounds is non-polar?

A HCl B CCl4 C CH3Cl D CHCl3

9 Which one of the following contains a coordinate bond?

A N2 B NH3 C NH2

–

D NH+ 4

10 The enthalpy of neutralisation when an acid reacts with an alkali is, by definition, the number of kilojoules released by

A the formation of one mole of water. B the formation of one mole of salt. C the neutralisation of one mole of acid. D the neutralisation of one mole of alkali.








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Section B

Answer all six questions in the spaces provided.

11 There are three accepted intermolecular forces i.e.

● van der Waals forces (attractions between induced dipoles) ● permanent dipole attractions ● hydrogen bonding

Complete the following table.

✓ = present ✘ = not present

[3]

liquid van der Waals permanent dipole hydrogen bonding

water ✓ ✘ ✓

ammonia

xenon

hydrogen chloride








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12 Rocket fuels need to supply a large amount of energy, yet have a low mass i.e. a high power to weight ratio.

One potential reaction is that of fluorine with diborane (a boron hydride).

H H � � H— B—B —H � � H H

diborane

B2H6(g) + 6F2(g) → 6HF(g) + 2BF3(g)

(a) Calculate the enthalpy change when one mole of diborane reacts completely with fluorine given the following bond enthalpies.

Bond kJ mol–1

F — F 158 B — H 389 B — B 293 H — F 566 B — F 627

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]

(b) The bond enthalpy of hydrogen fluoride is 566 kJ mol–1 whereas that of hydrogen iodide is 299 kJ mol–1. State what would be observed when hydrogen fluoride and hydrogen iodide are heated.

__________________________________________________________

__________________________________________________________

_______________________________________________________ [2]








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(c) Calculate the total volume of gas produced at 20 °C and one atmosphere pressure by the complete reaction of 7 g of diborane with fluorine.

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]

(d) State and explain the shape of the boron trifluoride molecule.

__________________________________________________________

__________________________________________________________

_______________________________________________________ [2]








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13 Calcium fluoride, CaF2, occurs naturally as fluorite or fluorspar. Impurities give a blue variety known as Blue John. Fluorspar is the major source of hydrogen fluoride and fluorine. It can be prepared in the laboratory by precipitation or direct combination of the elements.

(a) What is the colour of pure calcium fluoride?

_______________________________________________________ [1]

(b) Explain the formation of calcium fluoride from calcium and fluorine atoms using dot and cross diagrams showing outer electrons only.

[3]

(c) Write an equation for the precipitation of calcium fluoride by mixing solutions of calcium chloride and sodium fluoride.

_______________________________________________________ [1]

(d) Calcium fluoride reacts with concentrated sulphuric acid to form hydrogen fluoride and calcium sulphate. Write an equation for the reaction.

_______________________________________________________ [2]








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(e) Calcium fluoride is sparingly soluble in water; 0.0025 g dissolves in 100 cm3 of water at 18 °C.

Calculate the concentration of fluoride ions in moles per litre using the following headings:

(i) relative formula mass of calcium fluoride

____________________________________________________ [1]

(ii) number of moles of calcium fluoride in 0.0025 g

____________________________________________________ [1]

(iii) number of moles of fluoride ion in 0.0025 g of calcium fluoride

____________________________________________________ [1]

(iv) number of moles of fluoride ion in 100 cm3 of water

____________________________________________________ [1]

(v) number of moles of fluoride ion per litre of water

____________________________________________________ [1]

(f) The natural presence of fluoride ions in domestic water supplies is regarded as beneficial by some, but the deliberate addition of fluoride ions is controversial.

(i) State a benefit of fluoride ions in drinking water.

____________________________________________________ [1]

(ii) Explain why some people object to the addition of fluoride ions to drinking water.

______________________________________________________

____________________________________________________ [1]








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14 All naturally occurring sodium atoms are represented by the symbol 2311Na.

However, radioactive isotopes of sodium may be prepared.

(a) (i) State the number of electrons, protons and neutrons in an atom of 2311Na.

______________________________________________________

____________________________________________________ [2]

(ii) Explain why 2311Na and 24

11Na are regarded as isotopes.

______________________________________________________

____________________________________________________ [2]

(b) A sample of sodium from a nuclear reactor contains 2.00% of 2411Na

and 98.00% of 2311Na by mass. Calculate the relative atomic mass of the

sodium in the sample to two decimal places.

__________________________________________________________

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]








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(c) A major use of sodium metal is in sodium street lamps. The lamp contains mercury vapour which conducts electricity at high voltages. Sodium within the lamp vaporises and the electrical energy causes yellow (orange) light to be given out.

(i) Explain, in terms of energy levels, how the yellow light is generated.

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(ii) When the light from the sodium lamp is analysed the spectrum shows two bright yellow lines at wavelengths of 589 nm and 589.6 nm.

(1 nm = 1 × 10–9 m)

λ →

Using the equations E = hv and c = vλ, calculate the energy change (in joules) associated with the line at 589 nm.

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [3]

589 589.6








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(d) If larger amounts of energy are supplied to sodium vapour (gas) it ionises.

(i) Using state symbols, write the equation which represents the first ionisation energy of sodium.

____________________________________________________ [1]

(ii) The value of the first ionisation energy for sodium is 500 kJ mol–1. The second ionisation energy has a value of 4513 kJ mol–1. Explain why this is a much higher value.

______________________________________________________

______________________________________________________

____________________________________________________ [2]

(iii) The first ionisation energies of the alkali metals are shown in the graph below.

Explain why lithium has a higher first ionisation energy than sodium.

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [2]

Quality of written communication [2]

500

400

300Cs

Rb

K

Na

Li

firstionisation

energykJ mol–1








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15 Oxalic acid (ethanedioic acid) is a weak dicarboxylic acid. It is similar to acetic acid (ethanoic acid) which is a monocarboxylic acid.

COOH CH3COOH � COOH

oxalic acid acetic acid

(a) Suggest the meaning of the term dicarboxylic acid.

__________________________________________________________

_______________________________________________________ [2]

(b) Both of these acids react with alkalis. Write the equation for the reaction of oxalic acid with excess sodium hydroxide.

_______________________________________________________ [2]

(c) Weak acids such as oxalic acid can be titrated with strong alkalis using phenolphthalein as indicator.

State the colour of phenolphthalein in alkaline and acidic solution.

alkali __________________________________________________ [1]

acid ___________________________________________________ [1]

(d) When oxalic acid reacts with a small amount of phosphorus pentachloride, a mixture of gases is produced:

COOH PCl5 + � → CO2 + CO + 2HCl + POCl3 COOH

How could you show, using a named silver salt, that hydrogen chloride was a product?

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]








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(e) Oxalic acid reacts with excess phosphorus pentachloride to give oxalyl chloride which has the following percentage composition by mass.

(i) Calculate the empirical formula of oxalyl chloride.

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(ii) Explain why the empirical formula could not be the molecular formula.

______________________________________________________

____________________________________________________ [1]

element % composition

carbon 18.9

chlorine 55.9

oxygen 25.2








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16 In the laboratory, ammonia can be prepared by heating a mixture of ammonium chloride and calcium hydroxide, as shown in the diagram below.

The equation for the reaction is:

2NH4Cl + Ca(OH)2 → 2NH3 + CaCl2 + 2H2O

(a) The ammonia gas is collected upwards. Calculate the relative molecular masses of ammonia, oxygen and nitrogen and use them to explain why ammonia is collected in this way.

ammonia __________________________

oxygen __________________________

nitrogen __________________________

explanation ________________________________________________

_______________________________________________________ [3]

(b) Calculate the volume of ammonia produced, at 20 °C and one atmosphere pressure, if 1.07 g of ammonium chloride are heated with excess calcium hydroxide.

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]

ceramic wool

calciumoxideheat

ammonium chloride+ calcium hydroxide

test tube








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(c) Ammonia gas is alkaline.

(i) Describe the effect of ammonia on moist Universal Indicator paper.

____________________________________________________ [1]

(ii) Gases are normally dried using concentrated sulphuric acid. Explain why this drying agent would be unsuitable in the case of ammonia.

______________________________________________________

____________________________________________________ [2]

(iii) Ammonia may be detected using concentrated hydrochloric acid. Write the equation for the reaction and describe what is seen.

______________________________________________________

____________________________________________________ [2]

(d) Ammonia can act as a reducing agent. When passed over heated copper(II) oxide the following reaction occurs:

2NH3 + 3CuO → 3Cu + N2 + 3H2O

Deduce the oxidation numbers of nitrogen and copper in the reactants and products and use them to explain the redox change.

__________________________________________________________

__________________________________________________________

_______________________________________________________ [3]








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(e) The shape of the ammonia molecule is explained using electron pair repulsion theory.

(i) Draw the dot and cross diagram for ammonia showing only the outer electrons.

[2]

(ii) Draw and name the shape of the ammonia molecule.

name __________________________ [2]

(iii) Explain why the bond angle in ammonia is not 109°.

______________________________________________________

____________________________________________________ [2]








ADVANCED SUBSIDIARY (AS)General Certifi cate of Education

2006

Chemistry



[ASC11]


MARK SCHEME








Section A

1 A

2 C

3 C

4 A

5 A

6 A

7 B

8 B

9 D

10 A

[2] for each correct answer [20] 20








P

Section B

11

[3] 3

12 (a) Bonds broken B–B 1 × 293 = 293

6B–H 6 × 389 = 2334 6F–F 6 × 158 = 948 Total 3575 Bonds formed 6H–F 6 × 566 = 3396

6B–F 6 × 627 = 3762 Total 7158 Energy given out = –3583 kJ [3]

(b) HF no effect [1] HI violet (vapour) [1] [2]

(c) B2H6 = 2 ×11 + 6 × 1 = 22 + 6 = 28 7 g of borane = 0.25 mol [1] 1 mol of borane gives 8 mol of gas 0.25 mol of borane gives 2 mol of gas = 2 × 24 dm3 = 48 dm3 [3] (d) Trigonal planar [1] Bond (electron) pairs repel (as far as possible) [1] [2] 10

13 (a) White [1]

(b)

[3]

(c) Either the ionic or full equation is acceptable

CaCl2 + 2NaF → CaF2 + 2NaCl Ca2+ + 2F– → CaF2 [1]

(d) CaF2 + H2SO4 → CaSO4 + 2HF [2]

ammoniaxenonhydrogen chloride

F F–

F F–

Ca2+Ca →

8/9 correct [3]7 correct [2]6 correct [1]≤5 correct [0]








(e) (i) CaF2 = 40 + 2 × 19 = 78 [1]

(ii) 0.0025/78 = 0.000032 = 3.2 × 10–5 [1] (iii) 6.4 × 10-5 [1] (iv) 6.4 × 10-5 [1] (v) 6.4 × 10-4 [1] (f) (i) Prevent tooth decay [1]

(ii) Freedom of choice/against mass medication [1] 14

14 (a) (i) Electrons = 11 Protons = 11 Neutrons = 12 [–1] for each error [2] (ii) Same number of protons [1] different number of neutrons/different mass [1] [2]

(b) Sodium-24 2 × 24 = 48 Sodium-23 98 × 23 = 2254 = 2302 relative atomic mass = 23.02 [3]

(c) (i) Electrons raised to a higher (energy) level (by electrical energy) [1] accept any value for C as not given in the question paper Electrons fall down to lower/original (energy) level [1] Light given out/energy given out as light [1] [3]

(ii) E= hν = hc/λ = 6.63 × 10–34 × 3 × 108

589 × 10–9 = 3.38 × 10–19 J [3] (d) (i) Na(g) → Na+(g) + e [1]

(ii) 2nd e less shielded [1] 2nd e closer to nucleus [1] 2nd e from full shell/pair [1] essential + any 1 other 2nd e removed from +ve ion [1] [2]

(iii) Electron closer to lithium/electron further away in sodium [1] Electron more shielded in sodium/electron less shielded in lithium [1] [2]

Quality of written communication [2] 20








15 (a) Two [1] COOH groups [1] [2]

(b) COOH COONa + 2NaOH → + 2H2O [2] COOH COONa (c) Pink [1] Colourless [1] [2]

(d) Dissolve HCl in water [1] silver nitrate [1] white precipitate [1] [3]

(e) (i) carbon: 18.9 18.9/12 = 1.575 chlorine: 55.9 55.9/35.5 = 1.575 oxygen: 25.2 25.2/16 = 1.575 COCl [3]

(ii) O == C — Cl carbon should have a valency of four [1] 13

16 (a) Ammonia = 17 Oxygen = 32 Nitrogen = 28 [2] Ammonia is lighter than air (which is N2 + O2 ) [1] [3] (b) NH4Cl = 14 + 4 + 35.5 = 53.5 1.07g = 1.07/53.5 = 0.02 mol 0.02 × 24 = 0.48 dm3 [3]

(c) (i) Turns blue/violet [1] (ii) Ammonia alkaline [1] reacts with acid [1] forms ammonium sulphate [1] any 2 from 3 [2]

(iii) NH3 + HCl → NH4Cl [1] white smoke [1] [2] (d) NH3 –3 N2 0 [1] CuO +2 Cu 0 [1] nitrogen oxidised and copper reduced [1] [3]








(e) (i)

[2]

(ii) N

H H H [1]

Pyramidal [1] [2]

(iii) Lone pair [1] repels bond pairs [1] [2] 20

Section B 80

Total 100

N

H

HH








TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in question 13(d).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.


January 2007

Chemistry



[ASC11]

THURSDAY 18 JANUARY, MORNING








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Section A

For each of the questions only one of the lettered responses (A – D) is correct.


1 Which one of the following types of bonding can be described as intermolecular?

A covalent bonding B hydrogen bonding C ionic bonding D metallic bonding

2 Which one of the following gives the correct flame colours for each of the metal chlorides?

3 Which one of the following gases occupies the largest volume at 20 °C and one atmosphere pressure?

A 0.32 g of oxygen B 0.44 g of carbon dioxide C 0.02 g of hydrogen D 0.80 g of argon

4 Which one of the following contains hydrogen bonds?

A ice (s) B iodine (s) C methane (g) D quartz (s)

Metal chloride A B C D

CaCl2 crimson red red green

LiCl green green crimson red

BaCl2 red crimson green crimson








5 The deuterium ion, 21D+, contains

A one proton only. B one proton and one neutron only. C one proton and one electron only. D one proton and one neutron and one electron only.

6 Which one of the following molecules has a shape which is influenced by a lone pair of electrons?

A BeCl2 B CO2 C C2H4 D NH3

7 Which one of the following equations represents the third ionisation energy of aluminium?

A Al(g) → Al3+(g) + 3e–

B Al+(g) → Al2+(g) + e–

C Al2+(g) → Al3+(g) + e–

D Al3+(g) → Al4+(g) + e–

8 The enthalpy change for the formation of ammonia is –46.2 kJ mol–1. What is the enthalpy change, in kJ, for the reaction 2NH3(g) → N2(g) + 3H2(g)?

A –46.2 B +46.2 C –92.4 D +92.4








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9 The Pauling electronegativity values for the following elements are:

Which one of the following bonds will be the most polar?

A C—F B Cl—H C C—H D O—F

10 In which one of the following atoms is the number of protons greater than the number of neutrons?

A 2H B 3He C 10B D 238U

hydrogen carbon chlorine oxygen fluorine

2.1 2.5 3.0 3.5 4.0








ASC1W7 2739 5 [Turn over

Section B


11 Circle the three polar molecules in the table below.

[3]

N2 CO2 BF3 C2H4

H2O NH3 HCl CH4








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12 Naturally occurring bromine consists of two isotopes. The following mass spectrum shows the molecular ion region.

(a) (i) Define the term isotope.

______________________________________________________

____________________________________________________ [2]

(ii) State the mass numbers of the two bromine isotopes.

____________________________________________________ [1]

(iii) Deduce the percentage abundance of each isotope.

______________________________________________________

____________________________________________________ [1]

(b) The isotopes cannot be chemically distinguished. They both react with sodium and phosphorus in the same way to form sodium bromide and phosphorus tribromide respectively.

(i) Explain why the isotopes react in the same way.

______________________________________________________

____________________________________________________ [1]

100

50

158 160 162

relativeabundance

mass/charge








(ii) Write the equation for the reaction of sodium with bromine.

____________________________________________________ [2]

(iii) Write the equation for the reaction of phosphorus, P4, with bromine.

____________________________________________________ [2]

(c) Both sodium bromide and phosphorus tribromide can be used to prepare hydrogen bromide. Hydrolysis of phosphorus tribromide yields pure hydrogen bromide gas. The reaction of sodium bromide with concentrated sulphuric acid yields an impure gas.

PBr3 + 3H2O → 3HBr + H3PO3

(i) Calculate the volume of hydrogen bromide produced, at 20 °C and one atmosphere pressure, when 9.0 g of phosphorus tribromide are completely hydrolysed.

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(ii) Write the equation for the reaction of sodium bromide with concentrated sulphuric acid to form hydrogen bromide.

____________________________________________________ [2]

(iii) Explain why the hydrogen bromide, prepared using sulphuric acid, is impure and name the impurities present.

______________________________________________________

____________________________________________________ [2]

(d) Hydrogen bromide decomposes at high temperatures. Name the products.

_______________________________________________________ [1]








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13 The neutralisation of hydrochloric acid with potassium hydroxide can be used to illustrate Hess’s Law as shown by the diagram below.

(a) (i) Write the equation for the reaction of KOH(s) with HCl(aq).

____________________________________________________ [1]

(ii) Define the term standard enthalpy change of neutralisation.

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(b) (i) State Hess’s Law in words.

______________________________________________________

______________________________________________________

____________________________________________________ [2]

(ii) Write an equation to express Hess’s Law using ΔH1, ΔH2 and ΔH3 from the above enthalpy cycle.

____________________________________________________ [1]

ΔH1

KOH(s) KCl(aq)

ΔH3ΔH2 +H2O(l)

+HCl(aq)

KOH(aq)








(c) What solution must be added to KOH(aq) in order to measure ΔH3?

_______________________________________________________ [1]

(d) Explain how you would determine the value of ΔH1 experimentally using 5.6 g of potassium hydroxide.

__________________________________________________________

__________________________________________________________

__________________________________________________________

__________________________________________________________

__________________________________________________________

_______________________________________________________ [4]









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14 Silver metal is valued for its appearance and its conductivity. Its electrical conductivity is greater than that of any other metal.

Before 1919 British silver coinage contained 92.5% silver. At the end of 1946 silver coinage was replaced by cupro-nickel alloy coins.

Silver metal will dissolve in all concentrations of nitric acid to form silver nitrate.

(a) Explain the process by which silver metal conducts electricity.

__________________________________________________________

_______________________________________________________ [2]

(b) The equation for the reaction of silver with nitric acid is:

3Ag + 4HNO3 → 3AgNO3 + 2H2O + NO

(i) Determine the oxidation numbers of the underlined elements in the following species from the equation and use them to explain the redox reaction taking place.

Ag HNO3 AgNO3 NO

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(ii) Silver nitrate is amazingly soluble in hot water; 100 cm3 of boiling water dissolves 952 g of silver nitrate.

Calculate the molarity of this saturated solution, assuming no change in volume from 100 cm3.

______________________________________________________

______________________________________________________

____________________________________________________ [3]








(c) A coin containing silver can be dissolved in nitric acid and silver chloride, AgCl, precipitated using sodium chloride solution. After drying and weighing the mass of the precipitate is used to determine the amount of silver present in the coin.

(i) Write an equation for the reaction of silver ions with chloride ions.

____________________________________________________ [1]

(ii) Calculate the percentage of silver in a coin weighing 5.00 g which produces 6.11 g of silver chloride using the following headings:

relative formula mass of silver chloride

______________________________________________________

percentage of silver in silver chloride

______________________________________________________

mass of silver in 6.11g of silver chloride

______________________________________________________

percentage of silver in the coin

____________________________________________________ [4]

(d) The silver chloride precipitate is white when first obtained but changes its appearance if left for some time in the laboratory.

(i) State the change in appearance and explain why it takes place.

______________________________________________________

____________________________________________________ [2]

(ii) Suggest why this could affect the result for the percentage determination of silver.

____________________________________________________ [1]








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(e) Silver nitrate reacts with iodine as shown in the following equation. The reaction involves disproportionation.

6AgNO3(aq) + 3I2(aq) + 3H2O(l) → AgIO3(aq) + 5AgI(s) + 6HNO3(aq)

(i) Explain the term disproportionation.

______________________________________________________

____________________________________________________ [1]

(ii) Which element is undergoing disproportionation?

____________________________________________________ [1]

(iii) Describe what would be observed.

______________________________________________________

____________________________________________________ [2]








15 Ammonia does not burn in air but it will burn in oxygen.

(a) The equation for the combustion of ammonia in oxygen is:

4NH3(g) + 3O2(g) → 2N2(g) + 6H2O(g)

(i) If 20 cm3 of ammonia are reacted with the required amount of oxygen, what is the total volume of gases produced at the same temperature and pressure?

____________________________________________________ [1]

(ii) Using the table of bond enthalpies below, calculate the enthalpy change for the combustion of one mole of ammonia.

bond kJ mol–1

N—H 391 O——O 498 N———N 945 O—H 464

______________________________________________________

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(iii) Explain why this reaction is exothermic.

______________________________________________________

____________________________________________________ [1]








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(b) The partial oxidation of ammonia, using a platinum catalyst, is used in the manufacture of nitric acid.

The equation for the catalysed oxidation of ammonia is:

4NH3 + 5O2 → 4NO + 6H2O

This is also an exothermic reaction. Draw a labelled enthalpy level diagram for the reaction.

[3]

(c) Ammonia also reacts with chlorine in two different ways. Excess ammonia results in the formation of nitrogen and ammonium chloride.

The following equation is not balanced.

NH3 + Cl2 → N2 + NH4Cl

(i) Write the balanced equation and insert state symbols.

____________________________________________________ [2]

(ii) Write the formulae for the ions present in ammonium chloride.

____________________________________________________ [2]








(d) Ammonia reacts with a limited amount of chlorine to form nitrogen trichloride, an extremely explosive substance.

(i) Draw the dot and cross diagram for nitrogen trichloride showing outer electrons only.

[2]

(ii) Draw and explain the shape of a nitrogen trichloride molecule.

______________________________________________________

______________________________________________________

____________________________________________________ [3]








[Turn over

16 A solution of iodine in hexane can be titrated with standard sodium thiosulphate solution to determine the concentration of iodine.

(a) State the colour of iodine in hexane.

_______________________________________________________ [1]

(b) (i) State the indicator used, when this indicator is added to the conical flask and the colour change observed at the endpoint.

______________________________________________________

______________________________________________________

____________________________________________________ [3]

(ii) Suggest why it is important to swirl the flask vigorously after the addition of sodium thiosulphate solution.

____________________________________________________ [1]

sodium thiosulphatesolution

solution of








(c) The sodium compound produced during the titration has the following percentage composition.

(i) Complete the table.

[2]

(ii) Determine the empirical formula.

____________________________________________________ [1]

(iii) Write the formula of the sodium compound.

____________________________________________________ [1]

Grams of element in 100 g Moles of Ratio of moles Element of compound element of element

sodium 17.03

sulphur 47.41

oxygen 35.60

[Turn over[Turn over








advanced subsidiary (as)

General certificate of education

January 2007

MarkscheMe

chemistry

Assessment Unit AS 1

assessing


[asc11]

Thursday 18 January, MorninG








Section A

1 B

2 C

� D

4 A

5 B

6 D

7 C

8 D

9 A

10 B


Section A 20








Section B

11 H2O [1]NH3 [1]HCl [1] [3] 3

12 (a) (i) Same number of protons [1] different number of neutrons [1] [2]

(ii) 79 + 81 [1]

(iii) 50% [1]

(b) (i) Same electron structure [1]

(ii) 2Na + Br2 → 2NaBr [2]

(iii) P4 + 6Br2 → 4PBr3 [2]

(c) (i) PBr3 = 31 + 3 × 80 = 271 9.0 g = 9/271 = 0.033 mol

volume of hydrogen bromide = 3 × 24 × 0.033 dm3

= 2.376 dm3 [3]

(ii) NaBr + H2SO4 → NaHSO4 + HBr [2]

(iii) Br-/HBr reduces H2SO4 [1] bromine and sulphur dioxide and (water) [1] [2]

(d) Hydrogen and bromine [1] 17

13 (a) (i) KOH(s) + HCl(aq) → KCl(aq) + H2O(l) [1]

(ii) Heat/enthalpy produced [1] measured under standard conditions/or stated [1]

when one mole of water formed [1] in the reaction of an acid with an alkali [3]

(b) (i) If a given change can be brought about in more than one way the overall enthalpy change is the same for each way [2]

(ii) ∆H2 + ∆H3 = ∆H1 [1]

(c) HCl(aq) [1]

(d) Calorimeter/polystyrene cup [1] known volume of known M hydrochloric acid [1] measure temperature at start and end [1] stir [1] [4] Quality of written communication [2] 14








14 (a) Delocalised electrons [1]move/carry charge [1] [2]

(b) (i) Ag = 0; Ag in AgNO3

= +1 N in HNO3

= +5 N in NO = +2 [2] 0 → +1 is oxidation +5 → +2 is reduction [1] [3]

(ii) AgNO3

= 108 + 14 + 48 = 170 952 g = 952/170 = 5.6 mol

molarity = 56 M [3]

(c) (i) Ag+ + Cl– → AgCl [1]

(ii) AgCl = 108 + 35.5 = 143.5 [1]

% of silver in silver chloride =108 × 100

143.5 = 75.3% [1]

mass of silver in 6.11g of AgCl =75.3 × 6.11

100 = 4.6 g [1]

% of silver in the coin =4.6 × 100

5.0 = 92% [1] [4]

(d) (i) Goes grey [1] effect of light [1] [2]

(ii) mass of silver chloride less [1]

(e) (i) A redox reaction in which both reduction and oxidation of the same element occurs [1]

(ii) Iodine [1]

(iii) Brown solution [1] goes colourless [1] yellow precipitate [1] to max of [2] [2] 20








6ASC1W7 2739.01

15 (a) (i) 40 cm3 [1]

(ii) Bonds broken 12 N H = 12 × 391 = 4692 3 O O = 3 × 498 = 1494

= +6186

Bonds formed 2 N N = 2 × 945 = 1890 12 O H = 12 × 464 = 5568

= –7458

H = –7458 + 6186 = –1272 Bond energy = –1272/4 = –318 kJ [3]

(iii) Energy needed to break bonds less than that needed to form them [1]

(b)

enthalpy

reaction coordinate [3]

(c) (i) 8NH3(g) + 3Cl2(g) 6NH4Cl(s) + N2(g) [2]

(ii) NH4+

Cl– [2]

(d) (i)

[2]

(ii) N

Cl Cl Cl [1]

Lone pairs repel bond pairs [2] [3] 17

4NO + 6H2O

4NH3 + 5O2

Cl Cl

Cl








16 (a) Violet/pink/purple [1]

(b) (i) Add starch [1] when the colour of iodine is faint/straw colour [1] blue black to colourless [1] [3] (ii) To mix the hexane and water [1]

(c) (i)

[2]

(ii) NaS2O3 [1] (iii) Na2S4O6 [1] 9

Section B 80

Total 100

Element Moles of elementRatio of moles

of element

sodium 17.03/23 = 0.74 1

sulphur 47.41/32 = 1.48 2

oxygen 35.60/16 = 2.23 3








advanced subsidiary (as)

General certificate of education

January 2007

MarkscheMe

chemistry


assessing

Module 2: Organic, Physical and Inorganic Chemistry

[asc21]

Tuesday 23 January, MorninG








Section A

1 D

2 D

3 B

4 D

5 A

6 A

7 C

8 C

9 B

10 C


Section A 20








Section B

11

[3]

12 (a) (i) petrol, kerosene, diesel [3]

(ii) boiling point/condensing point [1]

(b) (i) C18H36 [1]

(ii) unpaired electron [1]

(iii) C16H34 2C2H4 + C3H6 + C9H20

or C16H34 4C2H4 + 2C3H6 + C2H6 [2]

(iv) bromine water [1] goes colourless [1] [2] 10

13 (a) (i) glowing splint relights [2]

(ii) removes the NO2 which has a low boiling point/easily condensed [1]

(b) (i) the equilibrium is continually moving both ways [1]

(ii) moves to the right-hand side [1] reaction is endothermic/absorbs the heat [1] [2]

(iii) moves to the left-hand side [1] two molecules on the RHS and one molecule on LHS [1] [2]

(iv) more yellow than brown/lightens [1]

(c) passed over a catalytic converter [1] nitrogen formed [1] carbon dioxide formed [1] heterogeneous/chemisorption/adsorption [1] (3 from 4) [3]

(d) (i) 720 dm3 of nitrogen dioxide = 720/24 = 30 mol [1]

(ii) 30 mol [1]

(iii) 30/2 = 15 dm3 [1] 15

low high high

high low high[1] per column.








14 (a) (i) lactic acid contains the structure CH3CHOH [1]

(ii) (heated with) iodine/NaClO + KI [1] and sodium hydroxide [1] yellow precipitate of iodoform [1] [3]

(b) CH3COOCH2COOH [1]

CH2BrCOOH [1]

CHOCOOH [1]

COOHCOOH [1] [4]

(c) CH3CHOHCOOH = C3H6O3 = 90 12 g = 12/90 = 0.133 mol

molarity = 1.33 M [3] 11

15 (a) (i) measured volume of water and temperature measured with thermometer/water bath at 25° C [1] add Sr(OH)

2until no more dissolves [1]

determine mass added, e.g. titration or gravimetric [1] [3]


(ii) hydration energy and lattice energy [1] lattice enthalpy decreases more rapidly [1] [2]

(iii) 1.72 × 10–3 mol of magnesium hydroxide in 100 cm3 1.72 × 10–2 mol of magnesium hydroxide in 1000 cm3 3.44 × 10–2 mol of hydroxide ion in 1 dm3

0.034 M [2] (iv) Ca(OH)2 = 40 + 34 = 74 [1] not g (–1) 0.15/74 = 0.00203 mol calcium hydroxide [1] 0.00203 mol carbon dioxide [1] 24 × 103 × 0.00203 = 48.72 cm3 [1] [4]

(b) (i) Ba(OH)2 + 2HNO3 Ba(NO3)2 + 2H2O [2]

(ii) green [1]

(iii) barium sulphate is insoluble [1]

(c) (i) stability increases down the group [1]

(ii) cation size [1] lattice enthalpy [1] [2]








(d) (i) BaCO3 BaO + CO2 [1]

(ii) limewater [1] goes milky [1] [2]

16 (a) (i) chlorine [1] (ii) nickel [1] (iii) iron [1]

(iv) vanadium pentoxide [1]

(v) sulphuric acid [1]

(b)

diagram [2] explanation [1] [3]

(c) (i) enzyme [1]

(ii) enzyme molecule has special shape [1] into which the reactant/substrate fits [1] [2]

(d) catalyst causes stereospecific polymerisation/leading to unbranched polymer [1] [1]

(e) catalyst speeds up reaction but is not used up/is reformed [1] light is “used up” [1] [2]

enthalpy

reactants

reaction coordinate

uncatalysed

catalysed

products








(f) (i) catalyst and reactants are in the same state [2]

(ii) e.g. esterification [1]

(iii) e.g. hardening of oils/Haber/catalytic converter [1] 18

Section B 80

Total 100









2007Chemistry


assessing


[ASC11]


TIME

1 hour 30 minutes.


Write your Centre Number and Candidate Number in the spaces

provided at the top of this page.

Answer all seventeen questions.

Answer all ten questions in Section A. Record your answers by

marking the appropriate letter on the answer sheet provided. Use only

the spaces numbered 1 to 10. Keep in sequence when answering.

Answer all seven questions in Section B. Write your answers in the

spaces provided in this question paper.


The total mark for this paper is 100.

Quality of written communication will be assessed in

Question 15(a)(ii).In Section A all questions carry equal marks, i.e. two marks for each

question.

In Section B the figures in brackets printed down the right–hand side

of pages indicate the marks awarded to each question or part question.

A Periodic Table of Elements (including some data) is provided.








Section A

For each of the questions only one of the lettered responses (A–D) is correct.


1 Which one of the following elements has a giant covalent structure?

A calcium

B carbon

C chlorine

D sulphur

2 The mass spectrum of an element X is shown below:

What is the relative atomic mass of X?

A 10.2

B 10.5

C 10.8

D 10.9

3 Which one of the following molecules has a bond angle of 120°?

A BF3

B BeCl2

C H2O

D NH3

100100

50

00 2 4 6 8 10 12

Mass/charge

Relativeabundance








4 A solution of a metal salt, sprayed into a Bunsen flame, gave a red colour.

The same solution gave a yellow precipitate with acidified silver nitrate solution. The metal

salt was

A calcium bromide.

B lithium bromide.

C sodium iodide.

D strontium iodide.

5 Given the following standard enthalpies of combustion:

C(s) –393 kJ mol–1

H2(g) –286 kJ mol–1

CH3COOH(l) –487 kJ mol–1

What is the standard enthalpy of formation of ethanoic acid, CH3COOH(l)

A –192

B 192

C –871

D 871

6 The standard redox potentials for chlorine, bromine and iodine are:

E /V Cl

2(g) + 2e– 2Cl–(aq) +1.36

Br2(g) + 2e– 2Br–(aq) +1.09

I2(g) + 2e– 2I–(aq) +0.54

Which one of the following statements is correct?

A When chlorine is added to a solution of iodine the brown colour disappears.

B When bromine is added to a solution of chloride ions the red colour disappears.

C When iodide ions are added to a solution of chlorine a brown colour appears.

D When bromide ions are added to a solution of iodide ions a red colour appears.








7 0.84 g of a Group II carbonate, MCO3, reacts exactly with 20.0 cm3 of a 1.0 mol dm–3

solution of hydrochloric acid.

MCO3 + 2HCl → MCl

2 + H

2O + CO

2

What is the metal, M?

A barium

B beryllium

C calcium

D magnesium

8 Which one of the following molecules contains a triple bond?

A BF3

B C2H

4

C N2

D NH3

9 Which one of the following represents the electronic configuration of calcium in the ground

state?

A 1s22p63d104s2

B 1s22s22p63s23p64s2

C 1s22s22p63s23p63d2

D 1s22s22p63s23p63d14s1

10 Which one of the following indicators would not be suitable for the titration listed?

Acid solution Base solution Indicator

A 0.1 M sulphuric acid 0.1 M sodium carbonate phenolphthalein

B 0.1 M hydrochloric acid 0.1 M sodium carbonate methyl orange

C 0.1 M hydrochloric acid 0.1 M sodium hydroxide methyl orange

D 0.1 M ethanoic acid 0.1 M sodium hydroxide phenolphthalein








Section B

Answer all seven questions in the spaces provided.

11 The possibility of life on Mars may be investigated by analysing the ratio

of carbon-12 to carbon-13 in the methane found in its atmosphere.

(a) (i) Complete the following table:

Number of protons Number of neutrons

Carbon-12

Carbon-13

[2]

(ii) Explain why carbon-12 and carbon-13 are isotopes.

______________________________________________________

___________________________________________________ [2]

(b) Draw and explain the shape of methane.

___________________________________________________________

_______________________________________________________ [3]








12 Chlorine dissolves in water and disproportionates as follows:

Cl2 + H

2O → HCl + HOCl

(a) Calculate the oxidation number of chlorine in the following molecules.

(i) Cl2 ________________

(ii) HCl ________________

(iii) HOCl ________________ [3]

(b) Explain the term disproportionation.

___________________________________________________________

_______________________________________________________ [2]

(c) (i) Write an ionic equation for the reaction which occurs when an

excess of a solution of chlorine is added to a solution of iron(II)

ions.

___________________________________________________ [2]

(ii) Explain whether the iron(II) is oxidised or reduced.

______________________________________________________

___________________________________________________ [2]








(d) Fluorine reacts with phosphorus, P4, to form phosphorus trifluoride,

PF3.

(i) Write an equation for this reaction.

___________________________________________________ [2]

(ii) Draw a dot and cross diagram to show the bonding in PF3; using

outer electrons only.

[2]

(iii) With reference to PF3 explain the octet rule.

______________________________________________________

___________________________________________________ [1]








13 People who live within 50 miles of a nuclear power plant are advised to

have a supply of “Anti-Rad” tablets readily available:

Each tablet weighs 0.085 g and contains potassium iodate.

In an experiment to determine the percentage of iodine in a tablet,

30 tablets were dissolved in water and added to excess acidified potassium

iodide solution, producing a solution of iodine.

(a) Write the equation for the reaction between iodide ions, I–, and iodate

ions, IO3

–, in the presence of acid.

_______________________________________________________ [2]

(b) The iodine solution was added to a 250 cm3 volumetric flask and made

up to the mark with deionised water. The number of moles of iodine

were found by titration of 25.0 cm3 portions against 0.02 mol dm–3

sodium thiosulphate solution.

(i) Write the equation for the reaction between iodine, I2, and sodium

thiosulphate, Na2S

2O

3.

___________________________________________________ [2]

(ii) Suggest a suitable indicator for the reaction and state when it

should be added to the titration flask.

______________________________________________________

___________________________________________________ [2]








(c) It was found that 22.5 cm3 of aqueous sodium thiosulphate were

required to react with the 25.0 cm3 portion of iodine solution.

(i) Calculate the number of moles of sodium thiosulphate used.

___________________________________________________ [1]

(ii) Calculate the number of moles of iodine, I2, in 25.0 cm3 of

solution.

___________________________________________________ [1]

(iii) Calculate the mass of iodine in 25.0 cm3 of solution.

___________________________________________________ [1]

(iv) Calculate the mass of iodine in 250 cm3 of solution.

___________________________________________________ [1]

(v) Calculate the mass of iodine in one tablet.

___________________________________________________ [1]

(vi) Calculate the percentage of iodine in each tablet.

___________________________________________________ [1]








14 Sodium was first isolated by Sir Humphrey Davy in 1807 by the

electrolysis of the molten hydroxide.

(a) Sodium hydroxide is an ionic solid.

(i) State the charge on a sodium ion.

___________________________________________________ [1]

(ii) Draw a diagram of a sodium ion showing all the electrons present.

[1]

(iii) State the charge on a hydroxide ion.

___________________________________________________ [1]

(iv) Draw a dot and cross diagram of a hydroxide ion showing the

outer electrons only.

[2]








(v) Explain why molten sodium hydroxide is able to conduct

electricity.

___________________________________________________ [1]

(vi) Write an equation for the conversion of sodium ions to sodium

metal.

___________________________________________________ [1]

(vii) Write an equation for the conversion of hydroxide ions to oxygen

and water.

___________________________________________________ [1]

(b) (i) Draw the structure of a typical metal showing the delocalised

electrons present.

[2]

(ii) State two properties of sodium, other than conductivity, which are

typical of a metal.

1. ____________________________________________________

2. __________________________________________________ [2]








15 (a) Potassium hydroxide may be converted to potassium chloride by

neutralisation with hydrochloric acid.

(i) Write the equation for this reaction.

___________________________________________________ [1]

(ii) Briefly describe how the enthalpy of neutralisation of this reaction

may be determined. Give experimental details, including one

source of error and one safety precaution. Details of calculations

involved are not required.

______________________________________________________

______________________________________________________

______________________________________________________

______________________________________________________

___________________________________________________ [4]


(iii) Explain the term enthalpy of neutralisation.

______________________________________________________

___________________________________________________ [2]

(b) (i) Using s,p,d notation, write the electron arrangement for the

potassium ion and the chloride ion.

Potassium ion __________________________ [1]

Chloride ion __________________________ [1]








(ii) Draw an s and a p orbital on the axes below:

[2]

(c) When a solution of potassium chloride is sprayed into a flame a lilac

colour is observed. Use a labelled energy level diagram to explain the

origin of this colour.

[3]

s orbital p orbital








16 The melting points of the elements in the third period are shown below:

(a) (i) Explain why the melting point of silicon is much higher than that

of aluminium.

______________________________________________________

___________________________________________________ [2]

(ii) Explain why the melting point of sulphur is much higher than that

of argon.

______________________________________________________

___________________________________________________ [2]

371K

922K 933K 803K

1683K

393K172K 84K

NaNa Mg Al Si P S Cl Ar








(b) The atomic radii of the elements in the third period are shown on the

graph below:

Explain the trend.

___________________________________________________________

_______________________________________________________ [2]

(c) The chart below shows the atomic and ionic radii of the elements in Group VII:

(i) Explain the size of the bromine atom relative to the bromide ion.

______________________________________________________

___________________________________________________ [1]

(ii) Explain the increase in atomic radius from fluorine to iodine.

______________________________________________________

___________________________________________________ [2]

0.20

0.18

0.16

0.14

0.12

0.10

0.08

Na Mg Al Si P S Cl Ar

Element

Atomic radius/nm

300

200

100

0

F

F–

Cl Br

Br–

I

Cl–I–

Radius/nm








(d) The electronegativities of the Group VII elements decrease down the

group.

(i) Explain the term electronegativity.

______________________________________________________

___________________________________________________ [2]

The electronegativities of hydrogen, boron and fluorine are 2.2, 2.0

and 4.0 respectively.

(ii) Using these values, explain why hydrogen fluoride is polar.

______________________________________________________

___________________________________________________ [1]

(iii) Explain why, even though the boron fluoride bond is polar, the

boron trifluoride molecule is non-polar.

______________________________________________________

___________________________________________________ [2]








17 The graph below shows the first ionisation energies for the first twelve

elements:

(a) Explain why element atomic number 8 has a lower first ionisation

energy than element 7.

_______________________________________________________ [1]

(b) Explain why elements 3 and 11 have low values.

_______________________________________________________ [1]

(c) Explain why elements 2 and 10 have high values.

_______________________________________________________ [1]

(d) Write the equation, including state symbols, for the first ionisation

energy for element number 4 (beryllium).

_______________________________________________________ [2]

1 5 10Atomic number

Ionisationenergy








ADVANCED SUBSIDIARY (AS)General Certifi cate of Education

2007

MARKSCHEME

ASC1S7P

Chemistry



[ASC11]


ASC11S

1









Section A

1 B

2 C

3 A

4 D

5 C

6 C

7 D

8 C

9 B

10 A


Section A 20








Section B

11 (a) (i) [2]

(ii) Same number of protons [1] different number of neutrons [1] [2]

(b) Tetrahedral shape [1] maximum repulsion [1] from (bonded pairs) of electrons [1] [3] 7

12 (a) (i) 0 [1]

(ii) –1 [1]

(iii) +1 [1]

(b) Oxidation number of an element increases and decreases in the same reaction [2]

(c) (i) Cl2 + 2Fe2+ → 2Fe3+ + 2Cl– [2]

(ii) Oxidised [1] it has lost an electron/oxidation number increases [1] [2]

(d) (i) P4 + 6F2 → 4PF3 [2]

(ii)

[2] (iii) (In forming compounds atoms share, gain or lose electrons) P/F have an outer shell of eight electrons [1] [1] 14

13 (a) 5I– + IO3– + 6H+ → 3I2 + 3H2O [2]

(b) (i) 2Na2S2O3 + I2 → 2NaI + Na2S4O6 [2]

(ii) Starch [1] near the end point/when straw coloured [1] [2]

(c) (i) 22.5 × 0.02/1000 = 0.00045 or 4.5 × 10–4 [1]

(ii) 0.00045/2 = 0.000225 mole [1]

(iii) 0.000225 × 254 = 0.05715 g [1]

(iv) 10 × 0.05715 g = 0.5715 g [1]

6 66 7

× × × F × × • × ×

• • P • × × •

× × F × × × ×

× F × × × × ×








(v) 0.5715 g/30 = 0.019 g 0.019 ÷ 6 = 0.0031 [1]

(vi) 0.00310.085

14 (a) (i) 1+/Na+ [1]

(ii) (+)

[1]

(iii) 1–/OH– [1]

(iv) (–)

[2]

(v) Ions move/ carry charge/current [1] [1]

(vi) Na+ + e– → Na [1]

(vii) 4OH– → 2H2O + O2 + 4e– [1]

(b) (i)

[2]

(ii) Malleability [1] lustrous [1] ductility [1] [2] 12

15 (a) (i) KOH + HCl → KCl + H2O [1]

(ii) Measure initial and fi nal temperatures of solution/change in temperature [1] Add excess HCl(aq) to a known quantity of KOH [1] Error – lack of insulation [1] Safety – protect eyes [1] [4]


(iii) Enthalpy of neutralisation is the enthalpy change [1] when 1 mole of water is produced by reaction between acid and alkali [1] [2]

(b) (i) 1s2 2s2 2p63s23p6 [1] 1s2 2s2 2p6 3s2 3p6 [1] [2]

× ×× ×

× ×

××××

× ×

× ×

××× • H

+ + + + + ++ + + + + +

+ + + + + +– – –– –

–– – – – – ––

– – –––








(ii)

or

[2]

(c) The colours observed during the fl ame test are due to the excitement of the electrons caused by absorption of energy. The electrons “jump” from their ground state to a higher energy level. As they return to n = 2 they emit visible light . . .

[3] 16

16 (a) (i) Silicon giant covalent [1] aluminium metallic [1] [2] (ii) Sulphur S8 large molecular/strong van der Waals forces [1] argon, weak van der Waals forces [1] [2]

(b) Nuclear charge increases [1] electrons in same shell drawn closer to nucleus [1] [2]

(c) (i) Increased repulsion due to added electron in bromide ion [1]

(ii) Increases from F to I, electrons are going into shells further and further from nucleus therefore less nuclear attraction [2]

(d) (i) Electronegativity is a measure of the tendency of an atom to attract a bonding [1] pair of electrons [1] [2]

(ii) HF is polar as F has a higher/different electronegativity than H [1]

(iii) The (trigonal symmetrical planar) structure [1] of BF3 is such that the dipole (moments) cancel out [1] [2] 14

n = 3n = 2

n = 1








17 (a) Element 7, (N), has a stable half-fi lled shell [1] Element 8 has a pair of electrons in a p orbital [1]

(b) Elements 3 and 11 are alkali metals; with one electron in their outer shell (they lose it to gain a stable outer shell) [1]

(c) Elements 2 and 10 are (noble gases with) fi lled shells [1]

(d) Be(g) → Be+(g) + e [2] 5

Section B 80

Total 100








TIME

1 hour 30 minutes.




Answer all sixteen questions.




Answer all six questions in Section B. Write your answers in the




Quality of written communication will be assessed in question

13(a)(iii).In Section A all questions carry equal marks, i.e. two marks for each

question.

In Section B the figures in brackets printed down the right-hand side




January 2008

Chemistry


assessing


[ASC11]









Section A

For each of the following questions only one of the lettered responses (A – D) is correct.


1 Which one of the following solids consists of molecular covalent crystals?

A Diamond

B Graphite

C Quartz

D Rhombic sulphur

2 Which one of the following compounds does not exhibit hydrogen bonding?

A Ammonia, NH3

B Ethanol, C2H

5OH

C Methane, CH4

D Water, H2O

3 Which one of the following graphs represents the change in melting point across the third

period from sodium to argon?

atomic number

melting

point

B

atomic number

melting

point

A

atomic number

melting

point

D

atomic number

melting

point

C








4 The hydrogen spectrum is best described as a series of

A equally spaced lines.

B lines becoming closer together as the frequency increases.

C lines becoming closer together as the wavelength increases.

D lines becoming closer together and then further apart.

5 Magnesium has three isotopes as shown below.

A 24.00

B 24.08

C 24.32

D 24.33

6 In an experiment 3.425 g of lead oxide were reduced to form 3.105 g of lead.

What is the empirical formula of the lead oxide?

A PbO

B Pb3O

2

C Pb3O

4

D Pb4O

3

7 Which one of the following ions has the smallest radius?

A F–

B Mg2+

C Na+

D O2–

Relative atomic massPercentage abundance

(%)

24 78.6

25 10.1

26 11.3








8 The electronegativity values for carbon, fluorine, nitrogen and oxygen are given below.

Which one of these values is most likely to be that of nitrogen?

A 2.5

B 3.0

C 3.5

D 4.0

9 Which one of the following is responsible for the open structure of ice?

A Covalent bonds

B Dipole–dipole attractions

C Hydrogen bonds

D Van der Waals forces

10 Phenol is converted to trichlorophenol (TCP) according to the equation below.

C6H

5OH + 3Cl

2 → C

6H

2Cl

3OH + 3HCl

If 50.0 g of phenol produced 97.6 g of TCP, what is the percentage yield of the TCP?

A 47.6%

B 49.4%

C 51.2%

D 92.9%








Section B

Answer all six questions in this section.

11 Metal salts are used to produce colours in fireworks.

(a) Complete the table below to give the flame colour of each of the metal

chlorides listed.

[2]

(b) Using a diagram, explain how the flame colours of the metal chlorides

originate.

_________________________________________________________

_________________________________________________________

_________________________________________________________

______________________________________________________________ [3]

Metal chloride Flame colour

Barium chloride

Strontium chloride








12 Beryllium chloride, BeCl2, and ammonia, NH

3, are covalent compounds.

(a) State the octet rule and explain whether beryllium chloride obeys the

rule.

_________________________________________________________

_________________________________________________________

______________________________________________________________ [2]

(b) The ammonium ion, NH4

+, contains a coordinate bond.

(i) Explain what is meant by a coordinate bond.

________________________________________________________

_____________________________________________________ [1]

(ii) Draw a dot and cross diagram, using outer electrons only, to show

the bonding in an ammonium ion.

[2]








13 A dilute solution of hydrogen peroxide can be used as an antiseptic

mouthwash.

The concentration of the hydrogen peroxide solution can be measured by

reaction with excess acidified potassium iodide solution.

H2O

2 + 2H+ + 2I– → 2H

2O + I

2

The liberated iodine is then estimated by titration with standard sodium

thiosulphate solution.

(a) (i) Name the indicator used in this titration.

_____________________________________________________ [1]

(ii) Explain what is meant by the term standard solution.

________________________________________________________

_____________________________________________________ [1]

(iii) Describe, giving experimental detail, how you would carry out the

titration. Assume all the apparatus is clean and dry.

________________________________________________________

________________________________________________________

________________________________________________________

________________________________________________________

________________________________________________________

________________________________________________________

_____________________________________________________ [4]









(b) 25.0 cm3 of the mouthwash was diluted to 250 cm3. A 25.0 cm3

portion of the diluted mouthwash required 22.5 cm3 of 0.1 M sodium

thiosulphate for a complete reaction.

(i) Write an equation for the reaction between sodium thiosulphate,

Na2S

2O

3, and iodine I

2.

__________________________________________________ [2]

(ii) Calculate the concentration of the hydrogen peroxide in the

mouthwash in g dm–3 using the following headings:

Number of moles of sodium thiosulphate used.

________________________________________________________

Number of moles of iodine, I2, liberated in 25.0 cm3 of diluted

mouthwash solution.

________________________________________________________

Number of moles of H2O

2 present in 25.0 cm3 of diluted

mouthwash solution.

________________________________________________________

Number of moles of H2O

2 in 25.0 cm3 of undiluted mouthwash.

________________________________________________________

Mass of H2O

2 in 25.0 cm3 of undiluted mouthwash.

________________________________________________________

Concentration of H2O

2 in undiluted mouthwash in g dm–3.

_____________________________________________________ [4]








14 Kerosene is a mixture of hydrocarbons. When it is mixed with hydrogen

peroxide it can be used as a rocket fuel.

(a) The kerosene burns in oxygen released by the hydrogen peroxide.

(i) One component of kerosene is dodecane, C12

H26

. Write an

equation for the complete combustion of dodecane.

_____________________________________________________ [2]

(ii) Draw an enthalpy level diagram for the combustion of dodecane

using the axes below, labelling the y-axis.

[2]

(iii) Explain, in terms of enthalpy, why the combustion of dodecane is

exothermic.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]

Extent of reaction








(b) The enthalpy of combustion of hydrocarbons can be calculated using

Hess’s Law.

(i) Explain what is meant by the term standard enthalpy of combustion.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]

(ii) State Hess’s Law.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]

(iii) One of the constituents of petrol is octane, C8H18. Use the

information below to calculate the enthalpy of combustion of

octane.

C8H

18(l) + 12 1–

2 O

2(g) → 8CO

2(g) + 9H

2O(l)

C(s) + O2(g) → CO

2(g) ΔH = –393.5 kJ mol–1

8C(s) + 9H2(g) → C

8H

18(l) ΔH = –250.0 kJ mol–1

H2(g) + 1–

2 O

2(g) → H

2O(l) ΔH = –286.0 kJ mol–1

________________________________________________________

________________________________________________________

________________________________________________________

________________________________________________________

Enthalpy of combustion ___________________ kJ mol–1 [3]








(c) The enthalpy of combustion can also be calculated using bond enthalpy

values.

(i) Explain what is meant by the term bond enthalpy.

________________________________________________________

_____________________________________________________ [1]

(ii) Use the bond enthalpy values below to calculate the total bond

enthalpy for the products and the overall enthalpy change for the

combustion of 1 mole of octane given that the total of the bond

enthalpies of the reactants is 16 548 kJ.

Bond Bond enthalpy

C = O 750 kJ mol–1

H – O 463 kJ mol–1

________________________________________________________

________________________________________________________

Total bond enthalpy of products _______________

Enthalpy change for the

combustion of octane ________________ [2]

(iii) Explain the difference in the value of the enthalpy of combustion

from Hess’s Law with that obtained using bond enthalpies.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]








15 Sodium and fluorine are reactive elements which combine directly to form

the ionic salt, sodium fluoride.

(a) Sodium has the typical properties of a metal.

(i) Draw a diagram to show the bonding in sodium metal.

[2]

(ii) Explain why sodium

is malleable.

________________________________________________________

_____________________________________________________ [1]

conducts electricity.

________________________________________________________

_____________________________________________________ [2]

(b) (i) Write an equation, including state symbols, for the formation of

sodium fluoride from sodium and fluorine.

_____________________________________________________ [2]

(ii) Draw dot and cross diagrams, using outer electrons only, to show

the formation of sodium fluoride from sodium and fluorine atoms.

[3]








(iii) Suggest why fluoride is added to some toothpastes.

_____________________________________________________ [1]

(iv) Explain why some people object to the fluoridation of drinking

water.

________________________________________________________

_____________________________________________________ [1]

(c) The graph below represents the first ionisation energies of elements

from sodium to argon.

(i) Explain the general rise in the first ionisation energy across the

period.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]

(ii) Explain the decrease in the first ionisation energy from numbers

12 to 13 and from 15 to 16.

12 to 13 ________________________________________________

________________________________________________________

15 to 16 ________________________________________________

_____________________________________________________ [2]

15

1612

13

atomic number

first

ionisation

energy








(iii) Complete the graph of the successive ionisation energies of

sodium.

[2]

number of electron removed

1 2 3 4 5 6 7 8 9 10 11

log

ionisation

energy








(c) The sodium salts of these acids are colourless ionic solids.

(i) How could you distinguish between solid samples of sodium

chloride and sodium bromide using aqueous silver nitrate and

dilute ammonia solution?

________________________________________________________

________________________________________________________

________________________________________________________

________________________________________________________

_____________________________________________________ [4]

(ii) Write the equation for the reaction of sodium bromide with

concentrated sulphuric acid.

_____________________________________________________ [2]

(iii) Give two observations when concentrated sulphuric acid is added

to solid sodium iodide.

________________________________________________________

________________________________________________________

_____________________________________________________ [2]








(d) Chlorine reacts with cold dilute sodium hydroxide according to the

equation:

Cl2 + 2NaOH → NaCl + NaOCl + H

2O

(i) State the oxidation number of chlorine in each of the following

species.

Cl2 _________________

NaCl _________________

NaOCl _________________ [3]

(ii) Explain, using the oxidation numbers, why this is described as a

redox reaction.

________________________________________________________

_____________________________________________________ [2]

(iii) What name is given to this type of redox reaction?

_____________________________________________________ [1]









January 2008

MARKSCHEME

ChemistryAssessment Unit AS 1

assessing


[ASC11]









Section A

1 D

2 C

3 A

4 B

5 B

6 C

7 B

8 B

9 C

10 D


Section A 20








Section B

11 (a) Green [1] Red/Crimson [1] [2]

(b) Electron (excited) to higher energy level [1] Electron falls back to lower energy level [1] Diagram

(Energy) emitted (as a quantum of) light [1] [3] 5

12 (a) Only 4 electrons around the central beryllium atom [1] instead of 8 Atom gain/lose share electrons to obtain 8 electrons in the outer shell [1] [2]

(b) (i) Both electrons in a (covalent) bond come from the same atom [1]

(ii)

H

H

H

HN

X•

X•

••X

•

[–1] for each mistake [2] 5

13 (a) (i) Starch [1]

(ii) A solution the concentration of which is known (accurately) [1]

(iii) Any four: Pipette known volume of the iodine solution into a (conical) flask [1] Titrate with the thiosulphate [1] until the colour of the solution is pale yellow then add starch [1] Continue adding the thiosulphate until the solution goes from blue/black to colourless [1] [4]


(b) (i) 2Na2S2O3 + I2 → Na2S4O6 + 2NaI accept ionic equation ([–1] for each mistake) [2]

+








(ii) Carry any error from the equation through

(0.1 × 22.5)/1000 = 2.25 × 10–3

(2.25 × 10–3)/2 = 1.125 × 10–3

1.125 × 10–3 1.125 × 10–3 × 10 = 0.01125 0.01125 × 34 = 0.3825 (0.3825/25) × 1000 = 15.3 (g dm–3). (Award full marks directly for the correct answer. Carry any errors through, [–1] for each mistake) [4] 14

14 (a) (i) C12H26 + 1812O2 → 12CO2 + 13H2O

([–1] for each mistake) [2]

(ii) y-axis labelled enthalpy [1] reactants and products correctly shown [1] [2]

(iii) Enthalpy of the reactants is greater than the enthalpy of the products [1]. (Difference) given out as heat [1]. (Or correct explanation in terms of breaking and making bonds.) [2]

(b) (i) Enthalpy of combustion: Enthalpy change on the complete combustion in oxygen [1] of 1 mole [1] (under standard conditions) [2]

(ii) Hess’s Law: The overall enthalpy change for a reaction is independent of the number of steps taken [1] provided the conditions remain the same [1] [2]

(iii) 250.0 + ∆H = 8(–393.5) + 9(–286.0) ∆H = –5472 (kJ mol–1) ([–1] for each mistake) [3]

(c) (i) The energy required to break (1 mole of) bonds in the gaseous state to form the gaseous atoms [1]

(ii) (16 × 750) + (18 × 463) = 20 334 [1] 16 548 – 20 334 = –3786 [1] [2]

(iii) Bond enthalpy is an average term/Hess’ Law refers to a specific compound [1]/varies in compounds [2] 18








15 (a) (i)

+ + + +

+ + + +

+ + + +

e– e–e– e–e–

e–e–

e–

e–

e–

e–

[2]

(ii) Malleable: Layers can slide over one another (and the bonds reform without breaking) [1] [1]

Conducts electricity: Delocalised (valence) electrons [1] are free to move [1] [2]

(b) (i) 2Na(s) + F2(g) → 2NaF(s) [2]

(ii)

x x

xxx

x

x x

xx

+ –

Na F

x x

Na

x x

F

([–1] for each mistake) [3]

(iii) To help reduce tooth decay [1]

(iv) Lack of freedom of choice (or suitable alternative) [1]

(c) (i) Shielding remains constant/electrons add to same shell [1] Nuclear charge increases giving a greater attraction for the outer electrons [1] [2]

(ii) 12 to 13: The outer electron is being added to the p-subshell which is slightly further from the nucleus making it easier to remove the electron/stability of shell [1]

15 to 16: The outer electron is forming the first pair in the p-subshell and it is slightly easier to separate a pair of electrons/ stability of shell [1] [2]

e–








(iii)

([–1] for each mistake) [2] 18

16 (a) The ability of an atom in a covalent bond to attract the electrons [1]

(b) (i) H2 + F2 → 2HF [1]

(ii) [1]

(iii) Thermal stability decreases down the Group [1] as strength of the HHal bond decreases down the Group [1] [2]

(iv) HAt > HI > HBr > HCl > HF [1]

(c) (i) * Make a solution of the solid [1] Add silver nitrate White precipitate for chloride [1] Giving a cream ppt for bromide [1] White ppt soluble in dilute ammonia solution [1] Cream ppt insoluble in ammonia solution [1]

either

(maximum of [4]) [4]

(ii) NaBr + H2SO4 → NaHSO4 + HBr [2]

(iii) Steamy fumes/purple vapour/smell of rotten eggs/heat produced/ (yellow) solid formed (any two, [1] each) [2]

(d) (i) Cl2: 0 [1] NaCl: –1 [1] NaOCl: +1 [1] [3]

(ii) (ON of Cl changes from) 0 to –1 = Reduction [1] (ON of Cl changes from) 0 to +1 = Oxidation [1] [2]

(iii) Disproportionation [1] 20

Section B 80

Total 100

δ+

Hδ–

F









January 2008

MARKSCHEME


assessing

Module 2: Organic, Physical and Inorganic Chemistry

[ASC21]

TUESDAY 22 JANUARY, MORNING








Section A

1 B

2 B

3 C

4 A

5 B

6 A

7 B

8 A

9 D

10 C


Section A 20








Section B

11 (a) C ═ C

H── H─

CH3H3C─

C ═ CH

── H─

CH3

H3C

─

cis trans [2]

(b) Restricted rotation [1] about C═C [1] [2] 4

12 (a) l-iodobutane [1]

(b) C─I bond [1] is weaker/weakest [1] (than C─Br and C─Cl) Breaks more easily [1] [2]

(c) (i)

[1]

(ii)

H ─ C ─ C ─ C ═ C ─ H│H

│H

H│

H│

CH3│

H│

[1] 5

13 (a) (i) Rate of forward reaction = rate of reverse reaction Concentration of any reactant or product constant [2]

(ii) Equilibrium shifts to RHS [1] 4 moles (g) LHS → 2 moles (g) RHS [1] [2]

(iii) Negative [1] When temperature increases shifts in endothermic direction/reverse reaction is endothermic [1] [2]

(b) Increases the rate of the forward reaction and reverse reactions equally [1] 7

H ─ C ─ C ─ C ─ C ─ H│H

│H

H│

H│

CH3│

H│

│H

│OH








14 (a) (i)

Correct dipole [1] [3]

(ii) electrophilic [1] addition [1] [2]

(b) Moles of ethene = 120 ÷ 24 000 = 0.005 moles 1:1 ratio Moles of bromine (Br2) = 0.005 moles Mass of bromine (Br2) = 0.005 × 160 = 0.8 g Volume of pure liquid bromine Volume (cm3) = mass (g) ÷ density (0.8 ÷ 3.2) = 0.25 cm3 [4] 9

15 (a) (i) White flame/white solid [1]

(ii) 2Mg + O2 → 2MgO [2]

(b) Moles of HCl = 2 × 0.05 = 0.1 moles 2:1 ratio Moles of H2 = 0.05 moles Volume of H2 = 0.05 × 24 = 1.2 dm3 or 1200 cm3 [3]

(c) (i) White [1] precipitate [1] [2]

(ii) Mg2+(aq) + CO32–(aq) → MgCO3(s) [2]

(d) (i) Decreases down the group [1] Use of lattice enthalpy (X) [1] Use of hydration enthalpy (Y) [1] Variation in Y is greater than variation in X down the group [1] [4]

(ii) Barium chloride (or nitrate) [1] 15

C ═ CH

── H─

HH ─Hδ+

Brδ–│

→ H ─ C ─ C ─ H

H│

H│

│H Br–

+→ H ─ C ─ C ─ H

H│

H│

│H

│Br








16 (a) T [1]

H C C C OH

H│

CH3│

│H

│H

H│

│H

[1] P [1]

H C C C C H

H│

OH│

│H

│H

H│

│H

H│

│H

[1] S [1]

Butan-2-ol [1] [6] (b) (i) Molecules which have the same molecular formula but different structural formula [2]

(ii) An alcohol which has 3 carbons directly attached to the same carbon as the hydroxyl group [2]

(iii) Heat each alcohol (separately) [1] with acidified (potassium) dichromate (APD)/acidified (potassium) manganate VII [1] with butan-1-ol APD orange to green/purple to colourless [1] with 2-methylpropan-2-ol APD remains orange/remains purple [1] Maximum [3] [3]


(c) (i) methyl group and a hydrogen attached to the same carbon as the hydroxy group [1]

(ii)

or Butan-2-ol

[1]

H C C C C H

H│

OH│

│H

│H

H│

│H

H│

│H

(iii) Alkaline solution [1] of iodine [1] yellow ppt [1] [3] 20

17 (a) (i) carbon dioxide and nitrogen [2]

(ii) 2CO + 2NO → N2 + 2CO2 [1]

(iii) Platinum/Rhodium/Palladium [1]








(iv) Maximum surface area/efficiency [1] minimises cost [1] better catalysis [1] Any two from three [2]

(v) The catalyst is in a different physical state (solid) [1] than the reactants (gases) [1] [2]

(vi) Adsorbs to surface [1] Weakens bonds (in the reactants)/correct orientation/d-orbitals argument [1] [2]

(vii) The lead compounds bond very strongly (chemisorb) to the catalytic surface [1] poisoning the catalyst/preventing catalysis [1] [2]

(b) (i) y-axis = number of molecules [1]

(ii) The minimum amount of energy required for reaction [1]

(iii) Most collisions (between molecules of CO and NO) have energy less than activation energy [1]

(iv)

Numberof

molecules

Peak to RHS+

Lower

Broadening

Kinetic energy

T1

EA

Peak to RHS + Lower [2] Broadening [1] [3]

(v) A greater fraction of all collisions (between CO and NO) [1] have energy greater than or equal to the activation energy [1] [2] 20

Section B 80

Total 100









2008

Chemistry


assessing


[ASC11]


TIME

1 hour 30 minutes.




Answer all sixteen questions.




Answer all six questions in Section B. Write your answers in the




Quality of written communication will be assessed in

Question 14(b).In Section A all questions carry equal marks, i.e. two marks for each

question.

In Section B the fi gures in brackets printed down the right-hand side










Section A



1 An element X has 3 electrons in its outermost shell. Element Y has 6 electrons in its

outermost shell. What is the empirical formula of the compound formed when X and Y

combine?

A XY2

B X2Y

C X2Y

3

D X3Y

2

2 Which one of the following equations represents a redox reaction?

A 2NH3 → N

2 + 3H

2

B CuO + H2SO

4 → CuSO

4 + H

2O

C H2O + H+ → H

3O+

D AgNO3 + KI → KNO

3 + AgI

3 Chlorine exists as isotopes, 35Cl and 37Cl, resulting in a relative atomic mass of 35.5. Which

one of the following statements is correct?

A The isotopes have different chemical properties.

B The 37Cl isotope has a natural abundance of 75%.

C The mass spectrum of molecular chlorine, Cl2, includes peaks at m/e 70, 72 and 74.

D The nuclei of the two isotopes have the same number of neutrons.

4 Which one of the following represents the electronic structure for the V2+ ion in the ground

state?

A 1s22s22p63s23p63d3

B 1s22s22p63s23p63d24s1

C 1s22s22p63s23p63d14s2

D 1s22s22p63s23p64s24p1








5 A metal chloride solution reacts with aqueous sodium hydroxide according to the equation:

MClx + xNaOH → M(OH)

x + xNaCl

10.0 cm3 of 0.5 M MClx solution reacted with 7.5 cm3 of 2M sodium hydroxide. What is the

formula of the metal chloride?

A MCl

B MCl2

C MCl3

D MCl4

6 Which one of the following correctly shows the trends in atomic radii across and down the

Periodic Table?

7 Which one of the following gives the correct colour of the solution formed when the named

substance is dissolved in the named solvent?

substance solvent colour

A iron(II) chloride water red–brown

B potassium bromide water colourless

C iodine hexane brown

D chlorine hexane purple

A increase

increase

C increase

increase

B increase

increase

D increase

increase








8 Which one of the following statements is correct?

A 24 dm3 of hydrogen gas, H2, at 20 °C and 1 atmosphere contains 2.4 × 1024 atoms of

hydrogen.

B 500 cm3 of 1M copper(II) nitrate, Cu(NO3)2, solution contains 3.0 × 1023 nitrate ions.

C 256 g of rhombic sulphur contains 6.0 × 1023 S8 rings.

D 1 mole of sodium hydrogensulphate, NaHSO4, dissolves in water to form 1.8 × 1024 ions.

9 Given the following bond enthalpies:

bond bond enthalpy/kJ mol–1

C — C +347

C === C +612

C — O +358

C — H +413

O — H +464

What is the standard enthalpy change for the hydration of ethene?

C2H

4 + H

2O → C

2H

5OH

A –42 kJ mol–1

B +42 kJ mol–1

C –254 kJ mol–1

D +254 kJ mol–1

10 Which one of the following is not produced in the reaction between concentrated

sulphuric acid and sodium iodide at room temperature?

A I2

B HI

C NaHSO4

D SO3








Section B


11 (a) On the Periodic Table below label the s, p and d blocks.

[2]

(b) Complete the table below indicating whether the elements represented

by the letters X, Y and Z belong to the s, p or d blocks.

[3]

(c) Explain why sodium, an s block element, has a lower fi rst ionisation

energy than chlorine, a p block element.

___________________________________________________________

___________________________________________________________

_______________________________________________________ [2]

(d) (i) Draw the shape of an s orbital.

[1]

(ii) Draw the shape of a p orbital.

[1]

Atom/ion Electronic structure s, p or d block

X 1s22s22p63s23p63d34s2

Y2– 1s22s22p63s23p6

Z+ 1s22s22p63s23p6








12 Nitrogen is a colourless, odourless, unreactive gas which has the formula,

N2. It will react with hydrogen to a small extent if an electric discharge is

passed through a mixture of nitrogen and hydrogen.

(a) Nitrogen gas is believed to be unreactive because it has a triple bond.

(i) In terms of electrons, what is meant by a triple bond?

______________________________________________________

___________________________________________________ [2]

(ii) Why does the presence of a triple bond explain the lack of

reactivity of nitrogen?

______________________________________________________

___________________________________________________ [1]

(iii) Draw the structure of ethyne, C2H

2, which contains a triple bond.

Show all the bonds present.

[1]

(b) Nitrogen reacts with hydrogen in the presence of an iron catalyst to

give a reasonable amount of ammonia. The equation for the reaction is:

N2 + 3H

2 M 2NH

3 ΔH = –58 kJ

(i) Using the value of the enthalpy change, explain whether the

reaction is exothermic or endothermic.

______________________________________________________

___________________________________________________ [1]








(ii) On the axes below, draw an enthalpy level diagram for the

reaction. Label the y-axis.

[3]

(c) Ammonia reacts with acids to form the ammonium ion. In this the

ammonia molecule reacts with a hydrogen ion to form a dative bond.

(i) Write an equation for the formation of the ammonium ion.

___________________________________________________ [1]

(ii) Using a dot and cross diagram, draw the structure of the

ammonium ion.

[2]

(iii) Circle the dative bond in your structure. [1]

Extent of reaction








13 Hydrogen is the most abundant element in the Universe. It is diffi cult to

place hydrogen in the Periodic Table, because it is similar to Group I and

Group VII elements, being neither strongly electropositive, nor

electronegative.

(a) The existence of hydrogen in distant stars has been determined by

spectroscopic analysis. The Lyman series (ultraviolet region) in the

atomic line emission spectrum for hydrogen is shown below.

(i) With reference to specifi c energy levels, explain the origin of the

line labelled A.

______________________________________________________

______________________________________________________

___________________________________________________ [3]

(ii) The line, B, may be used to determine the ionisation energy for

atomic hydrogen. Write an equation, including state symbols, for

the ionisation of a hydrogen atom.

___________________________________________________ [2]

(b) The boiling points of the Group V hydrides are shown below:

frequency

A B

boiling

point/°CNH

3

AsH3

PH3

SbH3








(i) The electronegativity values for nitrogen, phosphorus and

hydrogen are 3.0, 2.1 and 2.1 respectively. Explain what is meant

by the term electronegativity.

______________________________________________________

___________________________________________________ [2]

(ii) Explain why ammonia has a much higher boiling point than

phosphine, PH3, naming the intermolecular forces present between

molecules in the liquid phase in each case.

______________________________________________________

___________________________________________________ [2]

(iii) Phosphine has the following structure.

State and explain the shape of the phosphine molecule.

______________________________________________________

______________________________________________________

___________________________________________________ [3]

(iv) Phosphine may be produced in the laboratory by the action of

water on calcium phosphide, Ca3P

2. Calcium hydroxide is also

produced.

Write an equation for this reaction.

___________________________________________________ [2]

H

PH H








(c) The Group VII hydrides are gaseous at room temperature and pressure,

dissolving in water to form acidic solutions.

(i) Place the hydrides HF, HCl, HBr and HI in order of increasing

hydrogen–halogen bond strength (weakest fi rst).

___________________________________________________ [1]

(ii) Which one of the 1M solutions of the acids HF, HCl, HBr and HI

has the highest pH?

___________________________________________________ [1]

(d) Many metallic hydrides such as sodium hydride, NaH, are ionic solids.

(i) Show the formation of a sodium ion, Na+, and a hydride ion, H–,

from their atoms.

[2]

(ii) State and explain the relative sizes of a hydrogen atom and a

hydride ion.

______________________________________________________

______________________________________________________

___________________________________________________ [2]

(iii) Sodium hydride reacts with water to form sodium hydroxide

solution and hydrogen gas. Write an equation for this reaction,

including state symbols.

___________________________________________________ [2]

(iv) Explain why sodium hydroxide solution conducts electricity.

______________________________________________________

___________________________________________________ [2]








14 Barium ferrate, BaFeO4, may be prepared by bubbling chlorine gas through

a suspension of iron(III) hydroxide in potassium hydroxide solution at

70°C. The mixture is fi ltered and aqueous barium chloride is added to the

fi ltrate to precipitate out barium ferrate which collects as red crystals.

(a) Deduce the oxidation number of iron in the FeO42– ion and use this to

explain the role of chlorine in the reaction.

_________________________________________________________

______________________________________________________ [2]

(b) Using a named silver salt, describe how you would test for the presence

of chloride ions in the solution remaining after separation of the

crystalline product, writing an ionic equation for the reaction.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

______________________________________________________ [4]


(c) State the fl ame colour expected when barium ferrate solution is sprayed

into a fl ame. Assume the presence of iron does not affect the colour.

______________________________________________________ [1]

(d) Ferrate ions react with acid to produce oxygen gas. Describe a test for

oxygen.

_________________________________________________________

______________________________________________________ [2]








(e) The concentration of barium ferrate in solution may be determined in

the following way:

● treatment with excess hydrochloric acid to liberate chlorine gas

● passing the chlorine formed through excess potassium iodide

solution to produce iodine

● titrating the iodine formed with standard sodium thiosulphate

solution:

2Na2S

2O

3(aq) + I

2(aq) → Na

2S

4O

6(aq) + 2NaI(aq)

(i) State what is observed when chlorine gas is bubbled through

potassium iodide solution.

______________________________________________________

___________________________________________________ [2]

(ii) Name the indicator used for the titration and state the colour

change at the end point.

______________________________________________________

___________________________________________________ [3]

(iii) In an experiment, 25.0 cm3 of a barium ferrate solution was treated

with excess hydrochloric acid and the resultant chlorine passed

through excess potassium iodide solution. The iodine liberated

required 23.45 cm3 of 0.1M sodium thiosulphate for complete

reaction. Calculate the concentration of barium ferrate in the

original solution in mol dm–3.

1 mole of ferrate ions is equivalent to 3 moles of thiosulphate ions.

Number of moles of thiosulphate used:

______________________________________________________

Number of moles of barium ferrate in 25.0 cm3 of solution:

______________________________________________________

Concentration of barium ferrate in mol dm–3:

___________________________________________________ [3]








15 The reaction between zinc and copper(II) sulphate solution is exothermic

and is represented by the following equation:

Zn(s) + CuSO4(aq) → ZnSO

4(aq) + Cu(s)

(a) Apart from a temperature rise, give two observations which could be

made.

_________________________________________________________

______________________________________________________ [2]

(b) The enthalpy change for this reaction may be estimated by adding a

known mass of zinc powder to a known volume of standard copper(II)

sulphate solution in a polystyrene cup and measuring the temperature

rise.

A spatula measure of zinc powder was added to a polystyrene cup

containing 75.0cm3 of aqueous copper(II) sulphate at 20.5 °C.

The temperature rose to 33.7 °C. The copper metal formed was

collected, washed, dried and found to weigh 0.42g. Calculate the

enthalpy change for this reaction in kJ per mole of copper formed

using the following headings. (Assume the energy required to raise

1 cm3 of solution by 1 °C is 4.2 J)

temperature rise:

______________________________________________________

energy released in joules:

______________________________________________________

moles of copper formed:

______________________________________________________

energy released per mole of copper in kJ:

___________________________________________________ [4]








[Turn over

16 Phosphorus pentachloride, PCl5, is a pale yellow solid which sublimes

above 160°C and fumes in moist air forming phosphorus oxychloride,

POCl3, and hydrogen chloride.

Phosphorus pentachloride may be prepared in the laboratory using the

apparatus below.

(a) (i) Suggest the function of the anhydrous calcium chloride tube.

___________________________________________________ [1]

(ii) State one essential safety precaution for this preparation.

___________________________________________________ [1]

(iii) Suggest why the reaction vessel must be cooled.

___________________________________________________ [1]

(iv) Write the equation for the reaction between phosphorus trichloride

and chlorine gas.

___________________________________________________ [1]

(b) (i) State the Octet Rule.

___________________________________________________ [1]

(ii) Suggest why PCl5 does not obey the Octet Rule.

___________________________________________________ [1]

anhydrous

calcium chloride tube

phosphorus trichloride

freezing mixture of ice

and water

phosphorus

pentachloride

dry chlorine








(c) Phosphorus pentachloride reacts with oxalic acid, H2C

2O

4, to form

phosphorus oxychloride as shown in the equation:

H2C

2O

4(s) + PCl

5(s) → POCl

3(s) + 2HCl(g) + CO

2(g) + CO(g)

(i) Suggest why this reaction is often used to synthesise pure

phosphorus oxychloride.

___________________________________________________ [1]

(ii) Phosphorus oxychloride is readily hydrolysed by water according

to the equation:

POCl3 + 3H

2O → H

3PO

4 + 3HCl

Calculate the volume of hydrogen chloride gas formed, at 20°C

and 1 atmosphere pressure, when 1.53g of phosphorus oxychloride

is hydrolysed.

______________________________________________________

______________________________________________________

___________________________________________________ [3]









2008

MARKSCHEME


assessing


[ASC11]









Section A

1 C

2 A

3 C

4 A

5 C

6 C

7 B

8 C

9 A

10 D


Section A 20








Section B

11 (a) s, d and p in correct order [2]

(b) d [1] p [1] s [1] [3]

(c) in sodium electron further from nucleus/in chlorine closer or Cl radius smaller than Na [1] conditional on comment on size less attraction/more attraction (of nucleus for electrons) [1] [2]

(d) (i) [1]

(ii) or or [1] 9

12 (a) (i) three (bonding/shared) pairs of electrons between the nitrogens [2] [2]

(ii) a lot of energy needed to break the bonds [1] [1]

(iii) H C C H [1]

(b) (i) It is negative hence exothermic [1]

(ii)

[3]

(c) (i) NH3 + H+ → NH4+ [1]

(ii)

[2]

(iii)

each error [–1] [1] 12

H• ×

• × H •

× N •• H

H• ×

• × H •

× N •• H

H

H

enthalpyreactants

products








13 (a) (i) electron excited to n = 2 [1] electron falls to n = 1 [1] (u.v.) radiation emitted [1] [3]

(ii) H(g) → H+(g) + e– [2]

(b) (i) ability of an atom to attract electrons [1] in a covalent bond [1] [2]

(ii) ammonia has hydrogen bonding and phosphine has van der Waals [1] comment on strength of hydrogen bonds [1] [2]

(iii) pyramidal [1] electron pairs repel [1] move as far apart as possible [1] [3]

(iv) Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3 [2] unbalanced = [1]

(c) (i) HI, HBr, HCl, HF [1]

(ii) HF [1] (d) (i) Na° + Hx → [Na]+ [H°]– [2]

x

(ii) hydride ion is larger [1] electrons repel [1] [2]

(iii) NaH(s) + H2O(l) → NaOH(aq) + H2(g) [2]

(iv) ions [1] are able to move [1] [2] 24

14 (a) +6 [1] chlorine is an oxidising agent [1] [2]

(b) add nitric acid [1] add silver nitrate solution [1] white precipitate [1] add NH3(aq) [1] ppt. dissolves [1] to a maximum of [3] Ag+ + Cl– → AgCl [1] – essential [4]


(c) green [1]

(d) glowing splint [1] relights [1] [2]

(e) (i) solution changes from colourless [1] to brown (not clear)/black solid [1] [2]








(ii) starch [1] blue/black [1] to colourless [1] [3]

(iii) moles of thiosulphate = 23.45 × 0.1/1000 = 2.345 × 10–3 [1] moles of ferrate = 2.345 × 10–3/3 = 7.82 × 10–4 [1] concentration = 7.82 × 10–4 × 40 = 3.12 × 10–2 M = 0.0312 M [1] [3] 19 each error [–1], carry error through

15 (a) red/brown solid forms [1] blue solution [1] fades/goes colourless [1] zinc dissolves [1] Any two from four [2]

(b) 13.2 °C 13.2 × 75 × 4.2 = 4.158 kJ 0.00656 mole Cu 633.8 kJ mol–1 [4] 6

16 (a) (i) prevents moisture entering the apparatus (and reacting with PCl5) [1]

(ii) fume cupboard [1]

(iii) prevent loss/sublimation of product [1]

(iv) PCl3 + Cl2 → PCl5 [1]

(b) (i) 8 electrons in the outer shell [1]

(ii) 10 electrons (in the outer shell around phosphorus atom) [1]

(c) (i) other products are gases [1]

(ii) moles POCl3 = 1.53/153.5 = 0.01 moles HCl = 0.03 volume HCl = 0.03 × 24 = 0.72 dm3

error [–1] [3] 10

Section B 80

Total 100









January 2009

Chemistry


assessing


[ASC11]

FRIDAY 16 JANUARY, MORNING

TIME

1 hour 30 minutes.



Quality of written communication will be assessed in question 15(d)(ii).In section A all questions carry equal marks, i.e. two marks for each

question.

In Section B the figures in brackets printed down the right-hand side of

pages indicate the marks awarded to each question or part question.


















Section A



1 How many electrons are there in a calcium ion?

A 18

B 20

C 22

D 40

2 Which one of the following equations represents the second ionisation energy for barium?

A Ba(s) → Ba2+(g) + 2e–

B Ba(g) → Ba2+(g) + 2e–

C Ba+ (s) → Ba2+(g) + e–

D Ba+ (g) → Ba2+(g) + e–

3 Which one of the following shows the trend in electronegativity values of the elements in the

Periodic Table?

Across a Period Down a Group

A decrease decrease

B decrease increase

C increase decrease

D increase increase

4 The element astatine lies immediately below iodine in the Periodic Table and is likely to

A be pale yellow.

B be a volatile liquid at room temperature and pressure.

C form a hydride which dissolves in water to give an acidic solution.

D oxidise iodide ions to iodine.








5 If the price of one tonne (1000 kg) of sulphur is £160, what is the cost (to the nearest pound)

of the sulphur needed to make one tonne of sulphuric acid, H2SO4?

A £52

B £98

C £160

D £490

6 Which one of the following does not obey the octet rule?

A beryllium chloride

B carbon dioxide

C nitrogen

D oxygen

7 The orbitals of a nitrogen atom may be represented as shown.

1s 2s 2p

Which one of the following diagrams represents the arrangement of electrons in the ground

state of the nitrogen atom?

A

B

C

D






8 Potassium iodide is formed when potassium is warmed in iodine vapour. Which one of the

following describes the bonding in the three species?

potassium iodine potassium iodide

A ionic covalent ionic

B metallic ionic covalent

C covalent covalent ionic

D metallic covalent ionic

9 Which one of the following has a bond angle of 109.5 °?

A BeCl2 B BF3 C CH4 D CO2

10 Which one of the following ions has the largest radius?

A F–

B Mg2+

C Na+

D O2–






Section B


11 Metal ions are responsible for the flame colours produced by fireworks.

Complete the table below by inserting the flame colour for each metal ion.

metal ion flame colour

barium

potassium

sodium [3]

12 (a) Draw the shape of an s orbital.

[1]

(b) Draw the shape of a p orbital.

[1]

(c) Write the electronic configuration of a carbon atom in terms

of s and p electrons.

_______________________________________________________[1]






13 Rocket fuels need to supply a large amount of energy, yet have a low mass

i.e. a high power to weight ratio.

One potential reaction is that of fluorine with diborane (a boron hydride).

H H

� �H�B�B�H

� �H H

diborane

B2H6(g) + 6F2(g) → 6HF(g) + 2BF3(g)

(a) Calculate the enthalpy change when one mole of diborane reacts

completely with fluorine given the following bond enthalpies.

bond kJ mol–1

F�F 158

B�H 389

B�B 293

H�F 566

B�F 627

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]

(b) The bond enthalpy of hydrogen fluoride is 566 kJ mol–1 whereas that

of hydrogen iodide is 299 kJ mol–1. State what would be observed

when hydrogen fluoride and hydrogen iodide are heated.

_________________________________________________________

_________________________________________________________

_______________________________________________________[2]






(c) Both HF and BF3 are gases at room temperature and pressure.

Calculate the total volume of gas produced at 20 °C and one

atmosphere pressure by the complete reaction of 7.0 g of diborane

with fluorine.

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]

(d) State and explain the shape of the boron trifluoride molecule.

_________________________________________________________

_________________________________________________________

_______________________________________________________[2]






14 Silver occurs in nature as the sulphide and as the chloride. Although

less malleable and less ductile than gold, the thermal and electrical

conductivities of silver are greater than those of any other metal.

(a) (i) Explain what is meant by the term malleability.

_____________________________________________________

___________________________________________________[2]

(ii) Explain what is meant by the term ductility.

_____________________________________________________

___________________________________________________[2]

(iii) Explain how silver is able to conduct electricity.

_____________________________________________________

___________________________________________________[2]

(b) Silver hardly dissolves in hydrochloric acid and the usual method of

preparing silver chloride is to add a soluble metal chloride to a soluble

silver salt.

(i) Name a soluble metal chloride.

___________________________________________________[1]

(ii) Name a soluble silver salt.

___________________________________________________[1]

(iii) Write an ionic equation for the reaction of the soluble metal

chloride with the soluble silver salt.

___________________________________________________[1]






(c) Silver chloride is virtually insoluble in water. Its solubility has been

estimated at 0.4 mg in 250 cm3 of water at 20 °C.

(i) Calculate the molar mass of silver chloride.

___________________________________________________[1]

(ii) Calculate the number of moles of silver chloride in 0.4 mg.

_____________________________________________________

___________________________________________________[1]

(iii) Calculate the solubility of silver chloride in moles per litre.

_____________________________________________________

___________________________________________________[1]

(iv) Name a solution that will dissolve silver chloride.

___________________________________________________[1]

(v) Silver chloride is a white solid which is affected by light. Describe

and explain the effect of light on silver chloride.

_____________________________________________________

___________________________________________________[2]






15 Bromine was discovered in the residues from the manufacture of sea salt at

Montpellier, France. The residues contain magnesium bromide.

The addition of chlorine liberates bromine.

(a) (i) Write an equation for the reaction of chlorine with magnesium

bromide.

___________________________________________________[2]

(ii) Using electron transfer, explain why this can be considered to be

redox reaction.

_____________________________________________________

_____________________________________________________

___________________________________________________[3]

(b) Bromine, Br2, is a liquid at room temperature. Liquid bromine has a

high density: 1 mol of bromine, Br2, occupies 51 cm3.

Calculate the density of bromine in g cm–3.

_________________________________________________________

_______________________________________________________[2]

(c) Bromine is miscible with organic (non-aqueous) solvents, but its

solubility in water is more limited. The aqueous solution is known as

bromine water.

(i) Name a non-aqueous solvent that dissolves bromine.

___________________________________________________[1]

(ii) Compare the solubilities of bromine and iodine in water.

_____________________________________________________

___________________________________________________[2]






(iii) Describe what would be observed if bromine water was added

to concentrated solutions of sodium chloride and sodium iodide

respectively.

sodium chloride

___________________________________________________[1]

sodium iodide

___________________________________________________[1]

(d) The concentration of bromine in solution can be measured by reacting

it with excess potassium iodide and the liberated iodine determined by

titration with sodium thiosulphate solution.

(i) Write an equation for the reaction of sodium thiosulphate with

iodine.

___________________________________________________[2]

(ii) Describe, with practical details, how this titration would be

carried out. Assume all the apparatus is clean and dry. Details of

calculations are not required.

_____________________________________________________

_____________________________________________________

_____________________________________________________

_____________________________________________________

_____________________________________________________

_____________________________________________________

_____________________________________________________

___________________________________________________[5]









16 Oxalic acid (ethanedioic acid) is a weak dicarboxylic acid. It is similar to

acetic acid (ethanoic acid) which is a monocarboxylic acid.

COOH CH3COOH

� COOH

oxalic acid acetic acid

(a) Suggest the meaning of the term dicarboxylic acid.

_________________________________________________________

_______________________________________________________[2]

(b) Both of these acids react with alkalis. Write the equation for the

reaction of oxalic acid with excess sodium hydroxide.

_______________________________________________________[2]

(c) Weak acids such as oxalic acid can be titrated with strong alkalis using

phenolphthalein as indicator.

State the colour of phenolphthalein in alkaline and acidic solution.

alkali __________________________________________________[1]

acid ___________________________________________________[1]








(d) If oxalic acid reacts with a small amount of phosphorus pentachloride

a mixture of gases is produced:

How could you show that hydrogen chloride was a product?

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]

(e) Oxalic acid reacts with excess phosphorus pentachloride to give oxalyl

chloride which has the following percentage composition by mass.

element % composition

carbon 18.9

chlorine 55.9

oxygen 25.2

Calculate the empirical formula of oxalyl chloride.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]








17 Hydrogen has three naturally occurring isotopes: protium 1H, deuterium 2H

and tritium 3H.

(a) Draw an atom of tritium showing and labelling all the sub-atomic

particles present.

[2]

(b) The relative atomic mass of hydrogen is 1.0079. Explain which one of

the hydrogen isotopes is the most abundant in nature.

_________________________________________________________

_______________________________________________________[2]








(c) When subjected to an electrical discharge, atomic hydrogen emits

electromagnetic radiation due to electron transitions between energy

levels. The line emission spectrum of atomic hydrogen in the visible

region is shown below.

(i) To which energy level do electrons return to produce lines in the

visible region?

___________________________________________________[1]

(ii) The lowest frequency line in the visible region has a frequency of

4.568 × 1014 s–1. Calculate the energy in kJ mol–1, associated with

this frequency.

_____________________________________________________

_____________________________________________________

_____________________________________________________

___________________________________________________[3]

(iii) Explain why the lines in the emission spectrum converge.

_____________________________________________________

___________________________________________________[2]

frequency








(d) It is possible to obtain pure deuterium oxide, D2O, from sea water.

Deuterium oxide boils at 101.4 °C compared to 100.0 °C for water.

(i) Name the two types of intermolecular forces which exist between

water molecules.

_____________________________________________________

___________________________________________________[2]

(ii) Draw a dot and cross diagram to show the bonding in deuterium

oxide, D2O, showing all the outer shell electrons.

[2]

(iii) Deuterium oxide can combine with deuterium ions, D+, to form

D3O+. Write an equation for this reaction.

___________________________________________________[1]

(iv) Name the type of bond formed between the oxygen atom and the

deuterium ion.

___________________________________________________[1]









January 2009

Chemistry



[ASC11]

FRIDAY 16 JANUARY, MORNING

MARKSCHEME

Not to be circulated beyond the Examining Team








Section A

1 A

2 D

3 C

4 C

5 A

6 A

7 D

8 D

9 C

10 D


Section A 20








Section B

11 green [1] lilac [1] yellow/orange [1]

12 (a) [1]

(b) or or [1]

(c) 1s22s22p2 [1] 6

13 (a) bonds broken

1B—B = 1 × 293 = 293 6B—H = 6 × 389 = 2334 6F—F = 6 × 158 = 948 total = 3575

bonds formed

6H—F = 6 × 566 = 3396 6B—F = 6 × 627 = 3762 total = 7158

enthalpy change = 3575 – 7158 = –3583 [3]

(b) hydrogen fluoride; no change observed [1] hydrogen iodide; violet vapour observed [1]

(c) B2H6 = 2 × 11 + 6 × 1 = 22 + 6 = 28

1 mol of diborane forms 8 mol of gases 0.25 mol of diborane forms 2.0 mol of gases

2.0 mol of gases occupies 2.0 × 24 dm3 = 48 dm3 [3]

(d) trigonal planar [1] bonding pairs repel equally [1] 10

7.0g of = =728

0 25. moldiborane








14 (a) (i) can be hammered/bent [1] into a new shape [1] [2]

(ii) can be stretched [1] into wires [1] [2]

(iii) electrons delocalised [1] move and current flows [1]

(b) (i) e.g. sodium chloride [1]

(ii) silver nitrate [1]

(iii) Ag+ + Cl– → AgCl [1]

(c) (i) AgCl = 108 + 35.5 = 143.5 [1]

(ii) [1]

(iii) 4 × 2.79 × 10–6 = 1.116 × 10–5 = 1.12 × 10–5 [1]

(iv) ammonia solution [1]

(v) goes grey/black [1] silver is formed/light energy causes electron transfer [1] 15

15 (a) (i) MgBr2 + Cl2 → MgCl2 + Br2 [2]

(ii) Br– loses electrons [1] Cl2 receives electrons [1] loss of electrons is oxidation and gain of electrons is reduction [1]

(b) 1 mol Br2 = 2 × 80 = 160 g 160 g occupy 51 cm3

density = 3.14 g cm–3 [2]

(c) (i) e.g. hexane [1]

(ii) bromine is slightly/moderately soluble [1] iodine is insoluble [1]

(iii) sodium chloride: no effect/orange colour diluted [1] sodium iodide: brown/black colour produced [1]

0 4 10143 5

2 79 103

6..

.–

–=× ×

16051g occupy 1cm3








(d) (i) 2Na2S2O3 + I2 → Na2S4O6 + 2NaI [2]

(ii) iodine in flask and thiosulphate in burette [1] add thiosulphate until straw yellow [1] add starch [1] add thiosulphate until the blue/blue-black colour [1] disappears/colourless [1]


16 (a) two [1] carboxylic acid groups [1]

(b) [2]

(c) alkali: pink/red [1] acid: colourless [1]

(d) concentrated [1] ammonia solution [1] white fumes [1]

(e) element moles ratio

C 1

Cl 1

O 1

empirical formula = COCl [3] 12

COOH COONa| + 2NaOH → | + 2H2OCOOH COONa

18 912

1 575

55 935 5

1 575

25 216

1 575

. .

.

..

. .

=

=

=








17 (a)

[2]

(b) protium [1] value nearest to 1 [1]

(c) (i) second energy level [1]

(ii) E = hf = 6.63 × 10–34 × 4.568 × 1014 = 3.0286 × 10–19 J

For one mole = 3.0286 × 10–19 × 6.0 × 1023 J

= 1.82 × 105 J mol–1

= 182 kJ mol–1 [3]

(iii) the higher energy levels get closer together [2]

(d) (i) hydrogen bonding [1] van der Waals forces [1]

(ii) x x D x O x D [2] x x

(iii) D2O + D+ → D3O+ [1]

(iv) dative bond [1] 16

Section B 80

Total 100

electron

proton

neutrons








TIME

1 hour 30 minutes.




Quality of written communication will be assessed in question 16(e).In Section A all questions carry equal marks, i.e. two marks for each

question.

In Section B the figures in brackets printed down the right-hand side












2009

Chemistry


assessing


[ASC11]









Section A

For each of the following questions only one of the lettered responses (A–D) is correct.


1 20 cm3 of 0.3 mol dm–3 potassium hydroxide solution is exactly neutralised by

A 10 cm3 of 0.3 mol dm–3 sulphuric acid.

B 10 cm3 of 0.6 mol dm–3 sulphuric acid.

C 20 cm3 of 0.3 mol dm–3 sulphuric acid.

D 20 cm3 of 0.6 mol dm–3 sulphuric acid.

2 A positively charged particle with the electron configuration 1s22s22p6 is

A an aluminium ion.

B a fluoride ion.

C an oxide ion.

D a potassium ion.

3 Which one of the following molecules contains a triple bond?

A C2H

4

B CO2

C N2

D NF3

4 Which one of the following sodium compounds produces a gas when treated with dilute

sulphuric acid?

A sodium carbonate

B sodium chloride

C sodium fluoride

D sodium iodide








5 Which one of the following contains the name of the reagent and that of the indicator used

in an iodine titration?

A sodium sulphate and starch

B sodium sulphate and methyl orange

C sodium thiosulphate and starch

D sodium thiosulphate and methyl orange

6 Which one of the following electron configurations has two unpaired electrons?

A 1s22s2

B 1s22s22p3

C 1s22s22p4

D 1s22s22p63s23p5

7 Which area of the Periodic Table contains elements which have only s electrons in their

outer shells?

8 Which one of the following chloro-compounds is non-polar?

A HCl

B CCl4

C CH3Cl

D CHCl3

AB

C D









A N2

B NH3

C NH2

–

D NH4

+

10 The enthalpy of neutralisation when an acid reacts with an alkali is, by

definition, the number of kilojoules released by

A the formation of one mole of salt.

B the formation of one mole of water.

C the neutralisation of one mole of acid.

D the neutralisation of one mole of alkali.







Section B


11 The electronic energy levels of atomic hydrogen are shown below. Draw an

arrow on the diagram which represents the energy change associated with

the lowest frequency line in the ultraviolet emission spectrum.

[3]

12 The electronegativity of atoms causes bonds to be polar. Indicate the

polarity of the following bonds. The first one has been completed for you.

δ+ δ–

C�Cl

O�H

Cl�Br

N��O

[3]







13 The female of the American cockroach (Periplaneta americana) secretes a

chemical (pheromone) of molecular formula, C11

H18

O2, to which the male

of the species is attracted. It is reported that the male may respond to as

few as 60 molecules of the pheromone.

What is the mass, in grams, of these 60 molecules?

Use the following headings to assist you in your calculation.

relative molecular mass

____________________________________________________________

____________________________________________________________

mass of one mole

____________________________________________________________

mass of one molecule

____________________________________________________________

____________________________________________________________

mass of sixty molecules in grams

____________________________________________________________

__________________________________________________________[4]







14 Boron forms giant covalent structures with other elements, for example,

boron nitride, BN. It is claimed that boron nitride is as hard as diamond.

(a) (i) Explain why diamond is so hard.

_____________________________________________________

_____________________________________________________

___________________________________________________[2]

(ii) Explain why graphite is so soft.

_____________________________________________________

_____________________________________________________

___________________________________________________[2]

(iii) State one other physical property, apart from hardness, which you

would expect boron nitride to possess.

___________________________________________________[1]







(b) Both boron and carbon combine with fluorine. Boron forms boron

trifluoride (BF3) and carbon forms carbon tetrafluoride, (CF

4).

(i) Write an equation for the formation of boron trifluoride from

boron and fluorine.

___________________________________________________[2]

(ii) Write an equation for the formation of carbon tetrafluoride from

methane and fluorine, the other product being hydrogen fluoride.

___________________________________________________[2]

(iii) Use dot and cross notation to draw the structures of boron

trifluoride and carbon tetrafluoride, showing outer electrons only.

[2]

(iv) Explain the octet rule and comment on its application to boron

trifluoride and carbon tetrafluoride.

_____________________________________________________

_____________________________________________________

_____________________________________________________

___________________________________________________[4]

(v) Draw the shapes of BF3 and CF

4 and explain them in terms of

their electron structure.

_____________________________________________________

_____________________________________________________

___________________________________________________[4]







15 Calcium fluoride, CaF2, occurs naturally as fluorite or fluorspar. Impurities

give a blue variety known as Blue John. Fluorspar is the major source of

hydrogen fluoride and fluorine.

It can be prepared in the laboratory by precipitation or direct combination

of the elements.

(a) What is the colour of pure calcium fluoride?

_______________________________________________________[1]

(b) Explain the formation of calcium fluoride from calcium and fluorine

atoms using dot and cross diagrams showing outer electrons only.

[4]

(c) Write an equation for the precipitation of calcium fluoride by mixing

solutions of calcium chloride and sodium fluoride.

_______________________________________________________[1]

(d) Calcium fluoride reacts with concentrated sulphuric acid to form

hydrogen fluoride and calcium sulphate. Write an equation for the

reaction.

_______________________________________________________[2]







(e) Calcium fluoride is sparingly soluble in water; 0.0025 g dissolves in

100 cm3 of water at 18°C.

Calculate the concentration of fluoride ions in moles per litre using the

following headings:

(i) relative formula mass of calcium fluoride

_____________________________________________________

(ii) number of moles of calcium fluoride in 0.0025 g

_____________________________________________________

(iii) number of moles of fluoride ion in 0.0025 g of calcium fluoride

_____________________________________________________

(iv) number of moles of fluoride ion in 100 cm3 of water

_____________________________________________________

(v) number of moles of fluoride ion in 1000 cm3 of water

___________________________________________________[5]

(f) The presence of fluoride ions in domestic water supplies is regarded

as beneficial by some, but the deliberate addition of fluoride ions is

controversial.

(i) State one benefit of fluoride ions in drinking water.

___________________________________________________[1]

(ii) Explain why some people object to the addition of fluoride ions to

drinking water.

_____________________________________________________

___________________________________________________[1]







16 All naturally occurring sodium atoms have a relative atomic mass of 23

i.e. the atoms are represented by the symbol 23Na. However, radioactive

isotopes of sodium, e.g. 24Na, may be prepared.

(a) (i) State the number of electrons, protons and neutrons in an atom of 23Na.

_____________________________________________________

___________________________________________________[2]

(ii) Explain why 23Na and 24Na are regarded as isotopes.

_____________________________________________________

___________________________________________________[2]

(b) A sample of sodium from a nuclear reactor contains 2.00% of 24Na

and 98.00% of 23Na by mass. Calculate the relative atomic mass of the

sample to two decimal places.

_________________________________________________________

_________________________________________________________

_______________________________________________________[2]







(c) (i) A major use of sodium metal is in street lamps. The lamp contains

mercury vapour which conducts electricity at high voltages.

Sodium within the lamp vaporises and the electrical energy causes

yellow (orange) light to be given out. When the light from the

sodium lamp is analysed, the spectrum shows two bright yellow

lines at wavelengths of 589 nm and 589.6 nm.

(1 nm = 1 × 10–9 m)

589 589.6

λ →

Using the equations E = hv and c = vλ, calculate the energy

change (in joules) associated with the line at 589 nm.

(c = 3 × 108 m s–1).

_____________________________________________________

_____________________________________________________

_____________________________________________________

___________________________________________________[3]

(ii) Explain how you could carry out a flame test and a test for

chloride ions to identify a white solid as sodium chloride. Write

equations for any reactions taking place.

_____________________________________________________

_____________________________________________________

_____________________________________________________

_____________________________________________________

___________________________________________________[5]







(d) If larger amounts of energy are supplied to sodium vapour (gas) it

ionises.

(i) Write the equation which represents the first ionisation energy of

sodium including state symbols.

___________________________________________________[2]

(ii) The value of the first ionisation energy for sodium is

500 kJ mol–1. The second ionisation energy has a value of 4513 kJ

mol–1. Explain why this is a much higher value.

_____________________________________________________

_____________________________________________________

___________________________________________________[2]

(e) Using diagrams, explain why sodium is able to conduct electricity

whether solid or molten, while sodium chloride conducts only when

molten or dissolved in water.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________[5]








17 In the laboratory, ammonia can be prepared by heating a mixture of

ammonium chloride and calcium hydroxide as shown in the diagram below.

The equation for the reaction is:

2NH4Cl + Ca(OH)

2 → 2NH

3 + CaCl

2 + 2H

2O

(a) The ammonia gas is collected upwards. Calculate the relative

molecular masses of ammonia, NH3, oxygen, O

2 and nitrogen, N

2, and

use them to explain why ammonia is collected in this way.

ammonia ___________________________________

oxygen _____________________________________

nitrogen _____________________________________

explanation _______________________________________________

_______________________________________________________[2]

(b) Calculate the volume of ammonia produced, at 20 °C and one

atmosphere pressure, if 1.07 g of ammonium chloride are heated with

excess calcium hydroxide.

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]

ammonium chloride+

calcium hydroxide

heat

ammonia







(c) Ammonia gas is alkaline.

(i) Describe the effect of ammonia on moist Universal Indicator

paper.

___________________________________________________[1]

(ii) Ammonia may be detected using concentrated hydrochloric acid.

Write the equation for the reaction and describe what is observed.

_____________________________________________________

___________________________________________________[2]

(d) Ammonia can act as a reducing agent. When passed over heated

copper(II) oxide, the following reaction occurs:

2NH3 + 3CuO → 3Cu + N

2 + 3H

2O

Deduce the oxidation numbers of nitrogen and copper in the reactants

and products and use them to explain the redox change.

_________________________________________________________

_________________________________________________________

_______________________________________________________[3]







ADVANCEDGeneral Certificate of Education

2009

Chemistry



[ASC11]


MARKSCHEME







Section A

1 A

2 A

3 C

4 A

5 C

6 C

7 A

8 B

9 D

10 B


Section A 20







Section B

11 arrow down [1] from level 2 [1] to level 1 [1] 3

12 δ− δ+ O — H [1]

δ− δ+ Cl — Br [1]

δ+ δ− N O [1] 3

13 relative molecular mass C11H18O2 = 11 × 12 + 18 × 1 + 2 × 16 = 132 + 18 + 32 = 182 [1]

182 g [1]

182/6 × 1023 = 30.3 × 10–23 = 3.02 × 10–22 g [1]

60 × 3.02 × 10–22 g = 181.2 × 10–22 g = 1.812 × 10–20 g [1] 4







14 (a) (i) strong bonds [1] (between carbon atoms) and giant structure [1] [2]

(ii) layers [1] (of carbon atoms) can slide over each other/ weak van der Waals [1] [2]

(iii) e.g. high melting point/boiling point [1] [1]

(b) (i) 2B + 3F2 → 2BF3 [2]

(ii) CH4 + 4F2 → CF4 + 4HF [2]

(iii)

[1]

[1]

(iv) octet rule: eight [1] outer [1] electrons (around a central atom) [2] boron has 6 [1] carbon has 8 [1]

(v)

[1]

[1]

bonding electron (pairs) repel [1] [1]

as far apart as possible [1] [1] 19

F

F FB

F

F

F

F

C

F

BFF

F

F

CFF







15 (a) white [1]

(b)

[4]

(c) CaCl2 + 2NaF → CaF2 + 2NaCl [1]

(d) CaF2 + H2SO4 → CaSO4 + 2HF [2]

(e) (i) CaF2 = 40 + 2 × 19 = 78 [1]

(ii) 0.0025/78 = 3.2 × 10–5 [1]

(iii) 2 × 3.2 × 10–5 = 6.4 × 10–5 [1]

(iv) 6.4 × 10–5 [1]

(v) 10 × 6.4 × 10–5 = 6.4 × 10–4 [1]

(f) (i) helps to prevent tooth decay [1]

(ii) takes away freedom of choice [1] 15

F

F

Ca

F

F

Ca2+

–

–







16 (a) (i) 11 electrons 11 protons 12 neutrons [2]

(ii) They have the same number of protons [1] different number of neutrons [1]

(b) 2 × 24 = 48 98 × 23 = 2254 total = 2302

relative atomic mass = 23.02 [2]

(c) (i) c = νλ 3 × 108 = ν × 589 × 10–9

ν = 5.093 × 1014

E = hν = 6.63 × 10–34 × 5.093 × 1014 = 3.38 × 10–19 J [3]

(ii) nichrome wire cleaned with concentrated hydrochloric acid [1] placed in blue Bunsen flame [1] yellow [1] dissolve in water [1] add silver nitrate solution [1] to produce a white precipitate [1] to a maximum of [4] NaCl + AgNO3 → NaNO3 + AgCl [1] [5]

(d) (i) Na(g) → Na+(g) + e– [2]

(ii) (more energy is needed to) remove an electron from a full shell [1] the electron(s) are nearer to the nucleus [1]

from a positive ion [1] Any 2 from 3 [2]

(e) sodium metal has delocalised electrons [1] which move and thus a current flows [1] sodium chloride consists of ions [1] in fixed positions which cannot move [1] can move in solution or when molten [1]








17 (a) ammonia = 17 oxygen = 32 nitrogen = 28 [1]

ammonia is lighter [1]

(b) NH4Cl = 14 + 4 + 35.5 = 53.5 1.07/53.5 = 0.02 mol volume = 0.02 × 24 dm3 = 0.48 dm3 [3]

(c) (i) blue/violet [1]

(ii) NH3 + HCl → NH4Cl [1] White smoke/fumes [1]

(d) reactants: Cu = +2 N = –3 [1] products: Cu = 0 N = 0 [1] copper has been reduced and nitrogen has been oxidised [1] 11

Section B 80

Total 100







TIME

1 hour 30 minutes.


Write your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all seventeen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all seven questions in Section B. Write your answers in the spaces provided in this question paper.


The total mark for this paper is 100.Quality of written communication will be assessed in question 16(a)(i).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in the brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.

New

Specifi

catio

n

ADVANCED SUBSIDIARY (AS)

General Certificate of Education

2009

Chemistry


Inorganic Chemistry

[AC111]








Section A

For each of the following questions only one of the lettered responses (A – D) is correct.

Select the correct response in each case and mark its code letter by connecting the dots as illustrated on the answer sheet

1 How many electrons are present in a potassium ion, K+?

A 18 B 19 C 20 D 39

2 Which one of the following represents the first five ionisation energies in kJ mol–1 of an s-block element?

1st 2nd 3rd 4th 5th

A 580 1800 2700 11 600 14 800 B 740 1500 7700 10 500 13 600 C 1000 2300 3400 4600 7000 D 14 800 11 600 2700 1800 580

3 A sample of 4.64 g of hydrated sodium carbonate, Na2CO3.xH2O, was dissolved in 1 dm3 of water. 25.0 cm3 of this solution required 20.0 cm3 of 0.05 mol dm–3 hydrochloric acid for neutralisation. Which one of the following is the value of x?

A 0.5 B 5 C 7 D 13


A Ammonium, NH4+

B Boron trifluoride, BF3 C Sulphur hexafluoride, SF6 D Water, H2O







5 Which one of the following lists the colour of solid iodine and of iodine dissolved in the solvent stated?

Solid Water Hexane

A grey/black purple yellow/brown

B dark purple yellow/brown purple

C yellow/brown grey/black yellow/brown

D grey/black yellow/brown purple

6 Which one of the following does not show the number of each bond present in the named molecules?

MoleculeSingle

bond

Double

bond

Triple

bond

A Ethene, C2H4 2 1 0

B Nitrogen, N2 0 0 1

C Carbon dioxide, CO2 0 2 0

D Beryllium chloride, BeCl2 2 0 0

7 In which one of the following molecules does the named element have two lone pairs of electrons?

A Beryllium in BeCl2 B Carbon in CH4 C Nitrogen in NH3 D Oxygen in H2O







8 Using the half-equations below, which one of the following is the balanced ionic equation for the reaction between acidified manganate(VII) ions and ethanedioate ions?

Acidified manganate(VII) ions:

MnO4– + 8H+ + 5e– → Mn2+ + 4H2O

Ethanedioate ions:

C2O42– → 2CO2 + 2e–

A 2MnO4– + 16H+ + C2O4

2– → 2Mn2+ + 8H2O + 2CO2

B MnO4– + 8H+ + 5C2O4

2– → Mn2+ + 4H2O + 10CO2

C 2MnO4– + 16H+ + 5C2O4

2– → 2Mn2+ + 8H2O + 10CO2

D 5MnO4– + 40H+ + 2C2O4

2– → 5Mn2+ + 20H2O + 4CO2

9 Which one of the following molecules is non-polar?

A Ammonia, NH3 B Carbon dioxide, CO2 C Hydrogen fluoride, HF D Water, H2O

10 The extraction and purification of uranium from its ore involves the following reaction between uranium(IV) fluoride and magnesium.

2Mg + UF4 → U + 2MgF2

What mass of uranium can be extracted from 500 tonnes of uranium(IV) fluoride and 50 tonnes of magnesium?

A 192 tonnes B 246 tonnes C 379 tonnes D 495 tonnes







Section B

Answer all seven questions in this section

11 (a) Complete the table naming the strongest intermolecular force between molecules in each of the following liquids.

Liquid Intermolecular force

Ammonia, NH3(l)

Hydrogen chloride, HCl(l)

Methane, CH4(l) [3]

(b) Explain why ice has a lower density than water.

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [2]

(c) Draw and explain the shape of an ammonia molecule.

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [3]







12 Neon has several isotopes.

(a) Complete the table below.

Number of

protons

Number of

electrons

Number of

neutrons

Neon-20

Neon-21

Neon-22 [2]

(b) The table below gives the abundance of each isotope of neon.

Calculate the relative atomic mass of neon to two decimal places.

Isotope % abundance

Neon-20 90.92

Neon-21 0.26

Neon-22 8.82

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [2]

(c) Name the isotope used as the standard to compare the relative atomic mass of atoms.

___________________________________________________________ [1]

(d) Label the sub-shells below and draw the electronic structure of neon in the ground state.

[2]







(e) Draw the shape of an s and of a p orbital.

s orbital p orbital [2]







13 The percentage of calcium carbonate present in egg shells can be found by back titration using excess hydrochloric acid and standard sodium hydroxide solution.

(a) Write an equation for the reaction between calcium carbonate and hydrochloric acid.

___________________________________________________________ [2]

(b) Explain what is meant by a standard solution.

_____________________________________________________________

___________________________________________________________ [1]

(c) 1.12 g of an egg shell was reacted with 20.0 cm3 of 2M hydrochloric acid and the solution formed made up to 250 cm3 in a volumetric flask. 25.0 cm3 of this solution completely reacted with 18.6 cm3 of 0.1 M sodium hydroxide.

Calculate the percentage of calcium carbonate in the egg shell using the headings below.

Moles of hydrochloric acid added to the egg shell

_____________________________________________________________

Moles of sodium hydroxide used

_____________________________________________________________

Moles of hydrochloric acid in 250 cm3

_____________________________________________________________

Moles of hydrochloric acid which reacted with the egg shell

_____________________________________________________________

Mass of calcium carbonate in the egg shell

_____________________________________________________________

Percentage of calcium carbonate in the egg shell

___________________________________________________________ [6]







(d) Name a suitable indicator for the titration of hydrochloric acid with sodium hydroxide solution. Give the colour change observed at the end point.

Indicator: __________________________

Colour change:

from __________________________ to __________________________ [3]







14 The Periodic Table identifies various relationships between elements.

(a) (i) What property is used to order the elements in the Periodic Table?

_______________________________________________________ [1]

(ii) Explain why transition metals are classified as d-block elements.

_______________________________________________________ [1]

(b) A number of distinct trends can be seen in the 3rd period from sodium to argon.

(i) Describe the change in melting point across this period.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(ii) Describe and explain the change in atomic radius across this period.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(iii) On the axes below sketch the change in the 1st ionisation energy across the 3rd period.

[3]11 12 13 14 15atomic number

1st ionisationenergy

(kJ mol–1)

16 17 18







15 (a) Diamond and graphite have giant covalent structures.

(i) Explain what is meant by the term covalent.

_______________________________________________________ [1]

(ii) Describe the structures of diamond and graphite.

Diamond: ________________________________________________

_________________________________________________________

_______________________________________________________ [2]

Graphite: ________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(iii) Explain why graphite conducts electricity.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(iv) Explain why diamond is exceptionally hard.

_________________________________________________________

_______________________________________________________ [1]







(b) Carbon dioxide, CO2, and beryllium chloride, BeCl2, are both covalent compounds.

(i) Draw dot and cross diagrams for carbon dioxide and for beryllium chloride.

carbon dioxide beryllium chloride

[2]

(ii) State the octet rule and explain why beryllium chloride does not obey it.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]







16 Rock salt, impure sodium chloride, is found in large underground deposits at Kilroot.

(a) (i) Describe how you would carry out chemical tests used to show that solid rock salt contains sodium chloride.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [5]


(ii) Draw dot and cross diagrams to show how sodium chloride is formed from sodium and chlorine atoms.

[4]

(b) Chlorine is manufactured by the electrolysis of concentrated sodium chloride solution.

(i) Explain why sodium chloride solution conducts electricity but solid sodium chloride does not.

_________________________________________________________

_______________________________________________________ [1]







(ii) Household bleach is manufactured by reacting chlorine with sodium hydroxide solution.

Cl2 + 2NaOH → NaCl + NaOCl + H2O

Using oxidation numbers, explain why this reaction is described as disproportionation.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [3]

(iii) Describe what you would observe when chlorine is bubbled through a solution of potassium bromide.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(iv) Write an ionic equation for the reaction of chlorine with potassium bromide.

_______________________________________________________ [1]

(c) Concentrated sulphuric acid reacts with sodium halides to form the corresponding hydrogen halide.

(i) Write an equation for the reaction of concentrated sulphuric acid with sodium chloride.

_______________________________________________________ [2]

(ii) Give two observations when concentrated sulphuric acid is added to sodium iodide.

_________________________________________________________

_______________________________________________________ [2]







17 The electronic structure of atoms has been interpreted from analysis of emission spectra.

The diagram below shows the emission spectrum of hydrogen in the ultraviolet region.

(a) Draw the electron transition responsible for the line at 122 nm.

n = 4 _____________________________

n = 3 _____________________________

n = 2 _____________________________

n = 1 _____________________________ [2]

(b) Explain what is meant by the convergence limit.

___________________________________________________________ [1]

90 95 100 105 110

122 nm

Wavelength (nm)

Convergence limit91.1 nm

115 120 125







(c) The convergence limit can be used to calculate the ionisation energy for hydrogen.

(i) Write an equation, including state symbols, for the ionisation of atomic hydrogen.

_______________________________________________________ [2]

(ii) Use the information below to calculate the frequency of the line at the convergence limit.

(speed of light = 3 × 108 m s–1, 1 nm = 1 × 10–9 m)

speed of light = frequency × wavelength

_________________________________________________________

_______________________________________________________ [1]

(iii) Use this frequency value to calculate the energy required to ionise one mole of hydrogen atoms.

Energy required to ionise one hydrogen atom

_________________________________________________________

Energy required to ionise one mole of hydrogen atoms in kJ mol–1

________________________________________________kJ mol–1 [2]








2009

Chemistry


Module 1: Basic Concepts in Physicaland Inorganic Chemistry

[AC111]


MARKSCHEME

New

Specifi

catio

n







Section A

1 A

2 B

3 C

4 A

5 D

6 A

7 D

8 C

9 B

10 B


Section A 20







Section B

11 (a)Molecule Attractive force

Ammonia Hydrogen bonds

Hydrogen chloride Dipole-Dipole

Methane Van der Waals [1] each [3]

(b) More/longer/fixed Hydrogen bonds (between the water molecules) [1] give ice a more open structure (and so a lower density) [1] [2]

(c)

Diagram [1] Repulsion between electron pairs [1] Mention of four pairs or comment on lone pair [1] [3] 8

12 (a) Number ofprotons

Number ofelectrons

Number of neutrons

Neon-20 10 10 10

Neon-21 10 10 11

Neon-22 10 10 12 [–1] for each mistake [2]

(20 x 90.92) + (21 x 0.26) + (22 x 8.82) (b) ––––––––––––––––––––––––––––––– 100

= 20.18 [–1] for each mistake [2]

(c) Carbon-12 isotope [1]

(d) 2p

2s

1s

Subshell labels [1] electronic arrangement [1] [–1] for each mistake

N

H HH







(e)

s orbital p orbital [1] each [2] 9

13 (a) CaCO3 + 2HCl → CaCl2 + H2O + CO2 [2], [–1] for each mistake [2]

(b) A solution of known concentration [1]

(c) (2 x 20)/1000 = 0.04 mole (0.1 x 18.6)/1000 = 0.00186 mole 0.0186 0.04 – 0.0186 = 0.0214 (0.0214/2) x 100 = 1.07(g) (1.07/1.12) x 100 = 95.5% [–1] for each mistake [6]

(d) phenolphthalein/methyl orange [1] from colourless [1] to pink/from red [1] to yellow [1] [3] 12

14 (a) (i) Atomic number [1]

(ii) Their (outer) electrons are in the d-subshell [1]

(b) (i) Melting point increases to silicon [1] then decreases (to argon) [1] [2]

(ii) Atomic radius decreases across the period [1] Shielding remains the same but nuclear charge increases [1] [2]

(iii)

General rise across the period [1] Fall between Groups 2 and 3 [1] Fall between Groups 5 and 6 [1] [3] 9

12 14 16 18







O C

xx

xx

xx

xx

O

Cl Be Cl

XOXO

XOXO

x

x

x

15 (a) (i) Pair(s) of electrons shared between (two) atoms [1]

(ii) Diamond: Carbon atoms joined to 4 others [1] tetrahedrally [1] Graphite: hexagonal rings of carbon atoms [1] in layers [1] [4]

(iii) Free electrons [1] are able to move [1] around the layers [2]

(iv) Strong (covalent) bonds [1] throughout the giant (tetrahedral) structure [1]

(b) (i)

[–1] for each mistake [1]

[–1] for each mistake [1]

(ii) Octet rule: eight electrons in the outer shell (when bonded) [1] Be (has less than 8) in beryllium chloride/has only 4 electrons in its outer shell [1] [2] 12

16 (a) (i) Sodium: nichrome wire [1]/(conc HCl) blue flame [1]/yellow [1] Chloride: (make a solution) silver nitrate [1] white precipitate [1] or dissolve in HNO3 [1] (solution) [5]


(ii) [4]

(b) (i) In the solution the ions are free to move, (they cannot move in the solid) [1]

(ii) Chlorine atoms are both oxidised (0 to +1) [1] and reduced (0 to –1) [1] this is disproportionation [1] [3]

(iii) Colourless solution [1] turns yellow [1] [2]

(iv) Cl2 + 2Br – → 2Cl – + Br2 [1]

Na+ •Cl Na+ Cl –• • • •

• • • ••• ••×•







(c) (i) H2SO4 + NaCl → NaHSO4 + HCl (–1 if Na2SO4) [2]

(ii) Steamy fumes/purple vapour/yellow solid/fizzing/heat evolved/ grey-black solid/ rotten egg smell/choking gas (SO2) any two, [1] each [2] 22

17 (a) Level 2 to level 1 [1] indication of downwards [1] [2]

(b) An electron leaves the atom/energy levels come together [1]

(c) (i) H(g) → H+(g) + e–

(–1 for each mistake) [2]

(ii) 3 × 108 = 91.1 × 10–9 × frequency 3.29 × 1015 [1]

(iii) E = hf = (6.63 × 10–34) × (3.29 × 1015) = 2.18 × 10–18 [1] (2.18 × 10–18) × (6.02 × 1023) = 1314 kJ mol–1 [1] [2] 8

Section B 80

Total 100







TIME

1 hour 30 minutes.

New

Specifi

catio

n


The total mark for this paper is 100.Quality of written communication will be assessed in question 13(c).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.


Write your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all fifteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all five questions in Section B. Write your answers in the spaces provided in this question paper.

ADVANCED SUBSIDIARY (AS)

General Certificate of Education

January 2010

Chemistry


Basic Concepts in Physicaland Inorganic Chemistry

[AC111]








Section A


Select the correct response in each case and mark its code letter by connecting the dots

as illustrated on the answer sheet.

1 2.65 g of anhydrous sodium carbonate, Na2CO3, was dissolved in water and the solution made up to 250 cm3 in a volumetric flask. The concentration of the solution was

A 0.025 mol dm–3

B 0.050 mol dm–3

C 0.100 mol dm–3

D 0.200 mol dm–3

2 In which one of the following molecules does the central atom obey the octet rule?

A BeCl2 B BF3 C CF4 D SF6

3 Which one of the following statements about iodine is not correct?

A It has a molecular covalent structure. B It contains non-polar molecules. C It exists as a grey-black shiny solid. D It is more soluble in water than hexane.

4 Elements Q and R have ground state electron structures 1s22s22p63s2 and 1s22s22p5 respectively. Q and R combine to produce a compound with the formula

A QR B QR2 C Q2R D Q2R5







5 Which one of the following molecules is polar?

A BF3 B CF4 C OF2 D F2

6 4.88 g of hydrated barium chloride, BaCl2.xH2O, was heated to a constant mass of 4.16 g. What is the value of x?

A 1 B 2 C 3 D 4

7 Which one of the following represents the emission spectrum of atomic hydrogen in the ultraviolet region?

C D

A B

wavelength

frequency

wavelength

frequency







8 When burned in a plentiful supply of oxygen, propane (C3H8) produces carbon dioxide and water.

C3H8 + 5O2 → 3CO2 + 4H2O

What is the number of molecules of carbon dioxide produced when 4.4 g of propane are burned?

A 6.02 × 1022

B 1.81 × 1023

C 6.02 × 1023

D 1.81 × 1024

9 A compound produces a lilac colour in a flame test. When chlorine is bubbled into an aqueous solution of the compound, the solution changes from colourless to yellow-orange. The compound is

A potassium bromide B potassium iodide C sodium bromide D sodium iodide

10 Iron(III) oxide can be reduced by carbon to form iron.

2Fe2O3 + 3C → 4Fe + 3CO2

What is the maximum mass of iron which can be produced when 3.20 kg of iron(III) oxide is heated with 0.72 kg of carbon?

A 1.12 kg B 2.24 kg C 3.36 kg D 4.48 kg







Section B

Answer all five questions in this section.

11 There are five isotopes of germanium.

(a) Atoms of the 74Ge isotope contain 32 protons, 32 electrons and 42 neutrons. Complete the following table which shows the properties of each of these particles.

Particle Relative mass Relative charge

Proton

Electron

Neutron

[3]

(b) State, in terms of protons and neutrons, the meanings of the following terms:

Mass number ________________________________________________

___________________________________________________________ [1]

Atomic number _______________________________________________

___________________________________________________________ [1]

Isotopes _____________________________________________________

___________________________________________________________ [1]







(c) The terms atomic number and mass number can be used to deduce the numbers of protons, neutrons and electrons in an atom or ion.

(i) An atom has 15 protons and 20 neutrons fewer than there are in the 70Ge isotope. Deduce the symbol and mass number of this atom.

Symbol _______________ Mass Number _______________ [2]

(ii) Complete the table for the ions of the elements X, Y and Z. The letters X, Y and Z are not the symbols of the elements.

IonAtomic

Number

Mass

Number

Number of

Neutrons

Electronic

Structure

X2+ 20 1s22s22p63s23p6

Y– 18 1s22s22p63s23p6

Z2– 16 1s22s22p6

[3]

(d) The mass spectrum of germanium is used to calculate its relative atomic mass.

(i) Define the term relative atomic mass.

_________________________________________________________

_______________________________________________________ [2]







(ii) The table below gives the percentage abundance of each isotope in the mass spectrum of germanium.

Relative Isotopic Mass

70 72 73 74 76

% Abundance 20.55 27.37 7.67 36.74 7.67

Use this information to calculate the relative atomic mass of germanium to one decimal place.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]







12 Phosphorus and nitrogen are in Group V of the Periodic Table. Nitrogen forms a hydride called ammonia and the hydride of phosphorus is called phosphine, PH3.

(a) (i) Draw a dot and cross diagram to show the bonding in phosphine.

[2]

(ii) Draw and name the shape of a phosphine molecule.

_______________________________________________________ [2]

(iii) Explain why a phosphine molecule has the shape you have drawn.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]







(iv) Suggest a value for the bond angles in phosphine. Explain your answer.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(b) Phosphine reacts with hydrogen ions to form phosphonium ions, PH4

+. Ammonia similarly forms ammonium ions, NH4+.

(i) Name the type of bond formed when phosphine reacts with a hydrogen ion.

_______________________________________________________ [1]

(ii) Explain how this bond is formed.

_________________________________________________________

_______________________________________________________ [2]

(c) Suggest, in terms of intermolecular forces, why ammonia has a higher boiling point than phosphine.

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [2]







13 Analysis of a vinegar solution was carried out using the following procedure:

Transfer 25.0 cm3 of undiluted vinegar into a 250 cm3 volumetric flask

25.0 cm3

and add a few drops of indicator to each flask. Titrate each solution with 0.1 mol dm–3 sodium hydroxide until an end point is reached.

A student obtained the following results:

Initial burette

reading (cm3)

Final burette

reading (cm3)Titre (cm3)

Rough 0.0 21.7 21.7

1st accurate 21.7 43.1

2nd accurate 0.0 21.3

(a) (i) Name a suitable indicator for this titration.

_______________________________________________________ [1]

(ii) State the colour change which would be obtained at the end point.

from ________________________ to ________________________ [2]

(b) (i) Write the equation for the reaction between vinegar (ethanoic acid) and sodium hydroxide.

_______________________________________________________ [2]

(ii) Complete the results table and calculate the average titre.

_______________________________________________________ [2]

(iii) Use the average titre to calculate the number of moles of sodium hydroxide used in the titration.

_______________________________________________________ [1]







(iv) Calculate the concentration of ethanoic acid in the diluted vinegar.

______________________________________________ mol dm–3 [1]

(v) Calculate the concentration of ethanoic acid in the undiluted vinegar.

______________________________________________ mol dm–3 [1]

(c) Describe, giving practical details, how you would prepare the solution of diluted vinegar and then transfer 25.0 cm3 to a conical flask.

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [4]








14 Ionisation energies provide evidence for the existence of shells and subshells in atoms.

(a) State the meaning of the term first ionisation energy of an element.

_____________________________________________________________

___________________________________________________________ [2]

(b) There is a general increase in the first ionisation energies across Period 3. The graph below shows the variation of the first ionisation energies of some of the elements in Period 3.

(i) Use crosses to mark the relative positions of the first ionisation energies for the elements Mg, Al and P. Complete the graph by joining the crosses. [2]

(ii) Explain the general increase in first ionisation energy across the period.

_________________________________________________________

_______________________________________________________ [2]

(iii) Using s, p and d notation give the ground state electronic configuration of a magnesium atom.

_______________________________________________________ [1]


Elements of Period 3

Firs

t ion

isa

tion

ener

gy (

kJ m

ol–1

)







(iv) Explain the position of the first ionisation energy of magnesium relative to that of aluminium in your graph.

_________________________________________________________

_______________________________________________________ [2]

(c) The first four ionisation energies of aluminium are 578, 1817, 2745 and 11 578 kJ mol–1.

(i) Label the subshells in the following diagram for an aluminium atom and use the electrons-in-boxes notation to show how the electrons are arranged in the Al2+ ion.

[2]

(ii) Write the equation, including state symbols, for the fourth ionisation energy of aluminium.

_______________________________________________________ [2]

(iii) Explain why the third ionisation energy of aluminium is much smaller than the fourth ionisation energy.

_________________________________________________________

_______________________________________________________ [2]







15 The elements in Group VII are all reactive non-metals.

(a) There is a trend in the electronegativity of the elements in Group VII.

(i) Define the term electronegativity.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(ii) State and explain the trend in the electronegativity of the elements down Group VII from fluorine to iodine.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [3]

(b) When concentrated sulphuric acid is added to solid sodium bromide, the acid reacts with bromide ions to form sulphur dioxide and bromine.

(i) State the change in the oxidation number of sulphur in this reaction.

_______________________________________________________ [2]

(ii) Write the half-equation to show how bromine is formed from bromide ions.

_______________________________________________________ [1]

(iii) Complete the half-equation to show how sulphur dioxide is formed from sulphuric acid.

H2SO4 + H+ → SO2 + H2O

[2]







(iv) Write the overall ionic equation for the reaction of bromide ions with sulphuric acid.

_______________________________________________________ [1]

(v) State one observation in the above reaction.

_______________________________________________________ [1]

(vi) State the role of bromide ions in this reaction.

_______________________________________________________ [1]

(vii) Why are sulphur dioxide and chlorine not formed when concentrated sulphuric acid is added to solid sodium chloride?

_________________________________________________________

_______________________________________________________ [1]

(viii) Write the equation for the reaction of solid sodium chloride with concentrated sulphuric acid.

_______________________________________________________ [2]

(c) When sodium bromide is dissolved in water, the presence of bromide ions can be established by using aqueous silver nitrate followed by concentrated ammonia solution.

(i) What is observed when aqueous silver nitrate is added to sodium bromide solution?

_______________________________________________________ [2]

(ii) Write the ionic equation, including state symbols, for the reaction.

_______________________________________________________ [2]

(iii) What is observed when an excess of concentrated ammonia solution is added?

_______________________________________________________ [1]








January 2010

MARKSCHEME

New

Specifi

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n

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[AC111]








Section A

1 C

2 C

3 D

4 B

5 C

6 B

7 B

8 B

9 A

10 B


Section A 20







Section B

11 (a)

1 mark per particle [3]

(b) Number of protons + Number of neutrons [1]

Number of protons [1]

Atoms which have the same number of protons but different numbers of neutrons [1]

(c) (i) CI [1] 35 [1] [2]

(ii)

1 mark per particle [3]

(d) (i) The average mass of an atom of an element [1] relative to (one twelfth of) the mass of an atom of carbon-12. [1] [2]

(ii) [(70 × 20.55) + (72 × 27.37) + (73 × 7.67) + (74 × 36.74) + (76 × 7.67)] ÷ 100

72.7 [2] 15

Particle Relative mass Relative charge

Proton 1 + 1

Electron – 1

Neutron 1 0

11840

Ion Atomic Number

Mass Number

Number of Neutrons

Electronic Structure

X2+ 20 40 20 1s22s22p63s23p6

Y– 17 35 18 1s22s22p63s23p6

Z2– 8 16 8 1s22s22p6







12 (a) (i)

[2]

(ii)

[1]

Pyramidal [1]

(iii) In the outer shell of the central P atom there are (3) bonding pairs and (1) lone pair of electrons [1] which repel [1] [2]

(iv) approximately 107° [1] Lone pair-bonding pair repulsion is greater than

bonding pair-bonding pair repulsion [1] [2]

(b) (i) coordinate (dative covalent) bond [1]

(ii) the phosphorus atom shares its lone pair with the H+ ion [2]

(c) Hydrogen bonding in ammonia [1] is stronger than dipole-dipole (or van der Waals’) in phosphine [1] [2] 13

H P H

H

H HH

P







13 (a) (i) Phenolphthalein [1]

(ii) colourless [1] to pink [1] [2]

(b) (i) CH3COOH + NaOH → CH3COONa + H2O [2]

(ii) 21.4 and 21.3 into table [1] Average titre = 21.35 cm3 [1]

(iii) 0.1 × (21.35/1000) = 2.135 × 10−3 [1]

(iv) (2.135 × 10−3) ÷ (0.025) = 0.0854 [1]

(v) 0.854 [1]

(c) Rinsing the pipette (2-3 times) with undiluted vinegar Use of pipette filler Meniscus on mark Mixing/inverting solution in volumetric flask Rinsing the pipette (2-3 times) with diluted vinegar Max = 4 [4]








14 (a) Energy required to convert one mole of gaseous atoms to gaseous ions with single positive charges. [2]

(b) (i)

Point for Mg above Na but not above that for Si point for Al below that for Mg but not below Na [1]

Point for P above that for S and below that for Cl [1]

(ii) increasing nuclear charge [1] Shielding approximately constant/atomic radius decreases [1]

(iii) 1s22s22p63s2 [1]

(iv) (stability of) filled subshell for Mg atom [1] Stability not present in aluminium/3p1/outer electron is in higher

energy subshell [1] [2]

(c) (i)

Labels [1] Configuration [1] [2]

(ii) Al3+ (g) → Al4+ (g) + e− [2]

(iii) Outer electron in Al2+ is further from nucleus/more shielded/ stability of sub shell full Al3+(higher energy level) 2 from 3 [2] 15


×

×

×

×

×

×

×

×

3p

3s

2p

2s

1s







15 (a) (i) The ability/power of an atom to attract bonding electrons towards itself in a covalent bond [2]

(ii) decreases [1] Shared pair is further from nucleus [1] more shielded [1] [3]

(b) (i) +6 [1] to +4 [1] [2]

(ii) 2Br− → Br2 + 2e− [2]

(iii) H2SO4 + 2H+ + 2e− → SO2 + 2H2O [1]

(iv) H2SO4 + 2H+ + 2Br− → SO2 + 2H2O + Br2 [1]

(v) red-brown gas [1]

(vi) reducing agent [1]

(vii) chloride ions are not strong enough reducing agents to reduce concentrated sulphuric acid [1]

(viii) NaCl + H2SO4 → NaHSO4 + HCl [2]

(c) (i) cream [1] precipitate [1] [2]

(ii) Ag+ (aq) + Br− (aq) → AgBr (s) Equation [1] State symbols [1] [2]

(iii) cream precipitate dissolves [1] 21

Section B 80

Total 100







New

Specifi

catio

n

TIME1 hour 30 minutes.

INFORMATION FOR CANDIDATESThe total mark for this paper is 100.Quality of written communication will be assessed in question 13(d).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.

INSTRUCTIONS TO CANDIDATESWrite your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all fifteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all five questions in Section B. Write your answers in the spaces provided in this question paper.


2010

Chemistry



[AC111]

MONDAY 7 JUNE, MORNING







Section A



1 Which one of the following represents the ground state electronic configurationof a nitrogen atom?

2p

2s

1s

2 When burned in a plentiful supply of oxygen, methane produces carbon dioxide and water.

CH4 + 2O2 → CO2 + 2H2O

What is the number of molecules of oxygen required for the complete combustion of1.6 g of methane?

A 6.0 × 1022

B 1.2 × 1023

C 6.0 × 1023

D 1.2 × 1024

3 In which one of the following molecules are the bond angles closest to 107°?

A BF3 B CH4 C H2O D NH3

A B C D







4 Which one of the following gives the correct flame colour for the named compound?

A barium chloride red B copper(II) chloride blue-green C potassium chloride yellow/orange D sodium chloride lilac

5 An element was analysed using a mass spectrometer. The spectrum showed that there were four isotopes. The relative isotopic masses and relative abundances are given below.

Relative isotopic mass Relative abundance

50 2

52 35

53 4

54 1

The relative atomic mass of this element is

A 52.00. B 52.05. C 52.25. D 52.50.

6 Which one of the following is produced when concentrated sulphuric acid reacts with solid sodium chloride?

A chlorine B hydrogen chloride C hydrogen sulphide D sulphur dioxide







7 Titanium is extracted in a two-stage process. The first stage involves the conversion of titanium(IV) oxide to titanium(IV) chloride. In the second stage, the titanium(IV) chloride is reduced using magnesium.

TiO2 + C + 2Cl2 → TiCl4 + CO2

TiCl4 + 2Mg → Ti + 2MgCl2

What is the maximum mass of titanium which can be obtained when 8.0 kg of titanium(IV) oxide is converted to titanium(IV) chloride and then reduced using 7.2 kg of magnesium?

A 2.4 kg B 4.8 kg C 9.6 kg D 14.4 kg

8 Which one of the following electron transitions is responsible for the lowest frequency line in the visible region of the emission spectrum of atomic hydrogen?

n = 4

n = 3

n = 2

n = 1

A

BD

C







9 Which one of the letters represents the first ionisation energy of an alkali metal?

10 The ionisation energy of hydrogen is 1312 kJ mol–1. Use this value to calculate the frequency at convergence in the hydrogen emission spectrum.

A 2.179 × 10–21 Hz B 1.312 × 106 Hz C 3.287 × 1015 Hz D 1.979 × 1036 Hz

B

A

C

D

atomic number

firstionisationenergy/kJ mol –1







Section B


11 Aluminium is the most abundant metal in the Earth’s crust. It is used in electrical cables and is present in high strength alloys.

(a) All atoms of aluminium have a mass number of 27. How many neutrons are present in the nucleus of these atoms?

_____________________________________________________ [1]

(b) (i) Write the equation, including state symbols, which represents the first ionisation energy of aluminium.

_________________________________________________ [2]

(ii) Explain why the first ionisation energy of boron has a larger value than the first ionisation energy of aluminium.

___________________________________________________

___________________________________________________

_________________________________________________ [2]

(iii) Explain why the first ionisation energy of magnesium has a larger value than the first ionisation energy of aluminium.

___________________________________________________

___________________________________________________

_________________________________________________ [2]

(iv) Give the ground state electronic configuration of the Al4+ ion.

_________________________________________________ [1]







(v) Sketch a graph to show the successive ionisation energies of aluminium.

[2]

log

(ioni

satio

n en

ergy

)

number of electrons removed1 2 3 4 5 6 7 8 9 10 11 12 13







12 Fluorine is the most reactive non-metallic element. It combines with both metals and non-metals.

(a) (i) Using dot and cross diagrams, explain how strontium atoms combine with fluorine atoms to form strontium fluoride. Show the outer electrons only.

[4]

(ii) Compare the electrical conductivity of solid strontium metal with that of solid strontium fluoride. Explain your answer.

____________________________________________________

___________________________________________________

___________________________________________________

_________________________________________________ [3]







(b) Sulphur and fluorine combine to form a non-polar molecule sulphur hexafluoride, SF6.

(i) Define the term electronegativity.

_________________________________________________ [2]

(ii) Label the diagram below to show the polarity of the S—F bond.

S—F

[1]

(iii) Draw a dot and cross diagram to show the bonding in SF6 using outer shell electrons only.

[2]

(iv) Explain whether the SF6 molecule obeys the octet rule.

___________________________________________________

___________________________________________________

_________________________________________________ [2]







(v) Draw, name and explain the shape of the SF6 molecule.

___________________________________________________

___________________________________________________

_________________________________________________ [4]

(vi) Suggest why SF6 is a non-polar molecule, even though it contains polar bonds.

___________________________________________________

_________________________________________________ [2]

(c) Boron trifluoride can combine with ammonia to form the following molecule.

(i) Name the type of bond formed between the boron and nitrogen atoms.

_________________________________________________ [1]

(ii) Explain how this bond is formed.

___________________________________________________

_________________________________________________ [1]

F

F

F

H

HNB

H







13 Concentrated nitric acid (HNO3) oxidises iodide ions to form iodine. In the reaction the nitric acid is reduced to form nitrogen monoxide (NO).

(a) Reduction and oxidation can be defined in different ways.

(i) Define oxidation in terms of electron transfer.

_________________________________________________ [1]

(ii) Define reduction in terms of changes in oxidation state.

_________________________________________________ [1]

(b) Deduce the oxidation number of nitrogen in

(i) HNO3 _________________________________________ [1]

(ii) NO _________________________________________ [1]

(c) The half-equation for the reduction of concentrated nitric acid is shown below.

HNO3 + 3H+ + 3e– → NO + 2H2O

(i) Write a half-equation for the oxidation of iodide ions to form an iodine molecule.

_________________________________________________ [1]

(ii) Combine the reduction and oxidation half-equations to give the overall ionic equation.

_________________________________________________ [2]







(d) At room temperature and pressure, iodine exists as a grey-black shiny solid. Describe the bonding in, and explain the structure of iodine crystals. Explain the relative solubilities of iodine in water and hexane.

_______________________________________________________

_______________________________________________________

_______________________________________________________

_______________________________________________________

_______________________________________________________

_______________________________________________________

_____________________________________________________ [5]








14 Chlorine is produced by the electrolysis of concentrated sodium chloride solution (brine). It is then used by other industries to produce a variety of useful products.

(a) The reaction between chlorine and cold dilute sodium hydroxide is used in the manufacture of bleach.


_________________________________________________ [2]

(ii) This reaction is described as disproportionation. Explain the meaning of this term.

___________________________________________________

___________________________________________________

_________________________________________________ [2]

(b) Chlorine reacts with hydrogen to produce hydrogen chloride.


_________________________________________________ [2]

(ii) Suggest why hydrogen chloride has a much lower boiling point than hydrogen fluoride.

___________________________________________________

___________________________________________________

_________________________________________________ [3]

(iii) Explain why hydrogen chloride is more thermally stable than hydrogen iodide.

___________________________________________________

___________________________________________________

_________________________________________________ [2]







(c) The presence of chloride ions in brine can be established by adding an aqueous solution of silver nitrate.

(i) What would be observed in this reaction?

_________________________________________________ [2]

(ii) Write an ionic equation, including state symbols, for the reaction.

_________________________________________________ [2]

(iii) State what is observed when an excess of dilute aqueous ammonia is then added.

_________________________________________________ [1]







15 The degree of hydration in samples of hydrated sodium carbonate (Na2CO3.xH2O) can be determined by different methods.

(a) When 10.04 g of a sample was heated to constant mass, 3.97 g of anhydrous sodium carbonate was obtained.

(i) Explain the term anhydrous.

___________________________________________________

_________________________________________________ [1]

(ii) Calculate the number of moles of anhydrous sodium carbonate obtained.

_________________________________________________ [1]

(iii) Calculate the mass of water present in the sample.

_________________________________________________ [1]

(iv) Calculate the number of moles of water present.

_________________________________________________ [1]

(v) Calculate the value of x in the sample.

_________________________________________________ [1]







(b) The degree of hydration can also be determined by dissolving the sample in water and titrating with a standard solution of hydrochloric acid.

(i) What is a standard solution?

_________________________________________________ [1]

(ii) Give the equation for the reaction between sodium carbonate and hydrochloric acid.

_________________________________________________ [2]

(iii) Name a suitable indicator for this titration.

_________________________________________________ [1]







(c) 3.57 g of a second sample of sodium carbonate was dissolved in water and the resulting solution was made up to 250 cm3 in a volumetric flask. A 25.0 cm3 sample of this solution required 28.5 cm3

of 0.1 mol dm–3 hydrochloric acid to reach the end point.

(i) Give the colour change which would be obtained at the end point, using the indicator given in (b)(iii).

From _____________________ to _____________________ [2]

(ii) Calculate the number of moles of hydrochloric acid used in the titration.

_________________________________________________ [1]

(iii) Calculate the number of moles of sodium carbonate present in 25.0 cm3 of solution.

_________________________________________________ [1]

(iv) Calculate the number of moles of sodium carbonate present in 250 cm3 of solution.

_________________________________________________ [1]

(v) Calculate the mass of sodium carbonate present in the second sample.

_________________________________________________ [1]

(vi) Calculate the mass of water present in the second sample.

_________________________________________________ [1]

(vii) Calculate the number of moles of water present in the second sample.

_________________________________________________ [1]

(viii) Calculate the value of x in the second sample.

_________________________________________________ [1]








2010

MARKSCHEME

New

Specifi

catio

n

Chemistry



[AC111]

MONDAY 7 JUNE, MORNING







1 D

2 B

3 D

4 B

5 B

6 B

7 B

8 B

9 D

10 C


Section A 20







Section B

11 (a) 14 [1]

(b) (i) AI (g) → AI+ (g) + e–

Equation [1] State symbols [1] [2]

(ii) Outer electron for boron is closer to nucleus [1] and is less shielded [1] (than for aluminium) [2]

(iii) stability of filled 3s shell for magnesium [1] aluminium has 3p1 configuration [1] [2]

(iv) 1s22s22p5 [1]

(v)

[2] 10

log (ionisation energy)







12 (a) (i)

Sr atom [1] F atom [1] 1 : 2 ratio [1] correct electron transfer + charges [1]

(ii) solid strontium – good electrical conductor [1] solid strontium fluoride – poor electrical conductor solid strontium – delocalised electrons [1] solid strontium fluoride – ions are not free to move [1]

(b) (i) The ability/power of an atom to attract bonding electrons in a covalent bond [2]

(ii)

[1]

(iii)

[2]

(iv) does not apply to sulphur (12 electrons in outer shell) [1] does apply to fluorine (8 electrons in outer shell)/octet rule – has 8 e– in outer shell [1]

Sr××

Sr 2+ – –

F

F

F

F× ×

d+ d–

S –– F

×F

×

F

×F×

F

×F

×F

S

}







(v)

[1]

Octahedral [1] (six) bonding pairs [1] repel equally [1] [2]

(vi) dipoles [1] cancel [1] [2]

(c) (i) coordinate/dative (covalent) bond [1]

(ii) both shared electrons come from nitrogen/lone pair of electrons on nitrogen shared (donated) [1] 22

13 (a) (i) loss of electrons [1]

(ii) decrease in oxidation state/number [1]

(b) (i) +5 [1]

(ii) +2 [1]

(c) (i) 2I− → I2 + 2e− [1]

(ii) 2HNO3 + 6H+ + 6I− → 2NO + 4H2O + 3I2 [2]

(d) reference to covalent bonding/crystalline molecular covalent structure/diatomic/I2 van der Waals’ attractions between molecules 1 mark for each two bold points mentioned – Max [3] marks more soluble in hexane than water [1] since iodine and hexane are non-polar [1] and water is polar [1] like dissolves like [1] Max [4] marks Max [5] marks


s

F

F

F FF F







14 (a) (i) Cl2 + 2NaOH → NaCI + NaOCI + H2O [2]

(ii) simultaneous reduction and oxidation [1] of the same element/in the same reaction [1] [2]

(b) (i) H2 + CI2 → 2HCI [2]

(ii) HF – hydrogen bonding/polar [1] HCl – polar/less polar than HF [1] Greater energy needed to separate molecules [1]

(iii) H−CI bond stronger than H−I bond [1] does not break as easily when heated [1]

(c) (i) white [1] precipitate [1] [2]

(ii) Ag+ (aq) + CI− (aq) → AgCI (s)

Equation [1] State symbols [1] [2]

(iii) (white) precipitate dissolves [1] 16







15 (a) (i) contains no water (of crystallisation) [1]

(ii) 0.0375 [1]

(iii) 6.07g [1]

(iv) 0.337 [1]

(v) 9 [1]

(b) (i) solution of known concentration [1]

(ii) Na2CO3 + 2HCl → 2NaCl + H2O + CO2 [2]

(iii) methyl orange [1]

(c) (i) yellow/orange [1] red/pink [1] [2]

(ii) 0.00285 [1]

(iii) 0.001425 [1]

(iv) 0.01425 [1]

(v) 1.5105 [1]

(vi) 2.06 [1]

(vii) 0.1144 [1]

(viii) 8 [1] 18

Section B 80

Total 100







TIME

1 hour 30 minutes.


Write your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all sixteen questions. Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all six questions in Section B. Write your answers in the spaces provided in this question paper.


The total mark for this paper is 100.Quality of written communication will be assessed in question 14(d).

In Section A all questions carry equal marks, i.e. two marks for each


January 2011

Chemistry



[AC111]








Section A



1 An atom in which the number of protons is greater than the number of neutrons is

A 2H. B 3He. C 10B. D 39K.

2 Which one of the following is a correct description of electronic transitions in a given series in the atomic emission spectrum of hydrogen?

A They all start from the ground state. B They all end at the ground state. C They all start from one particular energy level. D They all end at one particular energy level.

3 Which one of the following lists the first ionisation energies (in kJ mol–1) of the elements magnesium, aluminium, silicon, phosphorus and sulfur in this order?

A 496 736 577 786 1060 B 577 786 1060 1000 1260 C 736 577 786 1060 1000 D 786 1060 1000 1260 1520







4 The mass spectrum of molecular chlorine, Cl2, is shown below. An additional peak is in the spectrum which should not be present.

Which one of the following peaks should not be present?

70 35

37

m/e

7172

74

A 35 B 71 C 72 D 74

5 A solid melts sharply at 100–101 ºC. It does not conduct electricity even when molten. It dissolves in hydrocarbon solvents. The solid has

A an atomic structure. B a giant covalent structure. C an ionic structure. D a molecular covalent structure.

6 Which one of the following gaseous hydrides most readily decomposes into its elements on contact with a hot glass rod?

A ammonia B hydrogen fluoride C hydrogen iodide D steam







7 Arsine, AsH3, is a molecular hydride of arsenic which is found in Group V of the Periodic Table. Which one of the following is the structure of arsine in the vapour state?

H

Bond angles 120º

As

HA

H

Bond angles greater than 109.5º

B

D

Bond angles less than 109.5º

C

H

Bond angles 109.5º

As

HH

H

As

HH

H

As

HH

8 50 cm3 of 0.20 mol dm–3 sulphuric acid is exactly neutralised by

A 100 cm3 of 0.40 mol dm–3 potassium hydroxide solution. B 25 cm3 of 0.20 mol dm–3 potassium hydroxide solution. C 50 cm3 of 0.20 mol dm–3 potassium hydroxide solution. D 100 cm3 of 0.20 mol dm–3 potassium hydroxide solution.

9 Which one of the following is the number of electrons which have approximately the same mass as that of a proton?

A 20 B 200 C 2000 D 20000

10 Which one of the following oxides is not polar?

A CO B CO2 C H2O D NO







SectionB

Answerallsixquestionsinthissection.

11 Nitrogendioxide,NO2,isoneofthecomponentsofphotochemicalsmog.TheenergyrequiredtodissociatethismoleculeintoNOmoleculesandOatomsis305kJmol–1.

Usethefollowingheadingstocalculatethefrequencyofradiationrequiredtocausethedissociation.

(a) Convert305kJintoJoules

___________________________________________________________ [1]

(b) CalculatethenumberofJoulesrequiredtodissociateonemoleculeofnitrogendioxide.

___________________________________________________________ [1]

(c) UsetheequationE=hftoconvertthevalueinJoulesintoafrequencyandstatetheunits.

_____________________________________________________________

___________________________________________________________ [1]







12 The mineral beryl, Be3Al2Si6O18, is the principal source of beryllium. Although there are minerals richer in beryllium they are scarce and costly.

(a) Calculate the percentage, by mass, of beryllium in beryl.

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [3]

(b) The metal beryllium is obtained either by the electrolysis of a fused mixture of beryllium and potassium chlorides at 350 ºC or by the reduction of beryllium fluoride with magnesium.

(i) Write an equation for the formation of beryllium from beryllium ions.

_______________________________________________________ [1]

(ii) Write an equation for the formation of beryllium from the reduction of beryllium fluoride with magnesium.

_______________________________________________________ [1]

(c) The first element in a Group often has more distinctive properties than the elements in the rest of the Group. This is often as a result of the difference in electronegativities. The electronegativity values of the Group II elements are shown below.

electronegativityvalue

atomic number

Ba

Be

Mg

CaSr

(i) Explain the meaning of the term electronegativity.

_________________________________________________________

_______________________________________________________ [2]







(ii) Using electronegativity suggest why beryllium chloride is a covalent molecule and barium chloride is ionic.

_________________________________________________________

_______________________________________________________ [2]

(iii) State two physical properties which could be used to distinguish these two chlorides.

_________________________________________________________

_______________________________________________________ [2]

(d) Beryllium chloride may be prepared by the action of chlorine or hydrogen chloride on the metal.

(i) Write the equation for the reaction of beryllium with hydrogen chloride.

_______________________________________________________ [1]

(ii) Draw a dot and cross diagram to show the formation of beryllium chloride from beryllium and chlorine atoms. Use only the outer electrons of each atom.

[3]

(iii) State the octet rule.

_________________________________________________________

_______________________________________________________ [2]

(iv) Beryllium chloride can be said to obey the octet rule and also not to obey the octet rule. Explain this contradiction.

_________________________________________________________

_______________________________________________________ [2]







(v) Draw the shape of a beryllium chloride molecule.

[1]

(vi) State the shape of the beryllium chloride molecule.

_______________________________________________________ [1]

(vii) Explain the shape of the beryllium chloride molecule.

_________________________________________________________

_______________________________________________________ [2]







13 Chlorineformsaseriesofoxidessomeofwhicharelistedbelow.

chlorinemonoxide Cl2O

chlorinedioxide ClO2

chlorinehexoxide Cl2O6

chlorineheptoxide Cl2O7

(a) Deducethesystematicnameforchlorineheptoxideusingtheoxidationnumberofchlorine.

___________________________________________________________ [1]

(b) Chlorinedioxidedissolvesinwatertoformasolutionwhicheventuallyformsamixtureofchloricandhydrochloricacids.

6ClO2+3H2O→5HClO3+HCl

Thechlorineatomsinchlorinedioxideundergodisproportionationinthisreaction.

(i) Explainthemeaningofthetermdisproportionation.

_________________________________________________________

_______________________________________________________ [1]

(ii) Calculatetheoxidationnumberofchlorineinthereactantandintheproductsofthisreactionandusethemtoconfirmthatthereactionisadisproportionationreaction.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [3]







(c) Chlorine gas dissolves in water to the extent of 0.8 g in 100 cm3 at atmospheric pressure and 20 ºC.

(i) Calculate the molarity of the chlorine water, Cl2 (aq), produced.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(ii) Name another solvent in which chlorine will readily dissolve.

_______________________________________________________ [1]







14 An experiment was set up to investigate the displacement reactions of the halogens.

Solutions of sodium halides were prepared and reacted with other halogens. The results table is shown below.

sodium iodide (aq)

sodium bromide (aq)

sodium chloride (aq)

iodine solution

X X

bromine solution ✓

chlorine solution

✓ means that a reaction took place X means that no reaction took place

(a) Complete the three remaining places in the table. [2]

(b) (i) Both bromine and iodine solutions are coloured. Describe the observations which would indicate that a reaction took place when aqueous sodium iodide is added to a bromine solution.

_________________________________________________________

_______________________________________________________ [2]

(ii) Write the ionic equation for the reaction between bromine solution and aqueous sodium iodide.

_______________________________________________________ [1]

(c) (i) Describe what is observed when chlorine solution is added to aqueous sodium bromide.

_________________________________________________________

_______________________________________________________ [2]

(ii) Write the equation for the reaction between chlorine solution and aqueous sodium bromide.

_______________________________________________________ [1]







(d) If you had poured solutions of sodium iodide, bromide and chloride into beakers A, B and C and forgotten to label them, describe how, using aqueous silver nitrate and both dilute and concentrated ammonia solutions, you would determine which sodium salt was in which beaker. Each beaker must be tested.

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [6]








15 The structure of ice is shown below. The water molecules are held together by hydrogen bonds which are a type of intermolecular force.

(a) Name two other types of intermolecular force.

___________________________________________________________ [2]

(b) (i) Explain how hydrogen bonding takes place between the water molecules in ice.

_________________________________________________________

_______________________________________________________ [2]

(ii) Explain, using the structure above, why ice is less dense than water.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(c) Although water is capable of forming hydrogen bonds it does not form long chains of “polywater” at room temperature. However, in the liquid state, molecules such as hydrogen fluoride do form very short chains.

Suggest why water does not form chains and liquid hydrogen fluoride does.

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [2]







(d) Ammonia is another substance that can form hydrogen bonds. However, ammonia has a pyramidal structure.

(i) Draw two molecules of ammonia and show the hydrogen bond between the two molecules.

[2]

(ii) Explain why when ammonia reacts with a hydrogen ion it loses the ability to form hydrogen bonds.

_________________________________________________________

_______________________________________________________ [1]

(e) Explain why ammonia is extremely soluble in water.

_____________________________________________________________

___________________________________________________________ [2]







16 Lithium exists in nature as two isotopes, 6Li and 7Li. The composition of a sample of lithium in nature is shown in the table below.

isotope % abundance

lithium 6 7.42

lithium 7 92.58

(a) Draw the structure of a 7Li atom, labelling all the sub-atomic particles.

[3]

(b) State and explain to which of the s, p or d blocks lithium belongs.

_____________________________________________________________

___________________________________________________________ [2]

(c) Calculate the relative atomic mass of lithium to two decimal places.

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [3]







(d) Lithium sulphate is readily soluble in water and crystallises from solution as the hydrate.

(i) Explain what is meant by the term water of crystallisation.

_________________________________________________________

_______________________________________________________ [1]

(ii) Write the formula of anhydrous lithium sulphate.

_______________________________________________________ [1]

(iii) Calculate the formula of hydrated lithium sulphate if 3.76 g of the hydrated lithium salt produces 3.23 g of anhydrous lithium sulphate on heating.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [3]

(e) Lithium sulphate can be used in a flame test. Explain how a flame test could be carried out and state the expected colour of the flame.

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

_____________________________________________________________

___________________________________________________________ [4]








January 2011


assessingBasic Concepts in Physical

and Inorganic Chemistry

[AC111]


MARKSCHEME







Section A

1 B

2 D

3 C

4 B

5 D

6 C

7 C

8 D

9 C

10 B

[2] for each correct answer [20]

Section A

20

20







Section B

11 (a) 305 3 103 5 3.05 3 105 J [1]

(b) 3.05 × 105

6.02 × 1023 5 5.07 3 10–19 J [1]

(c) 0.507 3 10–18 5 6.63 3 10–34 3 f f 5 0.076 3 1016

5 7.6 3 1014 s–1/Hz [1]

12 (a) Be3Al2Si6O18 5 3 3 9 1 2 3 27 1 6 3 28 1 18 3 16 5 27 1 54 1 168 1 288 5 537

% Be 5 27537 3 100 5 5.03% [3]

(b) (i) Be2+ 1 2e– Be [1]

(ii) BeF2 1 Mg Be 1 MgF2 [1]

(c) (i) ability of an atom in a covalent bond to attract (bonding)electrons [2]

(ii) Be and Cl have similar EN values [1] Ba and Cl have (very) different EN values [1] [2]

(iii) melting point, boiling point, “reaction” with water, (electrical) conductivity etc.

2 from list [2]

(d) (i) Be 1 2HCl BeCl2 1 H2 [1]

(ii)

Be Be Clxxx

xxxx Cl

xxx x

xx

xxClxx

xxClxxx

x

xx

xxx

[3]

(iii) 8 electrons around an atom in outer shell [2]

(iv) 8 electrons around Cl [1] 4 electrons around Be [1] [2]

(v) Cl — Be — Cl [1]

(vi) linear/straight [1]

(vii) bond electrons repel [1] to minimise forces [1] [2]

3

23







13 (a) chlorine(VII) oxide [1]

(b) (i) atom raises and lowers its oxidation number during a chemical reaction [1]

(ii) ClO2 14 HClO3 15 HCl –1

[2]

{ 14 15 oxidation 14 –1 reduction

[1] [3]

(c) (i) Cl2 5 2 3 35.5 5 71

0.8 g 5 0.871 5 0.01127 mol

{ 0.113 M [2]

(ii) e.g. hexane [1]

14 (a) [2] (Note that the marking of the colour changes in this question will

be subject to the application of the “colour changes” scheme.)

(b) (i) compare colours with original solutionsit should go darker

colour of iodine is darker than bromine [2]

(ii) Br2 1 2I– 2Br – 1 I2 [1] no state symbols required

(c) (i) colourless solution [1] orange/yellow/brown colour produced [1] [2]

(ii) Cl2 1 2NaBr 2NaCl1 Br2 [1] no state symbols required

(d) iodide: yellow ppt insoluble in (both dilute and conc.) ammonia solution [2] bromide: cream ppt soluble in conc NH3 [2] chloride: white ppt soluble in dil NH3 [2] [6]


8

16







15 (a) van der Waals [1] dipole – dipole [1] [2]

(b) (i) attraction between lone pair on O and a H atom on another water molecule [2]

(ii) distance between water molecules in ice greater than in water [1] open structure [1] longer/expanded H-bonds [1] fixed H-bonds [1]

any 2 [2]

(c) F more electronegative than oxygen/H—F bond more polarised [1] F H bond stronger [1] movement of water molecules breaks H bonds [1] [2]

(d) (i)

H N H H

H

H

H

N| | | | | | | |

[2]

(ii) lone pair removed/forms a dative bond [1]

(e) N•• forms hydrogen bonds with

and N H forms hydrogen bonds with [2]

[1] OH H

13 HO H

[1]







16 (a)

+ + +

electronproton

neutron [3]

(b) s block [1] outer electron in s shell [1]

(c) 7.42 3 6 5 44.52

92.58 3 7 5 648.06

Total 5 692.58

4 100 5 6.9258

5 6.93 [3]

(d) (i) water chemically bonded in salt [1]

(ii) Li2SO4 [1]

(iii) Li2SO4 5 2 3 7 1 32 1 64 5 110

moles 5 3.23110 5 0.029

H2O 5 2 1 16 5 18

moles 5 0.5318 5 0.029

{ Li2SO4 • H2O [3]

(e) conc hydrochloric acid blue Bunsen flame nichrome (platinum) wire crimson flame [4]

Section B

Total

17

80

100

17

80

100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in question 15(f).

In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand

question.A Periodic Table of Elements (including some data) is provided.


2011

Chemistry



[AC112]

WEDNESDAY 15 JUNE, AFTERNOON







SectionA

Foreachofthefollowingquestionsonlyoneoftheletteredresponses(A–D)iscorrect.

Selectthecorrectresponseineachcaseandmarkitscodeletterbyconnectingthedotsasillustratedontheanswersheet.

1 Anelementwhichformsanionsmallerthanitsatomis

A chlorine. B potassium. C oxygen. D sulfur.

2 Acompoundwhichdoesnotconsistofindividualmoleculesis

A berylliumchloride. B calciumchloride. C hydrogenchloride. D phosphorustrichloride.

3 Whichoneofthefollowingelementscontainsthesamenumberofelectronsasanionofmagnesium,Mg21?

A calcium B fluorine C neon D sodium

4 Whichoneofthefollowingisthemassofcalciumcarbonatewhichwillexactlyneutralise500cm3of0.1Mhydrochloricacid?

A 1.25g B 2.50g C 12.50g D 25.0g







5 Which one of the following is the shape of the ammonia molecule?

HA B

C D

N HH

N

H

N

HH

H107°HH

N

H120°HH

6 If 30 g of water were completely converted into hydrogen and oxygen which one of the following would be the total mass of gases produced?

A 10 g B 30 g C 45 g D 90 g

7 Which one of the following can not be used to obtain hydrogen chloride in the laboratory?

A burning hydrogen in chlorine B heating concentrated hydrochloric acid C the reaction of chlorine with methane D bubbling chlorine through hexane at room temperature

8 The element europium reacts with hydrogen to form europium hydride. Atoms of europium have their outer electrons in levels 5 and 6 i.e. 5s2 5p6 6s2. Which one of the following formulae resembles europium hydride?

A AsH3 B CH4 C CaH2 D SnH4







9 Which one of the following atoms in the ground state contains no unpaired electrons?

A argon B fluorine C potassium D sulfur

10 Which one of the following species contains a coordinate bond?

A NH3

B NH4+

C NH2–

D NH2–







Section B


11 The molecule drawn below is that of a giant covalent structure. All the atoms are of the same element.

(a) Name the substance.

___________________________________________________________ [1]

(b) Explain whether the substance is hard or soft.

_____________________________________________________________

___________________________________________________________ [2]

(c) Explain whether the substance conducts electricity or not.

_____________________________________________________________

___________________________________________________________ [2]







12 Barium chloride crystallises from water to form a hydrate with the formula BaCl2.xH2O. On heating, the hydrate loses water to form anhydrous barium chloride. A solution of barium chloride is colourless.

(a) Barium chloride solution reacts with aqueous silver nitrate to form silver chloride.

(i) Write the equation for the reaction.

_______________________________________________________ [1]

(ii) Write the ionic equation for the reaction.

_______________________________________________________ [1]

(iii) Describe what is observed during the reaction.

_______________________________________________________ [1]

(b) 3.05 g of BaCl2.xH2O were dissolved in water to make 250 cm3 of solution in a graduated flask. 20 cm3 of this solution were titrated with 0.1 M silver nitrate solution. It was found that 20.0 cm3 were required.

(i) How many moles of silver ions were added during the titration?

_______________________________________________________ [1]

(ii) How many moles of chloride ions were there in 20 cm3 of the barium chloride solution?

_______________________________________________________ [1]

(iii) How many moles of anhydrous barium chloride were there in 250 cm3 of the solution?

_______________________________________________________ [1]

(iv) What is the relative formula mass of the hydrated barium chloride?

_______________________________________________________ [1]







(v) What is the relative formula mass of anhydrous barium chloride?

_______________________________________________________ [1]

(vi) Calculate the value of x in BaCl2.xH2O

_______________________________________________________ [1]

(c) Write the equation for the action of heat on the hydrated barium chloride.

___________________________________________________________ [1]







13 Astatine, the last element of the halogen group, was synthesised in 1940. Since then it has been stated that it is the rarest naturally occurring element on Earth with an estimated 30 g of astatine existing at any one time. It was named from the Greek word for “unstable”.

(a) The longest living isotope of astatine is astatine-210, 210At. However, half of this isotope disappears after about 8 hours.

(i) Define the meaning of the term isotope.

_________________________________________________________

_______________________________________________________ [2]

(ii) Name and calculate the numbers of the individual sub-atomic particles in one atom of astatine-210.

_________________________________________________________

_______________________________________________________ [3]

(b) Using Avogadro’s number calculate the number of astatine atoms that exist in 30 g of astatine. Assume that all of the atoms are of astatine-210.

_____________________________________________________________

___________________________________________________________ [2]

(c) Astatine was predicted to exist by Mendeleev in his original Periodic Table and was given the name eka-iodine.

Complete the table below by predicting some of the properties of astatine.

Property Result for Astatine

formula of an astatine molecule

physical state at room temperature

colour of astatine at room temperature

colour of astatine vapour/gas

solubility in water (yes or no)

solubility in hexane (yes or no)

[6]







(d) Write the equation for the reaction of iodine with sodium astatide.

___________________________________________________________ [1]

(e) Astatine was first made by bombarding bismuth with alpha particles. Explain how you would show, using mass spectrometry, that astatine had actually been formed.

_____________________________________________________________

___________________________________________________________ [1]







14 The diagram below shows a common method of preparing hydrogen chloride gas in the laboratory.

sodiumchloride

concentratedsulfuric acid

(a) Write the equation for the reaction of sodium chloride with concentrated sulfuric acid.

___________________________________________________________ [2]

(b) The hydrogen chloride is collected by “downward delivery” in which air, a mixture of oxygen and nitrogen, is displaced upwards.

(i) Calculate the relative molecular masses of oxygen, nitrogen and hydrogen chloride.

oxygen ___________________________________________________

nitrogen __________________________________________________

hydrogen chloride _______________________________________ [2]

(ii) Use the values of the calculated relative molecular masses to explain why hydrogen chloride is collected by downward delivery.

_________________________________________________________

_______________________________________________________ [2]







(c) Sometimes dry hydrogen chloride gas is required.

(i) Explain why sodium hydroxide would be inappropriate to use as a drying agent.

______________________________________________________________________________________________________________________________________________________________________________________________________________________________ [1]

(ii) Suggest the name of a substance which could be used.

______________________________________________________________________________________________________________________________________________________________________________________________________________________________ [1]

(d) Explain whether concentrated sulfuric acid could be reacted with sodium bromide in the preparation of hydrogen bromide using this method.

___________________________________________________________ [1]

(e) Which one of the following acids, hydrogen chloride, hydrogen bromide and hydrogen iodide is the strongest?

___________________________________________________________ [1]

(f) How could you prove that a gas jar you believed contained hydrogen chloride actually contained the gas.

_____________________________________________________________

___________________________________________________________ [2]







15 Some properties of the Group I elements from sodium to caesium are shown in the table below.

metal ionic radius/nm first ionisation energy/kJ mol–1

melting point/K

sodium 0.102 496 371

potassium 0.138 419 336

rubidium 0.149 403 312

caesium 0.170 376 302

(a) Explain why all of the Group I elements are described as being s-block elements.

___________________________________________________________ [1]

(b) Explain why the ionic radius increases down the group.

_____________________________________________________________

___________________________________________________________ [1]

(c) The first ionisation energy of the Group I elements may be determined using spectroscopic methods.

(i) If the frequency of the radiation needed to remove the outermost electron from a sodium atom is 1.25 1015 s–1 calculate the first ionisation energy of sodium in kJ per mole.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [3]

(ii) Write the equation, using state symbols, for the first ionisation energy of sodium.

_______________________________________________________ [2]







(iii) Give two reasons to explain why potassium has a lower first ionisation energy than sodium.

_________________________________________________________

_________________________________________________________

_______________________________________________________ [2]

(iv) Why is the second ionisation energy of all the Group I metals very much higher than the first?

_________________________________________________________

_______________________________________________________ [1]

(d) All of the Group I metals exhibit metallic bonding.

(i) Using a labelled diagram, explain what is meant by the term metallic bonding.

_________________________________________________________

_______________________________________________________ [3]

(ii) Suggest why the melting point of the metals decreases from sodium to caesium.

_________________________________________________________

_______________________________________________________ [1]

(iii) Using the concept of metallic bonding suggest why calcium should be a better electrical conductor than potassium.

_________________________________________________________

_______________________________________________________ [1]







(e) All of the Group I metals react with halogens to form ionic metal halides. Using outer electrons only draw diagrams to explain the formation of caesium chloride from caesium atoms and chlorine atoms.

[3]

(f) All of the Group I elements produce characteristic flame colours in a Bunsen burner flame which can be used to identify them.

(i) Describe how you would carry out a flame test.

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_________________________________________________________

_______________________________________________________ [4]


(ii) How would you distinguish between sodium chloride and potassium chloride using a flame test?

_________________________________________________________

_______________________________________________________ [1]







16 Carbon dioxide is the most frequently found oxide of carbon in nature. It is a colourless gas with a faint taste and smell. The structure of the molecule can be readily deduced by the application of the octet rule. Even though carbon and oxygen have different electronegativities the molecule does not have a permanent dipole.

(a) Explain the term octet rule.

_____________________________________________________________

___________________________________________________________ [2]

(b) Explain the term electronegativity.

_____________________________________________________________

___________________________________________________________ [2]

(c) (i) Using outer electrons only draw the dot and cross structure of carbon dioxide.

[2]

(ii) Draw and name the shape of a carbon dioxide molecule.

_______________________________________________________ [2]

(iii) Explain why carbon dioxide has the shape you have drawn.

_________________________________________________________

_______________________________________________________ [2]







(d) Carbon and oxygen have different electronegativities and form a polar bond.

Explain why a carbon dioxide molecule does not have a permanent dipole.

_____________________________________________________________

___________________________________________________________ [1]

(e) Although carbon dioxide does not have a dipole it is very soluble in water. Using intermolecular forces explain this extreme solubility.

_____________________________________________________________

___________________________________________________________ [2]








2011


assessing


[AC112]


MARKSCHEME







Section A

1 B

2 B

3 C

4 B

5 C

6 B

7 D

8 C

9 A

10 B


Section A

20

20







Section B

11 (a) diamond [1]

(b) hard [1] strong bonds [1] [2]

(c) no [1] electrons (in bonds) cannot move [1] [2]

12 (a) (i) BaCl2 1 2AgNO3 2AgCl 1 Ba(NO3)2 [1]

(ii) Ag1 1 Cl2 AgCl [1]

(iii) White precipitate [1]

(b) (i) 20 3 1023 3 1021 5 2 3 1023 mol [1]

(ii) 2 3 1023 mol [1]

(iii) 21 3 2 3 1023 3 20

250 5 1.25 3 1022 [1]

(iv) 0.0125 ; 3.05 g

1 mol ; 0.01253.05 5 244 [1]

(v) BaCl2 5 137 1 71 5 208 [1]

(vi) 244 2 208 5 36 36 5 H2O x 5 2 [1]

(c) BaCl2.xH2O BaCl2 1 xH2O [1] or BaCl2.2H2O BaCl2 1 2H2O

5

10







13 (a) (i) same number of protons [1] different number of neutrons [1] [2]

(ii) 85 protons 85 electrons 125 neutrons [21] for each wrong part [3]

(b) 30 g 5 21030 5 0.1429 mol

6.023 3 1023 3 0.1429 5 0.86 3 1023 5 8.6 3 1022 [2]

(c) At2 [1] solid [1] grey-black/black [1] purple/violet/dark violet [1] no [1] yes [1] [6]

(d) I2 1 2NaAt 2NaI 1 At2 [1]

(e) new peak at 210 or round about [1]

14 (a) NaCl 1 H2SO4 NaHSO4 1 HCl use of Na2SO4 5 [1] [2]

(b) (i) O2 5 32 N2 5 28 HCl 5 36.5 wrong value is [21] [2]

(ii) HCl is heavier than N2 or O2 [1] hence HCl sinks [1] [2]

(c) (i) NaOH reacts with HCl/acid 1 base [1]

(ii) conc H2SO4/anhydrous CuSO4 etc. [1] [2]

(d) no – HBr reacts with H2SO4 [1]

(e) hydrogen iodide [1]

(f) use AgNO3(aq) or conc NH3(aq) [1] white ppt/white smoke [1] [2]

15

12







15 (a) outer electrons are s electrons [1]

(b) increased number of shells [1]

(c) (i) E 5 hf 5 6.63 3 10234 3 1.25 3 1015

5 8.29 3 10219 J For one mole 5 6.023 3 1023 3 8.29 3 10219 J 5 49.9 3 104J 5 499 [3]

(ii) Na(g) Na1(g) 1 e2 [2]

(iii) outer electrons further away from the nucleus [1] shielded by increased shells of electrons [1] [2]

(iv) removal of second electron is from a full shell [1]

(d) (i)

[2] electrons are delocalised/can move/

(electrostatic) attraction between e2 and metal ion [1] [3]

(ii) forces of attraction decreases charge density decreases [1]

(iii) Ca has two outer electrons [1]

(e) .Cs 1 Cl

x xx xx x

x Cs1 1 Cl

x xx xx x . x

2 [3]

(f) (i) nichrome/platinum wire [1] blue flame (of Bunsen) [1] conc. hydrochloric acid [1] place compound on wire/put in blue flame [1] [4]


(ii) potassium lilac or sodium yellow/orange [1]

e2 e2 e2 e2

O1 O1 O1 O1e2 e2 e2 e2

O1 O1 O1 O1

25







16 (a) 8 electrons [1] in outer shell [1] [2]

(b) attraction of electrons by an atom [1] in a covalent bond [1] [2]

[2] (c) (i) O OCx xx x

(ii) O C O or O C O [1] linear/straight [1] [2]

(iii) electrons in the bonds [1] repel as much as possible [1] [2]

(d) the dipoles “cancel” out [1]

(e) attraction between d1 on C and d2 on O in H2O or d2 on O and d1 on H in H2O [2]

Section B

Total

13

80

100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in question 11.

In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.


January 2012

Chemistry



[AC112]

FRIDAY 13 JANUARY, AFTERNOON







Section A



1 Which one of the following bonds is the most polar?

A B-F B N-F C C-I D O-I

2 Which one of the following can not form hydrogen bonds?

A H2O B H3O

C NH3 D NH4

3 Which one of the following is the name of the species shown below?

is a protonis a neutronis an electron

��

��

�

A beryllium atom B beryllium ion C lithium atom D lithium ion







4 Which one of the following solids consists of molecular covalent crystals?

A Diamond B Graphite C Ice D Quartz

5 When excess chlorine is bubbled into hot concentrated alkali which one of the following lists the main products of the reaction?

A Cl, ClO, H2O B Cl, ClO3

, H2O C Cl, ClO4

, H2O D ClO, ClO3

, H2O

6 The elements X and Y are in Groups VI and VII respectively of the Periodic Table.

Which one of the following shows the formula and the bond type of the compound that they form?

A XY2, covalent B XY2, ionic C X2Y, covalent D X2Y, ionic

7 Which one of the following orbitals is occupied by an electron with the energy level n 5 2?

A A dumb-bell shaped orbital B A spherically shaped orbital C An s or d orbital D An s or p orbital







8

A crystalline solid melts sharply at 95 °C. It does not conduct electricity in the solid and liquid states. It dissolves in hexane.

Which one of the following is the structure of the crystal?

A giant molecular B ionic C metallic D molecular covalent

9 The diagram below shows a liquid escaping from a burette and passing a charged glass rod.

+++

Which one of the following liquids will be attracted to the glass rod?

A CCl4 B CHCl3 C CS2 D C5H12

10 The species Ar, K and Ca2 have the same number of electrons. Starting with the smallest, which one of the following is the order in which their radii increase?

A Ar Ca2 K

B Ar K Ca2

C Ca2 K Ar D K Ar Ca2







Section B


11 The geometry of covalent inorganic hydrides may be predicted using the electron structures of the molecules.

Draw the shapes of the following hydrides using their outer electron structures. Explain these shapes giving the values of the angles between bonds.

hydrogen chloride HClmethane CH4hydrogen sulfide H2Sammonia NH3

[6]








12 Solutions of acidified iodide ions are very easily oxidised to produce iodine molecules.

Even oxygen, from the air, will oxidise iodide ions to liberate iodine. The following half-equations represent the formation of iodine molecules

and the conversion of oxygen to water.

Equation 1 2I I2 2e

Equation 2 O2 4H 4e 2H2O

(a) (i) Using electron transfer explain which equation represents an oxidation reaction.

[1]

(ii) Using electron transfer explain which equation represents a reduction reaction.

[1]

(b) Write the equation for the reaction of acidified iodide ions with oxygen.

[2]

(c) The following solutions were added to a solution of acidified iodide ions in separate test tubes.

chlorine, iron(III) ions, ammonia, sodium hydroxide, sodium chloride.

Which two of these solutions would react with acidified iodide ions to produce iodine?

[2]







(d) The reaction of oxidising agents with potassium iodide can be used to prepare iodine in the laboratory. When heated with manganese dioxide and concentrated sulfuric acid, potassium iodide liberates iodine.

2KI MnO2 3H2SO4 2KHSO4 MnSO4 2H2O I2

(i) Using oxidation numbers, explain this redox reaction.

[3]

(ii) What observation would confirm that iodine had been produced?

[1]







13 Phosphorus, P, reacts with bromine at room temperature to form phosphorus tribromide, PBr3, which is a liquid with boiling point 173 °C. It reacts with water immediately forming hydrogen bromide and phosphoric(III) acid, H3PO3. The reaction is used as a method of preparing hydrogen bromide in the laboratory.

(a) Write the equation for the reaction of bromine with phosphorus.

[1]

(b) Calculate the maximum mass of phosphorus tribromide which can be formed when 6.2 g of phosphorus, which is an excess, reacts with 8.0 cm3 of bromine, Br2. The density of liquid bromine is 3.1 g cm3.

mass of bromine, Br2, in grams

[1]

moles of bromine, Br2

[1]

moles of phosphorus, P, in 6.2 g

[1]

moles of bromine, Br2 reacting

[1]

moles of phosphorus tribromide formed

[1]

mass of phosphorus tribromide formed

[1]

(c) Write the equation for the reaction of phosphorus tribromide with water.

[1]







(d) The apparatus shown below was used to prepare hydrogen bromide in the laboratory. Bromine is slowly added to a paste of phosphorus and water. Phosphorus tribromide is first formed and is immediately decomposed by the water present. The gases produced are passed through a U-tube containing glass beads coated in phosphorus.

bromine

cardboarddisc

P andwater

moist Pand glassbeads

(i) Suggest why the bromine is not added all at once.

[1]

(ii) An excess of water is not used in the experiment. What is the property of hydrogen bromide which is the reason for not using an excess?

[1]

(iii) Why is the hydrogen bromide collected as shown and not with the delivery tube pointing upwards?

[1]

(e) Hydrogen bromide is a colourless gas but produces fumes in moist air. Why does it fume in moist air?

[1]







(f) When hydrogen bromide is heated in a loosely corked test tube it produces a very pale red–brown colour after heating for several minutes.

(i) Name the substance responsible for the red–brown colour.

[1]

(ii) Explain what would be observed if a test tube of hydrogen iodide was heated.

[1]

(iii) Explain what would be observed if a test tube of hydrogen chloride was heated.

[1]

(iv) What do these observations indicate about the relative thermal stability of the hydrogen halides?

[1]







14 Only 0.08% of the Earth’s crust consists of carbon yet this element is an essential part of living organisms. It occurs naturally as the isotopes carbon-12 and carbon-13 although there is a radioactive isotope carbon-14. Carbon occurs in nature as two structures known as diamond and graphite.

(a) Naturally occurring carbon contains 98.89% of carbon-12 and 1.11% carbon-13. Calculate the relative atomic mass of carbon to three decimal places.

[3]

(b) Carbon-14 is not used in the calculation of the relative atomic mass because virtually none of it exists. It decomposes when a neutron in its nucleus changes into an electron and a proton forming a new element.

(i) What are the numbers of electrons, protons and neutrons in the new element?

[2]

(ii) Name the element produced when carbon-14 decomposes.

[1]

(c) Mass spectrometry uses carbon-12 as the international standard.

(i) What is the purpose of mass spectrometry?

[2]

(ii) Explain the meaning of the term carbon-12 standard.

[2]

(d) Explain why carbon-12 and carbon-14 are isotopes.

[2]







(e) Carbon may be produced in the laboratory in many ways. One is to heat cane sugar, C12H22O11, with concentrated sulfuric acid. Steam and carbon are produced together with diluted sulfuric acid.

(i) Write the equation for the reaction. Do not include sulfuric acid in the equation.

[1]

(ii) Explain the meaning of the terms hydrated and water of crystallisation.

[2]

(iii) Explain whether the cane sugar is hydrated.

[1]

(f) Diamond is oxidised when it burns in oxygen at about 700 °C.

(i) Name the product formed from the complete oxidation of diamond.

[1]

(ii) Name the product formed from the incomplete oxidation of diamond.

[1]

(iii) Explain whether graphite will form the same products when it is burned.

[2]







(g) Draw dot and cross diagrams, using outer electrons only, to show the formation of a carbon dioxide molecule from a carbon atom and an oxygen molecule.

[3]







15 Anhydrous copper(II) chloride, CuCl2, may be prepared by heating copper in chlorine gas. When prepared by dissolving copper(II) oxide in hydrochloric acid, copper(II) chloride crystallises with two molecules of water of crystallisation.

(a) Write the equation for the reaction of copper with chlorine.

[1]

(b) Write the equation for the reaction of copper(II) oxide with hydrochloric acid.

[1]

(c) Write the formula for hydrated copper(II) chloride.

[1]

(d) The purity of the copper(II) oxide may be determined by the process of back titration. Explain, without calculations, how this process would be carried out.

[4]







(e) The presence of copper in copper(II) chloride can be shown using a flame test.

(i) The diagram below shows the equipment needed for the test. Identify the acid W, the metal wire X, the colour Y of the flame before the test and the colour Z during the test.

WX

YZ

W [1]

X [1]

Y [1]

Z [1]

(ii) State two reasons for using W.

[2]

(iii) Explain the origin of the flame colour produced by copper(II) chloride.

[3]







(f) When copper(II) chloride dissolves in water its ions are surrounded by water molecules. The polar water molecules surround both anion and cation.

(i) State the formulae of the ions present in copper(II) chloride.

[2]

(ii) Draw a diagram showing two water molecules around the anion.

[1]

(iii) Draw a diagram showing two water molecules around the cation.

[1]

(iv) Explain how you would confirm the presence of chloride ions in the solution.

[3]








January 2012




[AC111]

FRIDAY 13 JANUARY, AFTERNOON

MARKSCHEME







Section A

1 A

2 D

3 B

4 C

5 B

6 A

7 D

8 D

9 B

10 C [2] for each correct answer [20] 20

Section A 20







Section B

11 Depending on the response of candidates it is likely that two marking points will be needed for each mark awarded.

shapes H–Cl C S N angles 180 109 105 107 dot and cross •

×H Cl •H

HC

HH

×•

× ×•× ×

•× ×

HS× H

× •× × ×

HN

HH×•×• ×•

diagram

apply the following to each compound lone pair v lone pair > lone pair v bond pair > bond pair v bond pair

the electron pairs repel to be as far apart as possible [6]

4 marking points per compound, i.e. shape, electron structure, angle, explanation, i.e. 16 marking points – count number of errors. Apply following: Even number of errors ÷ 2, subtract this from 6 (Odd number of errors –1) ÷ 2, subtract this from 6


12 (a) (i) 2I– → I2 + 2e oxidation because electrons are lost [1]

(ii) O2 → 2H2O reduction because electrons are gained [1]

(b) 4I– + 4H+ + O2 → 2I2 + 2H2O (electrons left in [–1]) [2]

(c) chlorine [1] iron(III) ions [1] [2]

(d) (i) I– = –1 MnO2 = +4 I2 = 0 MnSO4 = +2 [2]

iodide/iodine is oxidised and manganese is reduced [1] [3]

(ii) violet/purple vapour or grey/black solid at top of test tube [1] 10

(accept no angle)







13 (a) 2P + 3Br2 → 2PBr3 or P4 + 6Br2 → 4PBr3 [1]

(b) 8.0 3.1 = 24.8 g [1]

24.8/160 = 0.155 mol [1]

6.2/31 = 0.2 mol [1]

0.155 mol [1]

0.103 mol [1]

PBr3 = 31 + 3 80 = 271 0.103 271 = 27.9 g [1]

(c) PBr3 + 3H2O → 3HBr + H3PO3 [1]

(d) (i) reaction could be too vigorous [1]

(ii) hydrogen bromide is soluble (in water) [1]

(iii) hydrogen bromide is heavier (than air) [1]

(e) dissolves (in water vapour) to form hydrobromic acid [1]

(f) (i) bromine [1] (ii) violet/purple colour [1]

(iii) nothing observed/stays the same/remains colourless [1]

(iv) HCl > HBr > HI (mark is dependent on given observations) [1] 16

14 (a) 98.89 12 = 1186.68 1.11 13 = 14.43 = 1201.11 = 12.011 [3]

(b) (i) 7 electrons 7 protons 7 neutrons [2]

(ii) nitrogen [1]

(c) (i) to determine RAM and isotopic abundance/RMM [2]

(ii) atomic masses or RAM/mol mass/RMM are measured relative to C = 12.000 [2]

(d) same atomic number but different mass numbers [2]

(e) (i) C12H22O11 → 12C + 11H2O [1]

(ii) hydrated: contains water of crystallisation/water present [1] water of crystallisation: water chemically bonded [1] [2]

(iii) not hydrated, water is formed/no water in the sugar [1]







(f) (i) carbon dioxide [1]

(ii) carbon monoxide [1]

(iii) yes [1] it is also carbon [1] [2]

(g) C• •••

+ × ××

× O ××

× ××× O ×

× O× ×

O×× •×•×C•×•×

× ××× [3] 23

15 (a) Cu + Cl2 → CuCl2 [1]

(b) CuO + 2HCl → CuCl2 + H2O [1]

(c) CuCl2.2H2O [1]

(d) Weigh the CuO, add (known) excess (hydrochloric) acid (to CuO) [1] titrate excess hydrochloric acid [1] with (standard) alkali/sodium hydroxide [1] named indicator, e.g. phenolphthalein/methyl orange [1] [4]

(e) (i) W: concentrated hydrochloric acid [1] X: nichrome/platinum [1] Y: blue [1] Z: green-blue [1]

(ii) clean the wire [1] make the solid stick to the wire/dissolve the solid [1] [2]

(iii) electrons (in the energy levels) raised to higher levels [1] fall back down [1] to give out light [1] [3]

(f) (i) Cu2+; Cl– [2]

(ii)

........

Cl–H

O

H

O

H

H

.... ....

. . . . . .Cl–H

OH

OHH

[1]

(iii) HO

HCu2+

HO

H [1]

(iv) add silver nitrate (solution) [1] (add dilute nitric acid) white [1] precipitate/solid [1] [3] 23

Section B 80

Total 100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in question 14(b)(i).

In Section A all questions carry equal marks, i.e. two marks for each


2012

Chemistry



[AC112]








Section A



1 Part of the mass spectrum for an element is shown below:

Which one of the following is the relative atomic mass of the element?

A 20.0 B 20.2 C 21.0 D 22.8

2 Which one of the following metal compounds will produce a lilac flame colour?

A barium nitrate B calcium chloride C lithium chloride D potassium sulfate







3 The graph below shows how the second ionisation energy of elements varies across a period.

Which one of the elements is an alkali metal?

4 Which one of the following is the oxidation number of nitrogen in the nitrate ion, NO3?

A 21 B 23 C 15 D 17

5 Which one of the following molecules is the most polar?

A BF3

B CO2

C F2

D NH3

6 Which one of the following m/z values will not appear when a sample of chlorine gas is injected into a mass spectrometer?

A 35.0 B 35.5 C 37.0 D 74.0







7 Which one of the following molecules contains the smallest bond angle?

A BeCl2 B BF3

C CH4

D SF6

8 5.30 g of anhydrous sodium carbonate was dissolved in water and made up to 250 cm3 in a volumetric flask. Which one of the following is the concentration of sodium ions in mol dm23?

A 0.05 B 0.10 C 0.20 D 0.40

9 Which block in the Periodic Table contains silver?

A d block B f block C p block D s block

10 Which one of the following is involved in metallic bonding?

A electron delocalisation B electron transitions C gaining electrons to form ions D sharing electron pairs







Section B

Answer all five questions in the spaces provided.

11 The elements magnesium and chlorine are characterised by their atomic numbers. Chlorine has two isotopes each with a different mass number.

(a) Define each of the following in terms of protons, neutrons and electrons.

(i) atomic number

[1]

(ii) mass number

[1]

(iii) isotopes

[1]

(b) Magnesium reacts with chlorine to form magnesium chloride.

(i) How many protons, neutrons and electrons are present in each of the following ions?

IonNumbers of

protons neutrons electrons

24Mg21

35Cl2

[2]

(ii) Use the boxes below to complete the electronic configuration of the ions:

�

�







(iii) Use a dot and cross diagram to show, using outer electrons only, how magnesium atoms react with chlorine atoms to form magnesium chloride.

[4]

(c) Magnesium forms ions with a double positive charge.

(i) Define the term second ionisation energy.

[2]

(ii) Write an equation, including state symbols, which represents the second ionisation energy of magnesium.

[2]

(iii) Give reasons why the third ionisation energy of magnesium is much larger than the second.

[3]

(d) The Group II chloride, SrCl2, produces a characteristic red colour in a Bunsen flame. Explain, using energy levels, why this colour is observed.

[3]







12 Avogadro’s number has the value 6.02 3 1023.

(a) Define the term Avogadro’s number.

[2]

(b) X is an oxide of nitrogen.

(i) 2.30 g of X contains 3.01 3 1022 molecules of X. Calculate the molar mass of X.

[2]

(ii) Deduce the formula of X.

[1]

(c) Dinitrogen tetroxide (N2O4) reacts with water to form nitric acid, (HNO3) and nitrogen(II) oxide (NO). Write an equation for the reaction.

[1]







(d) Dilute nitric acid reacts with magnesium:

Mg 1 2HNO3 → Mg(NO3)2 1 H2

(i) Calculate the volume, in cm3, of 2.0 mol dm23 nitric acid required to react with 6.0 g of magnesium.

Number of moles of magnesium

[1]

Number of moles of nitric acid

[1]

Volume of nitric acid (in cm3)

[1]

(ii) Calculate the mass of magnesium nitrate produced.

Number of moles of magnesium nitrate produced

[1]

Mass of magnesium nitrate produced

[1]







13 Calcium carbonate is present in eggshells. The percentage of calcium carbonate may be determined by a back titration method. The eggshells are crushed, weighed and then treated with excess dilute hydrochloric acid.

CaCO3 1 2HCl → CaCl2 1 H2O 1 CO2

The unreacted acid is then titrated with standard sodium hydroxide solution.

HCl 1 NaOH → NaCl 1 H2O

(a) (i) Explain the term standard solution.

[1]

(ii) Name a suitable indicator for the titration and state the colour change occurring at the end point.

indicator [1]

from

to [2]

(b) A student weighed out 10.0 g of the crushed eggshells and added 100.0 cm3 of 2.0 mol dm23 hydrochloric acid. The resultant solution was transferred to a 250 cm3 volumetric flask and made up to the mark with distilled water. 25.0 cm3 portions of the solution were titrated with 0.10 mol dm23 sodium hydroxide solution. The average titre was found to be 18.0 cm3.

(i) Calculate the number of moles of sodium hydroxide used in the titration.

[1]

(ii) Calculate the number of moles of hydrochloric acid present in the 25.0 cm3 portion.

[1]

(iii) Calculate the number of moles of hydrochloric acid present in the 250 cm3 volumetric flask.

[1]







(iv) Calculate the total number of moles of hydrochloric acid added to the crushed eggshells.

[1]

(v) Calculate the number of moles of hydrochloric acid which reacted with the calcium carbonate in the crushed eggshells.

[1]

(vi) Calculate the number of moles of calcium carbonate in the crushed eggshells.

[1]

(vii) Calculate the mass of calcium carbonate in the crushed eggshells.

[1]

(viii) Calculate the percentage, by mass, of calcium carbonate in the crushed eggshells.

[1]







14 The halogens are reactive non-metals which often react by gaining electrons to form halide ions.

(a) Complete the table to show the colours and physical states of chlorine, bromine and iodine at room temperature and pressure.

Halogen Colour Physical State

Chlorine

Bromine

Iodine

[3]

(b) Solutions of silver nitrate and ammonia can be used to test for the presence of aqueous halide ions.

(i) Describe how you would use these reagents to distinguish between solutions of sodium chloride, sodium bromide and sodium iodide. State the expected result for each solution.

[6]


(ii) Give an ionic equation, including state symbols, for the reaction of aqueous sodium iodide with silver nitrate solution.

[2]







(c) Solid samples of sodium chloride, sodium bromide and sodium iodide can be distinguished using concentrated sulfuric acid.

(i) Write an equation for the reaction of sodium chloride with concentrated sulfuric acid.

[2]

(ii) Balance the following half-equation for the reduction of concentrated sulfuric acid to form hydrogen sulfide:

H2SO4 1 H1 1 → H2S 1 H2O

[2]

(iii) Combine the reduction half-equation in (c)(ii) with the following oxidation half-equation to produce a balanced redox equation.

2I2 → I2 1 2e2

[2]

(iv) Give one observation which indicates the formation of hydrogen sulfide.

[1]

(v) Name two other reduction products which are formed when concentrated sulfuric acid is added to sodium iodide.

[2]

(vi) Suggest why iodide ions are stronger reducing agents than chloride ions.

[2]







(d) (i) Write the equation for the reaction of chlorine with hot concentrated sodium hydroxide solution.

[2]

(ii) Name the type of redox reaction taking place.

[1]







15 The bonding and shape of a water molecule determines the properties of water.

(a) Draw a dot and cross diagram to show the bonding in a water molecule.

[2]

(b) (i) What is the bond angle in a water molecule?

[1]

(ii) State the shape of a water molecule and explain why it adopts this shape.

[3]

(iii) Why is the bond angle of water different to the bond angle in methane?

[1]

(c) Why does water have a higher boiling point than hydrogen sulfide?

[2]








2012




[AC112]


MARKSCHEME







Section A

1 B

2 D

3 B

4 C

5 D

6 B

7 D

8 D

9 A

10 A


Section A

20

20







Section B

11 (a) (i) number of protons [1] (ii) number of protons + number of neutrons [1]

(iii) atoms with the same number of protons but a different number of neutrons [1]

(b) (i) 12 12 10 17 18 18 error [–1] [2]

(ii)1s 2s 2p 3s 3p

24Mg2+ [1]

35CI– [1]

(iii)

Mg••

•Cl••

••

••Cl

_

••

••Mg 2+ ••

Cl _

••••

•••Cl••

•• ×•

×•

×× [4] [–1] each error

(c) (i) The energy required to convert one mole of gaseous ions with a single positive charge into (gaseous) ions with a double positive charge [2]

(ii) Mg+(g) → Mg2+(g) + e– [2]

(iii) electron closer to nucleus [1] less shielded [1] full shell [1] [3]

(d) electrons promoted from ground state/lower energy level to excited state/ higher energy level [1] as the electron drops back down [1] energy given out as (red) light [1] [3] 21

12 (a) number of atoms [1] present in 12.000 g of carbon-12 [1] [2]

(b) (i) moles of X = 0.05 [1] molar mass = 46 g [1] [2]

(ii) NO2 [1]

(c) 3N2O4 + 2H2O → 4HNO3 + 2NO [1]

(d) (i) moles of Mg = 0.25 moles of nitric acid = 0.50 volume of nitric acid = 250 cm3 each error [–1] [3]

(ii) moles of magnesium nitrate formed 0.25 molar mass = 148 mass (g) = 0.25 × 148 = 37g each error [–1] [2] 11







13 (a) (i) solution of known concentration [1]

(ii) e.g. phenolphthalein/methyl orange/colourless [1] to pink [1] or red/pink [1] to yellow/orange [1] [3]

(b) (i) 0.0018/1.8 × 10–3 [1]

(ii) 0.0018/1.8 × 10–3 [1]

(iii) 0.018/1.8 × 10–2 [1]

(iv) 0.2 [1]

(v) 0.2 – 0.018 = 0.182 [1]

(vi) 0.091 [1]

(vii) 9.1 g [1]

(viii) 91% [1] 12

14 (a) green/green-yellow/yellow-green gas red-brown liquid grey-black/black solid [3] [–1] for each error

(b) (i) white precipitate for solution of NaCl [1] dissolves in dilute ammonia [1] cream precipitate for solution of NaBr [1] dissolves in concentrated ammonia [1] yellow precipitate for solution of NaI [1] does not dissolve in concentrated ammonia [1] [6]


(ii) Ag+(aq) + I–(aq) → AgI(s) [2]

(c) (i) NaCl + H2SO4 → NaHSO4 + HCl [2]

(ii) H2SO4 + 8 H+ + 8e– → H2S + 4 H2O unbalanced [–1] [2]

(iii) H2SO4 + 8 H+ + 8I– → 4 I2 + H2S + 4 H2O unbalanced [–1] [2]

(iv) smell of rotten eggs [1]

(v) sulfur dioxide, sulfur [2]

(vi) the outer electrons in the iodide ions are further from the nucleus/ more shielded [1] iodide ions lose electrons more easily (than chloride ions) [1] [2]

(d) (i) 3Cl2 + 6NaOH → 5NaCl + NaClO3 + 3H2O [2]

(ii) disproportionation [1] 27







15 (a)

•H H

•O

× ×××

× × [2]

(b) (i) 104–105° [1]

(ii) V-shaped/angular/bent [1] repulsion [1] between lone (electron) pairs and bond pairs [1] [3]

(iii) no lone pairs in methane/only bond pairs in methane [1] [1]

(c) water molecules held together by hydrogen bonding [1] which is stronger than the intermolecular/polar forces between hydrogen sulfide molecules [1] [2] 9

Section B 80

Total 100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in question 11(b).

In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements (including some data) is provided.


January 2013

8197

Chemistry



[AC112]








Section A



1 Which one of the following is the electronic configuration for the Fe21 ion?

A 1s22s22p63s23p63d44s2

B 1s22s22p63s23p63d54s1

C 1s22s22p63s23p63d64s0

D 1s22s22p63s23p63d64s2

2 Which one of the following lists the first four ionisation energies of a Group II element?

A 584, 1823, 2751, 11584

B 744, 1457, 7739, 10547

C 793, 1583, 3238, 4362

D 1018, 1909, 2918, 4963

3 Which one of the following is the mass of zinc chloride produced when 8.1 g of zinc oxide, ZnO, is added to 150.0 cm3 of 1.0 mol dm23 hydrochloric acid?

A 10.2 g

B 13.6 g

C 20.4 g

D 40.8 g

4 Which one of the following substances has coordinate bonding in its structure?

A Ammonia

B Ammonium chloride

C Carbon dioxide

D Water







5 Which one of the following, in the liquid state, has van der Waals’ forces and permanent dipole attractions but not hydrogen bonds between the molecules?

A CH4

B CO

C H2O

D O2

6 Chlorine does not undergo disproportionation when reacted with

A cold dilute sodium hydroxide solution.

B hot concentrated sodium hydroxide solution.

C potassium bromide solution.

D water.

7 Which one of the following solids will react with concentrated sulfuric acid to give hydrogen sulfide?

A Calcium bromide

B Magnesium iodide

C Potassium chloride

D Sodium fluoride

8 For which one of the following titrations would phenolphthalein be a suitable indicator?

A Ethanoic acid and sodium carbonate

B Ethanoic acid and sodium hydroxide

C Hydrochloric acid and aqueous ammonia

D Hydrochloric acid and sodium carbonate







9 Which one of the following increases as Group VII is descended?

A Atomic radius

B Electronegativity

C First ionisation energy

D Reactivity

10 Which one of the following is the bond angle in a water molecule?

A 104.5°

B 107.0°

C 109.5°

D 112.0°







Section B


11 The graph below shows the successive ionisation energies of an element when all its electrons are removed.

1 2 3 4 5 6 7

log ionisation energy

number of electrons removed

XY

(a) Name the element that gives rise to this graph.

[1]

(b) (i) Explain why the ionisation energies increase in section X.

[2]

(ii) Explain the large difference in ionisation energies in section Y.

[2]








12 Caesium is a very reactive metal that has sky blue lines in the visible region of its emission spectrum (from the Latin caesius meaning “sky blue”).

(a) (i) How are the lines in the emission spectrum of caesium formed?

[3]

(ii) Why do the lines in the emission spectrum of caesium converge?

[1]

(b) Caesium has one of the lowest first ionisation energies of all the elements in the Periodic Table.

(i) Write an equation, including state symbols, for the first ionisation of caesium.

[2]

(ii) The frequency at the convergence limit of caesium is 9.41 3 1014 Hz. Calculate the first ionisation energy of caesium in kJ mol21.

[3]

(iii) Give two reasons why the first ionisation energy of caesium is low.

[2]







(c) Caesium is so reactive that it will react with gold to form caesium auride, CsAu. The gold can be obtained from caesium auride by reacting it with water. Caesium hydroxide and hydrogen are the other products. Write an equation for the reaction of caesium auride with water.

[2]







13 Fluorine is the most electronegative element in the Periodic Table. Although the element is extremely reactive, fluoride ions can be safely added to water supplies and toothpastes.

(a) What is the meaning of the term electronegativity?

[2]

(b) Fluorine reacts with boron to form boron trifluoride.

(i) Draw a dot and cross diagram, using outer shell electrons only, to show the bonding in boron trifluoride.

[2]

(ii) State and explain the shape of a boron trifluoride molecule.

Shape:

Explanation:

[3]

(iii) State the octet rule.

[2]

(iv) Explain whether or not the elements in boron trifluoride obey the octet rule.

[2]







(c) Sodium fluoride is added to toothpaste to strengthen tooth enamel.

(i) The data on a 50 g tube of toothpaste states that it contains “1450 ppm fluoride”; ppm means “parts per million” i.e. there would be 1450 g of fluoride ions in 106 (1,000,000) g of toothpaste. Use the following headings to work out the concentration of sodium fluoride in the toothpaste. The density of the toothpaste is 1.6 g cm–3.

Mass of fluoride ions in the 50 g tube

Number of moles of fluoride ions in the 50 g tube

Number of moles of sodium fluoride in the 50 g tube

Volume of toothpaste in the 50 g tube

Concentration of sodium fluoride in the toothpaste with units

[6]

(ii) Sodium fluoride is also added to some public water supplies. Why might some people be opposed to this?

[1]

(iii) Giving practical details, describe how you could prove that a sample of solid sodium fluoride contains sodium ions.

[4]







14 Halide ions can be displaced from aqueous solutions of their salts using a more reactive halogen.

(a) A student bubbled excess chlorine into sodium bromide solution and the following reaction took place.

Cl2(g) 1 2NaBr(aq) 2NaCl(aq) 1 Br2(aq)

(i) What change would be observed in the solution?

[2]

(ii) With reference to oxidation numbers explain why this is a redox reaction.

[3]

(iii) Describe how you could prove that there were no bromide ions remaining.

[4]

(b) Hydrogen halides are gases which are very soluble in water forming acidic solutions.

(i) State and suggest an explanation for the strength of hydrofluoric acid relative to the other hydrogen halides.

[2]

(ii) Silicon dioxide is used to make glass. Hydrofluoric acid has to be stored in plastic containers as it reacts with the “silicon dioxide” in glass to produce silicon tetrafluoride and water. Write the equation for this reaction.

[2]







15 Magnesium is an s-block metal which exists as the three isotopes, 24Mg, 25Mg and 26Mg in a ratio of 8:1:1. It reacts with oxygen to form magnesium oxide which has a wide variety of uses including cement manufacture and heartburn medication.

(a) (i) Why is magnesium in the s-block of the Periodic Table?

[1]

(ii) Explain the meaning of the term isotope.

[2]

(iii) Calculate the relative atomic mass of magnesium to one decimal place.

[3]

(b) Magnesium oxide can be formed by the combustion of magnesium metal in oxygen.

(i) Draw a dot and cross diagram, using outer shell electrons only, to show how magnesium oxide is formed from a magnesium atom and an oxygen atom.

[3]

(ii) What type of bonding exists in magnesium oxide?

[1]

(iii) State two physical properties of magnesium oxide.

[2]







(c) The amount of magnesium oxide in heartburn tablets can be determined by adding a known excess of hydrochloric acid to the tablets.

2HCl(aq) 1 MgO(s) MgCl2(aq) 1 H2O(l)

The amount of unreacted hydrochloric acid is determined by titrating it against sodium hydroxide.

NaOH(aq) 1 HCl(aq) NaCl(aq) 1 H2O(l)

(i) What is this method of titration called?

[1]

(ii) A student added 80 cm3 of 2.0 mol dm23 hydrochloric acid to five crushed heartburn tablets which contained magnesium oxide. The unreacted acid required 25 cm3 of 2.0 mol dm23 sodium hydroxide for complete neutralisation. Use the headings below to calculate the mass, in milligrams, of magnesium oxide in each tablet.

Number of moles of hydrochloric acid added to the tablets

Number of moles of unreacted hydrochloric acid

Number of moles of hydrochloric acid which reacted with the magnesium oxide

Number of moles of magnesium oxide present in five tablets

Mass of magnesium oxide per tablet in milligrams

[6]







16 The table below gives some information about three solids, aluminium, ice and diamond.

Aluminium Ice Diamond

Density (g cm23) 2.70 0.92 3.52

Electrical conductivity

High Low Low

Melting point (°C) 660 0 3550

Use your knowledge of bonding and intermolecular forces to answer the following questions.

(a) Why is the density of ice lower than the density of water?

[2]

(b) Explain the high electrical conductivity of aluminium.

[2]

(c) Why does diamond have a high melting point?

[2]








January 2013

Chemistry



[AC112]


MARKSCHEME







Section A

1 C

2 B

3 A

4 B

5 B

6 C

7 B

8 B

9 A

10 A [2] for each correct answer [20] 20

Section A 20







Section B

11 (a) Nitrogen [1]

(b) (i) X: proton to electron ratio increasing or effective nuclear charge increasing [1]; as each electron is removed the remaining electrons are held more tightly [1] [2]

(ii) Y: electron being removed from a full shell [1] closer to nucleus or less/no shielding represents the change in ionisation energy [1] [2]


12 (a) (i) Electrons in high energy level [1] drop back down [1] emit energy/light [1] [3]

(ii) Energy levels converge [1] (b) (i) Cs(g) → Cs+(g) + e– [2]

(ii) E = (6.63 × 10–34)(9.41 × 1014) = 6.239 × 10–19

IE = (6.239 × 10–19)(6.02 × 1023) 1000 = 375.6 kJ mol–1 [3]

(iii) The electron being removed is far from the nucleus [1] and experiences a lot of shielding [1] [2] (c) 2CsAu + 2H2O → 2CsOH + 2Au + H2 [2] 13







13 (a) The extent to which an atom attracts the (bonding) electrons [1] in a covalent bond [1] [2]

(b) (i) B× ××

[2]

(ii) Trigonal planar [1] Bonding electrons [1] repel to get as far apart as possible [1] [3]

(iii) Atoms bond in order to get eight electrons [1] in the outer shell [1] [2]

(iv) Boron does not obey the octet rule in BF3 as it has only six electrons [1] in its outer shell; fluorine does obey [1] [2]

(c) (i) 1450/106 × 50 = 0.0725 0.0725/19 = 3.816 × 10–3

3.816 × 10–3

50/1.6 = 31.25 cm3 or 0.03125 dm3

3.816 × 10–3 ÷ 31.25 × 1000 = 0.122 mol dm–3 or 0.000122 mol cm–3 [6]

(ii) Freedom of choice [1]

(iii) Nichrome wire [1] concentrated hydrochloric acid [1] place sample in blue Bunsen flame [1] orange/yellow [1] [4] 22

14 (a) (i) Colourless [1] to orange/yellow/brown [1] [2]

(ii) Cl 0 to –1 [1] Br –1 to 0 [1] Br oxidised and Cl reduced [1] [3]

(iii) Silver nitrate solution [1] (white) precipitate [1] add dilute ammonia [1] no precipitate remaining [1] [4]

(b) (i) HF is weakest acid [1] because the HF bond is strongest [1] [2]

(ii) 4HF + SiO2 → SiF4 + 2H2O [2] 13







15 (a) (i) Outermost electrons in an s orbital [1]

(ii) Atoms having the same atomic number [1] different mass numbers [1] [2]

(iii) RAM = (24 × 0.8) + (25 × 0.1) + (26 × 0.1) = 24.3 [3]

(b) (i) 2–

Mg O××

×××× O××

××××Mg2+

[3]

(ii) Ionic [1]

(iii) Conducts electricity when molten/aqueous [1] high melting point [1] [2]

(c) (i) Back [1] titration [1]

(ii) nHCl = (80 × 2)/1000 = 0.16 nNaOH = (25 × 2)/1000 = 0.05 = moles of unreacted HCl nHCl reacting with MgO = 0.16 – 0.05 = 0.11 nMgO = 0.11 ÷ 2 = 0.055 mass MgO = (0.055/5) × (24 + 16) = 0.44 g = 440 mg [6] 19

16 (a) Hydrogen bonds in ice are fixed [1] holding water molecules further apart/leading to open structure [1] [2]

(b) (Three/high number of) delocalised electrons (per atom) [1] move and carry charge [1] [2]

(c) A lot of heat/energy required [1] to break strong covalent bonds/to break strong bonds in giant covalent structure [1] [2] 6

Section B 80

Total 100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in Question 16(b)(iii).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of the Elements, containing some data, is included in this question paper.


2013

Chemistry



[AC112]








Section A



1 Which one of the following is the formula for nitrogen(I) oxide?

A NO

B NO2

C N2O

D N2O4

2 Which one of the following is the number of atoms present in 0.25 moles of C12H22O11?

A 6.8 3 1024

B 1.4 3 1025

C 2.7 3 1025

D 1.1 3 1026

3 Which one of the following is a molecular covalent substance?

A CaO

B CO

C Cr2O3

D CuO

4 A caesium atom differs from a caesium ion because the atom has a greater

A atomic number.

B mass number.

C number of electrons.

D number of protons.







5 Part of the mass spectrum for aspirin is shown below. Which one of the following numbers is the molecular ion peak?

percentageabundance

93

137

179

180

m/z

A 93

B 137

C 179

D 180

6 In which one of the following liquids are the van der Waals forces greatest?

A Argon

B Krypton

C Neon

D Xenon

7 Prozac tablets contain 20 mg of fluoxetine (C17H18F3NO) in each tablet. The number of moles of fluoxetine in each tablet is

A 6.47 3 1025

B 1.39 3 1024

C 6.47 3 1022

D 1.39 3 1021







8 Which one of the following does not have a total of 14 electrons?

A CO

B Li2O

C N2

D S22

9 Successive ionisation energies for elements X and Y are shown below.

Ionisationenergy

(kJ mol21)1st 2nd 3rd 4th 5th 6th 7th 8th

X 578 1817 2745 11 577 14 842 18 379 23 326 27 465

Y 1314 3388 5301 7469 10 990 13 327 71 330 84 078

Which one of the following is the formula for a compound of X and Y?

A XY2

B X2Y

C X2Y3

D X3Y2

10 Hexan-1-ol can be converted to hex-1-ene as follows:

C6H13OH C6H12 1 H2O

40.0 g of hexan-1-ol produced 24.7 g of hex-1-ene. Which one of the following is the percentage yield?

A 24.7%

B 50.8%

C 72.0%

D 75.0%







Section B


11 (a) Atoms consist of protons, neutrons and electrons.

(i) Complete the table below giving the properties of a proton, a neutron and an electron.

Relative mass Relative charge

Proton

Neutron

Electron

[3]

(ii) Element 116, ununhexium, was added to the Periodic Table in June 2011. Complete the table below.

Atomic number 116

Mass number

Number of protons

Number of neutrons 177

Number of electrons

[3]

(b) Iron is the sixth most abundant element in the Universe. It has four isotopes as shown in the table.

Isotope 54Fe 56Fe 57Fe 58Fe

Percentage abundance 5.84 91.76 2.12 0.28

(i) Explain what is meant by the term isotope.

[2]

(ii) Use the table to calculate the relative atomic mass of iron to two decimal places.

[2]







12 The emission spectrum for atomic hydrogen has been used to provide evidence for discrete electron energy levels in atoms.

(a) Complete the diagram to show the electron transitions associated with the first two lines of the hydrogen emission spectrum in the visible region.

n 5 5

n 5 4

n 5 3

n 5 2

n 5 1 [2]

(b) The convergence limit of the hydrogen spectrum in the ultraviolet region is at 3.28 3 1015 Hz. Calculate the ionisation energy of hydrogen in kJ mol21.

[3]

(c) The emission spectra of elements give rise to characteristic flame colours. Complete the table below.

Flame colour Formula of metal ion

Blue-green

Crimson

Green

[3]







13 Wood vinegar, which contains ethanoic acid, is formed when wood is heated. The percentage by mass of ethanoic acid in wood vinegar can be found by titration with standard sodium hydroxide solution.

(a) (i) What is meant by the term standard solution?

[1]

(ii) Write the equation for the reaction between ethanoic acid and sodium hydroxide.

[1]

(b) 25.0 cm3 of wood vinegar were diluted to 250 cm3 in a volumetric flask. 25.0 cm3 of the diluted wood vinegar required 30.3 cm3 of 0.1 mol dm23 sodium hydroxide solution for neutralisation.

(i) How many moles of sodium hydroxide were required?

[1]

(ii) How many moles of ethanoic acid were present in the 25.0 cm3 of diluted wood vinegar?

[1]

(iii) How many moles of ethanoic acid were present in 25.0 cm3 of undiluted wood vinegar?

[1]

(iv) What was the mass of ethanoic acid in the 25.0 cm3 of undiluted wood vinegar?

[1]

(v) What was the percentage of ethanoic acid by mass in the wood vinegar? Assume that the density of wood vinegar is 1.02 g cm23.

[1]







(c) Suggest a suitable indicator for the titration and state the colour change at the end point.

Indicator:

Colour change: from

to [3]







14 The Third Period from sodium to argon can be used to illustrate trends in the Periodic Table.

(a) In which block of the Periodic Table is argon found? Explain your answer.

[2]

(b) The graph below shows the melting points of the elements in the Third Period.


1800

1600

1400

1200

1000

800

600

400

200

0

mel

ting

poin

t/K

(i) Explain the rise in melting point from sodium to magnesium.

[2]







(ii) Explain why silicon has the highest melting point.

[2]

(iii) Explain why the melting point of sulfur, S, is higher than phosphorus, P.

[2]

(c) State and explain the trend in atomic radius across the Third Period.

[3]







15 Aluminium chloride exists as the molecule AlCl3 in the vapour state. This molecule contains covalent bonds and does not obey the octet rule.

(a) (i) Explain what is meant by the term covalent bond.

[2]

(ii) Explain what is meant by the term octet rule.

[2]

(b) Aluminium chloride reacts with chloride ions as follows:

AlCl3 1 Cl2 AlCl42

(i) Draw dot and cross diagrams, using outer electrons only, to show the bonding in AlCl3 and AlCl4

2.

[4]

(ii) What type of bond is formed between AlCl3 and the Cl2 ion?

[1]

(iii) Draw and name the shapes of AlCl3 and AlCl42.

[4]







16 The halogens form Group VII of the Periodic Table.

(a) The table below gives some of the physical properties of the halogens.

ElementAtomicradius(nm)

Boiling point(°C)

Electronegativity value

First ionisationenergy

(kJ mol21)

Fluorine 0.133 2187 4.0 1618

Chlorine 0.181 235 3.0 1256

Bromine 0.196 59 2.8 1143

Iodine 0.219 183 2.0 1009

(i) Explain why the atomic radii of the halogens increase as the Group is descended.

[1]

(ii) Explain the trend in the boiling points of the halogens.

[2]

(iii) Explain what is meant by the term electronegativity.

[1]

(iv) Explain the trend in electronegativity values of the halogens.

[2]







(v) Write an equation, including state symbols, for the first ionisation energy of fluorine.

[2]

(vi) Explain the trend in the first ionisation energy of the halogens.

[2]

(b) Chlorine is used to sterilise water.

(i) Write an equation for the reaction of chlorine with water.

[1]

(ii) Using changes in oxidation number explain why this is considered to be a disproportionation reaction.

[3]

(iii) Ultraviolet light does not react with water and is equally effective as chlorine at sterilising water. Suggest the advantages and disadvantages of storing and using chlorine to sterilise water.

[3]








(c) Iodide ions react with a variety of reagents. For each of the following state what you would observe and write an equation for the reaction.

(i) Chlorine gas with aqueous iodide ions.

Observation

Equation [3]

(ii) A solution containing excess Fe3+ ions with aqueous iodide ions.

Observation

Equation [3]

(iii) Silver nitrate solution with aqueous iodide ions.

Observation

Equation [3]








2013




[AC112]


MARKSCHEME







General Marking Instructions

IntroductionMark schemes are published to assist teachers and students in their preparation for examinations. Through the mark schemes teachers and students will be able to see what examiners are looking for in response to questions and exactly where the marks have been awarded. The publishing of the mark schemes may help to show that examiners are not concerned about fi nding out what a student does not know but rather with rewarding students for what they do know.

The Purpose of Mark SchemesExamination papers are set and revised by teams of examiners and revisers appointed by the Council. The teams of examiners and revisers include experienced teachers who are familiar with the level and stan-dards expected of students in schools and colleges.

The job of the examiners is to set the questions and the mark schemes; and the job of the revisers is to review the questions and mark schemes commenting on a large range of issues about which they must be satisfi ed before the question papers and mark schemes are fi nalised.

The questions and the mark schemes are developed in association with each other so that the issues of differentiation and positive achievement can be addressed right from the start. Mark schemes, therefore, are regarded as part of an integral process which begins with the setting of questions and ends with the marking of the examination.

The main purpose of the mark scheme is to provide a uniform basis for the marking process so that all the markers are following exactly the same instructions and making the same judgements in so far as this is possible. Before marking begins a standardising meeting is held where all the markers are briefed using the mark scheme and samples of the students’ work in the form of scripts. Consideration is also given at this stage to any comments on the operational papers received from teachers and their organisations. Dur-ing this meeting, and up to and including the end of the marking, there is provision for amendments to be made to the mark scheme. What is published represents this fi nal form of the mark scheme.

It is important to recognise that in some cases there may well be other correct responses which are equally acceptable to those published: the mark scheme can only cover those responses which emerged in the ex-amination. There may also be instances where certain judgements may have to be left to the experience of the examiner, for example, where there is no absolute correct response – all teachers will be familiar with making such judgements.







Section A

1 C

2 A

3 B

4 C

5 D

6 D

7 A

8 D

9 C

10 D


Section A 20







Section B

11 (a) (i) Relative mass Relative charge

Proton 1 1+Neutron 1 0Electron

18001

20001

- negligible 1–

error [–1] [3]

(ii) Atomic number 116Mass number 293

Number of protons 116Number of neutrons 177Number of electrons 116

([1] each) [3]

(b) (i) Atoms with the same atomic number/number of protons [1] but with a different mass number/number of neutrons [1] [2]

(ii) ((54 × 5.84) + (56 × 91.76) + (57 × 2.12) + (58 × 0.28))/100 = 55.91 ([–1] for each mistake, [–1] if not correct to 2 decimal places) [2] 10

12 (a) Arrow from n = 3 to n = 2 [1] Arrow from n = 4 to n = 2 [1] [2]

(b) E = hf = (6.63 × 10–34) × (3.28 1015) = 2.175 × 10–18

(2.175 × 10–18) × (6.02 × 1023) = 1 309 350 J mol–1 1309 (kJ mol–1) ([–1] for each mistake, [–1] if not kJ mol–1) [3]

(c) Flame colour Metal ionBlue-green Cu2+

Crimson Li+

Green Ba2+

(Must be formula not name, [1] each) [3] 8







13 (a) (i) A solution of (accurately) known concentration [1] (ii) CH3COOH + NaOH → CH3COONa + H2

(b) (i) (0.1 × 30.3)/1000 = 3.03 × 10–3 [1]

(ii) 3.03 × 10–3 [1]

(iii) 3.03 × 10–2 [1]

(iv) (3.03 × 10–2) × 60 = 1.82 g units needed [1]

(v) 1.8225 × 1.02 × 100 = 7.14(%) [1]

(c) Phenolphthalein [1] From colourless [1] to pink/red [1] [3] 10

14 (a) p(-block) [1] Outer electrons in the p-orbital [1] [2]

(b) (i) Increasing number of valence/outer/free electrons [1] Greater attraction between these and the (fixed) cations [1] [2]

(ii) Strong covalent bonds [1] throughout a giant structure [1] (reference to ionic bonding [0]) [2]

(iii) S8 → P4 – More atoms/greater mass/more electrons [1] Greater van der Waals forces (between the molecules) [1] [2]

(c) Atomic radius decreases across the Period/from sodium to argon [1] (Outer) electrons are in the same energy level/shielding remains the same [1] Nuclear charge increases causing greater attraction between the nucleus and the (outer) electrons [1] [3] 11







15 (a) (i) A shared pair of electrons (between two atoms) [1] Each atom provides one electron [1] [2]

(ii) When forming a compound an atom tends to gain, lose or share electrons to achieve eight [1] in its outer shell [1] [2]

(b) (i)

×•

Cl Al

Cl

×•

Cl

Cl

×•×•

Cl Al Cl

Cl

([2] each, [–1] for each mistake, [–1] if dot and cross not used) [4] (ii) Dative/co-ordinate bond

ClAlCl

Cl

[1]

(iii) trigonal planar (diagram [1], name [1])

tetrahedral (diagram [1], name [1]) [4] 13

ClAlClCl

Cl

–







16 (a) (i) As the Group is descended there are more energy levels [1]

(ii)electrons [1]

causing van der Waals forces between the molecules [1] [2]

(iii) The extent to which an atom attracts the (bonding) electrons in a covalent bond [1]

(iv) Fluorine has the smallest radius [1] greatest attraction between its nucleus and the bonding electrons [1] [2]

(v) F(g) → F+(g) + e–

([1] for equation, [1] for state symbols) [2]

(vi) Going down the Group the outer electrons are further from the nucleus [1]

Increased shielding (from the inner electrons) [1] [2] (b) (i) Cl2 + H2O → HOCl + HCl [1]

(ii) Cl2 = 0 HOCl = +1 4 oxidation numbers [2] HCl = –1 Cl is both oxidised, 0 → +1 and reduced, 0 → –1 [1] [3]

(iii) Disadvantages: Storing large quantities of chlorine causes problems/ Chlorine poisonous/toxic/ Freedom of choice

Advantages: Chlorine remains in the water after it leaves the treatment plant/ Chlorine gas can be compressed/chlorine is relatively cheap To a maximum of [3] [3]


(c) (i) Colourless solution [1] turns yellow/brown [1] Cl2 + 2l– → 2Cl– + I2 [1] [3]

(ii) Yellow/orange solution [1] turns brown/yellow [1] or colourless → yellow/brown not yellow → yellow 2Fe3+ + 2I– → 2Fe2+ + I2 [1] [3]

(iii) Yellow [1] precipitate [1] Ag+ + I– → AgI [1] [3] 28

Section B 80

Total 100







TIME

1 hour 30 minutes.




The total mark for this paper is 100.Quality of written communication will be assessed in Question 12(d)(iv).

In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included in this question paper.


January 2014

Chemistry



[AC112]

thURSDAY 9 JANUARY, moRNING







Section A



1 An element in the Periodic Table has the following successive ionisation energies (kJ mol–1).

590 1145 4912 6474 8144 10496 12320

In which one of the following groups is this element found?

A Group I

B Group II

C Group III

D Group IV

2 Which one of the following is the oxidation number of hafnium in Hf F732?

A 23

B 13

C 24

D 14

3 Boron consists of the isotopes 105B and 11

5B. The relative atomic mass of the element is 10.80. Which one of the following is the approximate ratio of the number of lighter atoms to heavier atoms?

A 1:3

B 1:4

C 1:9

D 4:1







4 Which one of the following equations shows hydrogen peroxide, H2O2, behaving as a reducing agent?

A 2Fe21 1 H2O2 1 2H1 → 2Fe31 1 2H2O

B 2I2 1 H2O2 1 2H1 → I2 1 2H2O

C MnO2 1 2H1 1 H2O2 → Mn21 1 2H2O 1 O2

D PbS 1 4H2O2 → PbSO4 1 4H2O

5 The electronegativity values, not in order, for caesium, cobalt, fluorine and nitrogen are listed below. Which one of the following is the value for the cobalt atom?

A 0.70

B 1.80

C 3.00

D 4.00

6 Which one of the following molecules is linear?

A CH3CH3

B CO2

C H2O2

D H2Te

7 Which one of the following is the reason why water boils at 100 C while the hydrides of the other Group VI elements boil below 0 C?

A Hydrogen bonding between water molecules

B Ionic bonding in water molecules

C The lower molar mass of water molecules

D The stability of the bonding in water molecules







8 The first ionisation energy is shown against increasing atomic number.

h

g

f

ed

c

b

a

ij

firstionisationenergy

atomic number

Which one of the following shows a Group I element together with a Group VII element?

Group I Group VII

A b f

B b g

C h f

D h g

9 Which one of the following properties is a characteristic of astatine?

A It has an electronegativity value greater than that of iodine.

B It is a solid at room temperature and pressure.

C It oxidises bromide ions to bromine.

D Its hydride exhibits more hydrogen bonding than hydrogen iodide.

10 3.12 g of MCl 2 were dissolved in water and made up to one litre of solution. 25.0 cm3 of this solution reacts with 7.5 cm3 of 0.100 M silver nitrate solution.

MCl2(aq) → M21(aq) 1 2Cl2(aq) Ag1(aq) 1 Cl2(aq) → AgCl(s)

Which one of the following Group II elements is M?

A barium

B calcium

C magnesium

D strontium







Section B


11 Complete the following table about the silver halides.

silverhalide

formula colourionic/

covalent

soluble in diluteammoniasolution

soluble in concentrated

ammoniasolution

silverfluoride

AgF white ionic yes yes

silverchloride

silverbromide

silveriodide

[4]







12 The creation of the friction match took many years and involved a variety of chemicals based on phosphorus.

The modern match is shown below. The head is a mixture of potassium chlorate, sulfur and phosphorus trisulfide held together by glue. The wood is soaked in ammonium phosphate which acts as a fire retardant.

wood

soaked in ammonium phosphate

potassium chloratesulfurphosphorus trisulfide

(a) Potassium chlorate reacts with the sulfur to form potassium chloride and sulfur dioxide as shown by the following equation.

2KClO3 1 3S → 2KCl 1 3SO2

(i) Deduce the oxidation number for each element in the reactants.

[1]

(ii) Deduce the oxidation number for each element in the products.

[1]

(iii) Explain, using these oxidation numbers, why this is a redox reaction.

[1]

(b) Potassium chlorate, KClO3, is manufactured using the reaction between chlorine and potassium hydroxide.

(i) Write the equation for the reaction.

[2]

(ii) State the conditions under which the reaction is carried out.

[1]







(c) Phosphorus trisulfide is easily ignited. It provides the heat to initiate the reaction between potassium chlorate and sulfur.

(i) Phosphorus has an oxidation number of 13 in phosphorus trisulfide. State the formula of phosphorus trisulfide.

[1]

(ii) Suggest whether phosphorus trisulfide is ionic or covalent. Explain your reasons.

[1]

(iii) Name the two products formed when phosphorus trisulfide is completely burnt. No oxidation numbers are needed.

[2]

(d) Ammonium phosphate is an ionic compound consisting of ammonium and phosphate ions, PO4

32.

(i) Write the formula of the ammonium ion.

[1]

(ii) Name and draw the shape of the ammonium ion stating the interbond angle.

[3]

(iii) Write the formula of ammonium phosphate.

[1]







(iv) State and explain three physical properties you would expect ammonium phosphate to have.

[3]








13 Francium is found in Group I of the Periodic Table and was discovered by Marguerite Perey in 1939 in the Curie Laboratory in France. It was isolated from uranium ore. Since then it has been synthesised by the nuclear reaction of oxygen atoms with gold atoms. It exists as 34 isotopes.

In the Periodic Table it has an atomic number of 87 and is given a relative atomic mass of 223.

(a) Francium is found in period 7 of the Periodic Table and is regarded as an s-block element.

Suggest the subshell occupied by the outermost electron in a francium atom.

[1]

(b) Francium was first synthesised according to the following equation.

197Au 1 18O → 210Fr 1 5n

The symbol n represents a neutron.

(i) What is the relative mass of a neutron?

[1]

(ii) Using the relative mass of the neutron from part (i) show, by calculation, that the equation is balanced according to mass.

[2]

(iii) Why are electrons not used when balancing the equation according to mass?

[1]







(c) Francium is one of the least electronegative elements in the Periodic Table.

(i) Explain the meaning of the term electronegativity.

[2]

(ii) State how electronegativity values change on going across a period.

[1]

(d) Francium has a melting point of 27 C and would melt in the hand just as caesium does. It has the highest electrical conductivity of the alkali metals.

(i) Explain, in terms of metallic bonding, why francium has a low melting point.

[2]

(ii) Explain, in terms of metallic bonding, why francium has the highest electrical conductivity.

[2]







(e) Francium loses electrons when it reacts with chlorine and the chlorine gains these electrons.

(i) Write the equation for the loss of an electron from a francium atom.

[1]

(ii) Write the equation for the formation of chloride ions from a chlorine molecule.

[1]

(iii) Write the equation for the reaction of francium with chlorine.

[1]

(iv) Francium chloride exists as a lattice structure similar to that of NaCl. Explain the term lattice structure.

[1]







14 The energy levels of a hydrogen atom are shown below and the arrows indicate the transition of electrons between successive energy levels.

X

Y

Z

n � 6n � 5

n � 4

n � 3

n � 2

n � 1

The electromagnetic spectrum is shown below.

Radiowaves

Microwaves Infrared Visible Ultraviolet X-raysGamma

rays

Energy increases →

(a) Write the equation that relates energy to frequency, explain the meanings of the symbols used and state the units in which they are measured.

[3]







(b) There are three series of lines. The first series X occurs in the ultraviolet region of the electromagnetic spectrum.

(i) In which part of the electromagnetic spectrum does the second series, Y, occur?

[1]

(ii) Suggest in which part of the electromagnetic spectrum the third series, Z, occurs.

[1]

(iii) What happens to the atom when its electron passes from energy level n 5 1 to an infinite energy level?

[1]

(c) (i) Use the energy level diagram below to show the ground state of a sodium atom. Use arrows to represent the sodium electrons.

3s

2p

2s

1s [2]

(ii) What is the predominant colour in the emission spectrum of sodium?

[1]







15 Iron reacts with dilute hydrochloric acid to form iron(II) chloride, FeCl2, and hydrogen. The solution deposits crystals of hydrated iron(II) chloride.

(a) Write the ionic equation, with state symbols, for the reaction of iron with hydrochloric acid.

[2]

(b) A solution of iron(II) ions is oxidised by chlorine water to form iron(III) ions.

(i) Write the ionic equation for the reaction.

[1]

(ii) Describe the colour of the solution after the reaction has taken place.

[1]

(iii) Explain whether iron(II) ions would react with bromine water.

[1]

(c) Iron(II) chloride is extremely soluble in water. 69 g of the anhydrous solid dissolve in 100 cm3 of water at 20 C. Assuming there is no volume change calculate the molarity of the resulting solution.

[3]







(d) 14.1 g of the hydrated iron(II) chloride crystals contain 6.5 g of water. Use these figures to calculate the formula of the crystals.

mass of iron(II) chloride

moles of iron(II) chloride

moles of water

ratio of moles of water to moles of iron(II) chloride

formula of iron(II) chloride crystals

[5]







16 Sulfur forms the following fluorides:

sulfur difluoride SF2

sulfur tetrafluoride SF4 sulfur hexafluoride SF6

Sulfur hexafluoride is the best known and can be used as a safe electrical insulator. The other fluorides are toxic.

(a) Draw the dot and cross diagrams showing the outer electrons only for each of the fluorides.

[3]

(b) (i) State the octet rule.

[2]

(ii) Explain whether sulfur is obeying the octet rule in each fluoride.

[2]

(c) Sulfur difluoride has the same shape as a water molecule but the bond angle is 6 smaller. Draw and name the shape of sulfur difluoride, stating its bond angle.

[3]







(d) The sulfur hexafluoride molecule has an octahedral shape.

(i) State the bond angle(s) in the sulfur hexafluoride molecule.

[1]

(ii) Explain why sulfur hexafluoride has an octahedral shape.

[2]

(iii) Explain why sulfur hexafluoride is a non-polar molecule.

[1]

(e) Sulfur tetrafluoride has a boiling point of –38 C whereas sulfur hexafluoride has a boiling point of –64 C.

(i) Which compound has the higher boiling point?

[1]

(ii) Explain, in terms of mass, which compound has the greater van der Waals forces.

[1]

(iii) Explain, in terms of intermolecular forces, the difference in boiling points.

[2]







MARKSCHEME


January 2014

Chemistry



[AC112]








Section A

1 B

2 D

3 B

4 C

5 B

6 B

7 A

8 C

9 B


Section A 20







Section B

11 Each mistake is [–1]

AgCl white ionic yes yes AgBr cream ionic no yes Agl yellow ionic no no [4] 4

12 (a) (i) K = +1; Cl = +5; O = –2; S = 0 [1]

(ii) K = +1; Cl = –1; O = –2; S = +4 [1]

(iii) Cl oxidation number goes down and S oxidation number goes up [1]

(b) (i) 6KOH + 3Cl2 → KClO3 + 5KCl + 3H2O [2]

(ii) hot and concentrated (potassium hydroxide solution) [1]

(c) (i) P2S3 [1]

(ii) covalent as both elements are non-metals/covalent as small difference in electronegativity between P and S [1]

(iii) phosphorus trioxide or pentoxide/phosphorus oxide [1] sulfur oxide or sulfur dioxide [1] [2]

(d) (i) NH4+ [1]

(ii) tetrahedral [1] 109º/109.5° [1]

H

N

HH H

[1] [3]

(iii) (NH4)3PO4/PO4(NH4)3 [1]

(iv) melting point which is high because of attraction between ions [1] boiling point which is high because of attraction between ions [1] electrical conductivity which is high in solution or molten because ions can move [1] solubility, ions are surrounded by H2O molecules [1] 3 from 4 [3]


+







13 (a) 7s1 [1]

(b) (i) 1 [1]

(ii) LHS = 197 + 18 = 215 RHS = 210 + 5 × 1 = 215 [2]

(iii) electrons have negligible mass/one two thousandth relative mass [1]

(c) (i) the extent to which an atom attracts the bonding electrons in a covalent bond [2]

(ii) electronegativity increases across a period [1]

(d) (i) sea of positive ions surrounded by electrons [1] attraction between positive ions and electrons is weak [1] [2]

(ii) electrons move and conduct/carry electricity [1] large radius of Fr means less attraction for electrons [1] [2]

(e) (i) Fr → Fr+ + e [1]

(ii) Cl2 + 2e → 2Cl– [1]

(iii) 2Fr + Cl2 → 2FrCl [1]

(iv) regular arrangement of (francium and chloride) ions [1] 16







14 (a) E = hf E is energy, h is Planck’s constant, f is frequency E in joules, h in J s, f in Hz or s–1 [3]

(b) (i) visible [1]

(ii) infrared [1]

(iii) the atom has ionised/lost the electron [1]

(c) (i) -

-.-.-.

-.

-. [2]

(ii) yellow/orange [1] 9

15 (a) Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g) [2]

(b) (i) 2Fe2+ + Cl2 → 2Fe3+ + 2Cl– [1]

(ii) yellow/orange [1]

(iii) yes, bromine is a sufficiently powerful oxidising agent [1]

(c) FeCl2 = 56 + 2 × 35.5 = 127 69 g = 69/127 = 0.54 mol = 5.4 M [3]

(d) mass of iron(II) chloride = 14.1 – 6.5 = 7.6 moles of iron(II) chloride = 7.6/127 = 0.0598 moles of water = 6.5/18 = 0.36 ratio of moles = 0.36:0.0598 = 6.02 FeCl2.6H2O [5] 13







16 (a) SF2

[1]

SF4

S×

××

×

× ×

[1]

SF6

S ××

×

××

×

[1] [3]

(b) (i) when forming a compound an atom tends to gain, lose or share electrons in its outer shell to achieve 8 [2]

(ii) SF2 yes; SF4 no; SF6 no, there are 10 and 12 electrons around S [2]

(c) shape [1]

bent [1] 104.5° – 6° = 98.5° [1] [3]

(d) (i) 90° [1]

(ii) the six bonds repel each other [1] as far apart from each other as possible [1] [2]

(iii) the molecule is symmetrical or dipoles cancel [1]

(e) (i) sulfur tetrafluoride [1]

(ii) sulfur hexafluoride has the greatest mass hence the greatest van der Waals [1]

(iii) SF4 is polar (SF6 is non-polar) [1] polar forces greater than van der Waals [1] [2] 18

Section B 80

Total 100

S× ×××

××

SF F









INFORMATION FOR CANDIDATESThe total mark for this paper is 100.Quality of written communication will be assessed in Question 12(b).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of the Elements, containing some data, is included in this question paper.


2014

Chemistry



[AC112]

MONDAY 9 JUNE, AFTERNOON








Section A



1 Which one of the following is not a redox reaction?

A 2Ca(NO 3)2 → 2CaO 4NO2 O2

B Cl2 2I− → I2 2Cl−

C Fe Cu2 → Fe2 Cu

D H2SO4 2NaOH → Na2SO4 2H2O

2 The graph of first ionisation energy against atomic number for a series of ten consecutive elements in the Periodic Table is shown below. Which one of the following indicates a Group II metal and a halogen?

a

b

c

d

f

e

g

h

i

j

atomic number

firstionisationenergy

Group II metal Halogen

A a h

B b g

C c h

D c i







3 Which one of the following is the strongest reducing agent?

A F−

B F2

C I−

D I2

4 4.35 g of potassium sulfate is dissolved in water and made up to 50.0 cm3. Which one of the following is the concentration of potassium ions in this solution?

A 0.025 mol dm–3

B 0.500 mol dm–3

C 0.644 mol dm–3

D 1.000 mol dm–3

5 Which one of the following describes the trend in bond energies of the halogen molecules down Group VII?

A Decreases

B Decreases to bromine then increases

C Increases

D Increases to chlorine then decreases

6 When 0.28 g of a basic oxide, MO, is reacted with 250 cm3 of 0.05 mol dm–3 hydrochloric acid the excess acid required 50 cm3 of 0.05 mol dm–3 sodium hydroxide solution for neutralisation. Which one of the following is the relative atomic mass of M?

A 12

B 28

C 40

D 56







7 Which one of the following diagrams represents the distribution of electrons in the 3d and 4s subshells in the ground state of an iron(III) ion?

A 1

1

1

1

1

B 1 1 1 1 1

C 1

1

1 1

1

D 1 1 1 1

1

8 Which one of the following describes the reaction between solid sodium chloride and concentrated sulfuric acid?

A Disproportionation

B Exothermic

C Neutralisation

D Redox

9 Chlorine was bubbled through a pale green solution causing the solution to turn yellow/orange. Which one of the following ions was in the original solution?

A Br –

B Fe2

C Fe3

D I–

3d 4s







10 Which one of the following molecules does not contain a polar bond?

A Fluorine

B Hydrogen fluoride

C Oxygen difluoride (OF2)

D Tetrafluoromethane (CF4)







Section B


11 Boron is the only element in Group III of the Periodic Table which is not a metal.

(a) On the axes below sketch a graph to show the successive ionisation energies of boron.

number of electrons being removed

logionisationenergy

1 2 3 4 5

[3]

(b) Boron trifluoride can react with a fluoride ion as shown in the equation below:

BF3 F– → BF4–

(i) Draw a dot and cross diagram for the BF4– ion and use it to

suggest the shape of the ion and its bond angle.

Shape _____________________________

Bond angle _____________________________ [4]

(ii) Name the type of bond formed between the fluoride ion and boron.

[1]







12 Phosphorus is a non-metal with a low melting point. It reacts explosively with liquid bromine and more gently with bromine vapour. In each case phosphorus tribromide is formed.

(a) (i) Write an equation for the reaction of phosphorus, P4, with bromine.

[2]

(ii) State the octet rule and explain whether or not phosphorus obeys the octet rule in phosphorus tribromide.

[3]

(b) The melting points of silicon, phosphorus and sulfur are given in the table below.

element Si P4 S8

melting point/°C 1410 44 113

With reference to the structures of silicon and sulfur explain why each has a higher melting point than phosphorus.

[4]








13 Sodium is a reactive, soft, silvery metal. Chlorine is a poisonous gas. The two react together to form sodium chloride that is essential to our diet.

(a) (i) Using a labelled diagram explain the bonding in sodium metal.

[3]

(ii) Metals are good conductors of electricity. Explain why the electrical conductivity of aluminium is greater than that of sodium.

[2]

(b) What type of structure is present in the element chlorine?

[1]

(c) (i) Draw dot and cross diagrams to show how sodium bonds with chlorine gas. Only outer shell electrons should be shown.

[3]

(ii) Name the type of bonding in sodium chloride.

[1]







(iii) The structure of sodium chloride is described as a lattice. Explain what is meant by the term lattice.

[2]

(iv) Apart from its appearance give three physical properties of sodium chloride.

[3]

(d) Sodium chloride can be made by reacting sodium carbonate with hydrochloric acid.


[2]

(ii) Using the following headings calculate the mass of sodium chloride formed when 5.3 g of sodium carbonate is reacted with

0.06 dm3 of 1.5 mol dm–3 hydrochloric acid.

Number of moles of sodium carbonate used

Number of moles of hydrochloric acid used

State which reagent is in excess

Number of moles of sodium chloride formed

Mass of sodium chloride formed in grams

[5]







14 Bromine tablets are used as a disinfectant in hot tubs and some swimming pools because of bromine’s ability to act as an oxidising agent.

(a) Bromine reacts with water in a similar way to chlorine.

(i) Suggest the equation for the reaction of bromine with water.

[1]

(ii) Using oxidation numbers explain why this reaction is an example of disproportionation.

[3]

(b) Manufacturers recommend maintaining the bromine concentration in swimming pools at 4 mg per litre. Calculate the molarity of bromine, Br2, in the water at this concentration.

[2]

(c) Occasionally a ‘shock treatment’ with chlorine is required to further disinfect the water.

(i) Suggest, in chemical terms, why chlorine is used for this purpose.

[1]

(ii) The compound used to provide the chlorine for the shock treatment is “sodium dichlor”, NaCl2C3N3O3. Calculate the percentage of chlorine in “sodium dichlor” to one decimal place.

[2]







(d) Bromine is produced from the reaction of sodium bromide with concentrated sulfuric acid. Name four other products formed when sodium bromide reacts with concentrated sulfuric acid.

1. ___________________________________________________

2. ___________________________________________________

3. ___________________________________________________

4. _________________________________________________ [4]

(e) Describe how you could show that a solution contains bromide ions.

[3]







15 (a) The first three ionisation energies of calcium are given in the table below.

1st ionisation energy

2nd ionisation energy

3rd ionisation energy

590 kJ mol–1 1145 kJ mol–1 4912 kJ mol–1

(i) Write the equation for the second ionisation of calcium including state symbols.

[2]

(ii) Using the following headings calculate the amount of energy, in kJ, required to form 8.0 g of Ca2(g) ions from Ca(g).

Energy required to form one mole of Ca2(g) from one mole of Ca(g)

[1]

Number of moles of Ca2(g) in 8.0 g

[1]

Energy required to form 8.0 g of Ca2(g)

[1]

(b) The Ca2 ion has the same electron arrangement as an argon atom. (i) Write the electron arrangement for the Ca2 ion.

[1]

(ii) The first ionisation energy of argon is 1520 kJ mol–1. Explain why the third ionisation energy of calcium is much higher than the first

ionisation energy of argon.

[2]







(c) The table below shows the relative abundance of the four main isotopes of calcium.

isotope 40Ca 42Ca 43Ca 44Ca

relative abundance 96.9% 0.6% 0.2% 2.3%

(i) What is meant by the term isotopes?

[2]

(ii) Calculate the relative atomic mass of calcium to two decimal places.

[2]

(iii) Complete the following table to show the number of subatomic particles in a 43Ca atom.

neutrons electrons protons43Ca

[2]

(d) A line emission spectrum of calcium, shown below, can be observed through a spectroscope.

frequency

(i) Draw an arrow in the box under ‘frequency’ pointing in the direction in which frequency increases. [1]







(ii) Describe how the movement of an electron within an atom gives rise to a line in an emission spectrum.

[3]

(iii) What flame colour is observed when calcium burns?

[1]

(iv) Using the following headings and the first ionisation energy of calcium, 590 kJ mol–1, calculate the frequency of the convergence limit of a calcium atom and state its units.

Energy, in joules, required to ionise one calcium atom

[2]

Frequency of the convergence limit of a calcium atom

[2]







MARKSCHEME

9453.01 F


2015




[AC112]


9453.01 F 2 [Turn over


IntroductionMark schemes are published to assist teachers and students in their preparation for examinations. Through the mark schemes teachers and students will be able to see what the examiners are looking for in response to questions and exactly where the marks have been awarded. The publishing of the mark schemes may help to show that examiners are not concerned about fi nding out what a student does not know but rather with rewarding students for what they do know.

The purpose of mark schemesExamination papers are set and revised by teams of examiners and revisers appointed by the Council. The teams of examiners and revisers include experienced teachers who are familiar with the level and standards expected of students in schools and colleges.



The main purpose of the mark scheme is to provide a uniform basis for the marking process so that all the markers are following exactly the same instructions and making the same judgements in so far as this is possible. Before marking begins a standardising meeting is held where all the markers are briefed using the mark scheme and samples of the students’ work in the form of scripts. Consideration is also given at this stage to any comments on the operational papers received from teachers and their organisations. During this meeting, and up to and including the end of the marking, there is provision for amendments to be made to the mark scheme. What is published represents the fi nal form of the mark scheme.

It is important to recognise that in some cases there may well be other correct responses which are equally acceptable to those published: the mark scheme can only cover those responses which emerged in the examination. There may also be instances where certain judgements may have to be left to the experience of the examiner, for example where there is no absolute correct response – all teachers will be familiar with making such judgements.

9453.01 F 3

AVAILABLEMARKS

Section A

1 B

2 B

3 C

4 A

5 D

6 C

7 C

8 D

9 A


Section A 20


AVAILABLEMARKS

Section B

11 (a) 2K + BeCl2 → 2KCl + Be [1]

(b) Be + Cl2 → BeCl2 [1] Be + 2HCl → BeCl2 + H2 [1]

(c) (i) the electrons in the bonds [1] repel equally [1] linear [1] [3] (ii) BeCl2: red, pH 0–2/strongly acidic. NaCl: green, pH 7/neutral solution. [4]

Quality of written communication [2] (d) (i) no, there are 6 electrons [1] (ii) yes, there are 8 electrons [1] (iii) electron pair donated by one atom to form a bond [2]

(iv) covalent or dipole-dipole bonds need to be broken [1] Needs energy to break bonds [1] 18

12 (a) (i) isotope protons neutrons electrons54Fe 26 28 2656Fe 26 30 2657Fe 26 31 26

[3] error [–1] (ii) (54 × 5.8) + (56 × 91.6) + (57 × 2.6)

100 = 55.9

[1] deducted for an error [2]

(iii) same number/structure of electrons 4 both needed

identical chemical properties [1]

(b) (i) 1s2 2s2 2p6 3s2 3p6 3d6 [1]

(ii) Cl2 + 2Fe2+ → 2Cl– + 2Fe3+

[1] for correct formula [1] for correct ratio [2]

(iii) electronic configuration of Fe3+ contains 3d5 [1] half-filled subshell has increased stability [1] [2] 11

9453.01 F 5

AVAILABLEMARKS

13 (a) (i) 2Na + O2 → Na2O2 [1] (ii) Na+ : O2–

1.2 × 1024 : 6.0 × 1023

2 : 1 Na2O [1]

(b) 500 000 6.02 × 1023 = 8.306 × 10–19 J

8.306 × 10–19

6.63 × 10–34 = 1.253 × 1015

3.0 × 108

1.253 × 1015

2.394 × 10–7 239.4 nm [4]

(c) (i) 2NH3 + 2Na → 2NaNH2 + H2 [1] (ii) 1s 2s 2p N ↑↓ ↑↓ ↑ ↑ ↑ [1]

N– ↑↓ ↑↓ ↑↓ ↑ ↑ [1]

(iii)

N

• •

H

H • •

or or correct dot and cross diagram

[1]

(iv) bent/V-shaped/non-linear [1]

(v) ammonia – 1 lone pair amide ion – 2 lone pairs4 both needed for [1] lone pair–lone pair repulsion on amide [1] > lone pair–bond pair repulsion on ammonia/pushes bond pairs closer/reduces bond angles [1] [3] 14

N

HH

• • • •


AVAILABLEMARKS

AVAILABLEMARKS14 (a) (i) 234 × 103

58.5 = 4 × 103

NaCl : NaHCO3 1 : 1 Moles of NaHCO3 = 4 × 103 [–1] for each error [2]

(ii) NaHCO3 : Na2CO3 2 : 1 Moles of Na2CO3 = 2 × 103

Mass of Na2CO3 = 2 × 103 × 106 = 212 kg [–1] for each error [2]

(b) (i) 24.31000 × 0.2 = 4.86 × 10–3 [1]

(ii) moles of Na2CO3 = 4.86 × 10–3 [1]

(iii) moles of Na2CO3 in 250 cm3 = 4.86 × 10–2 [1] (iv)

4.86 ×106.00

2- = 123.50 [1]

(v) 106 [1]

(vi) 106 + 18x = 123.50 x = 1 [1]

δ+ δ– δ– δ+

(c) (i) C═O O─H

[1] for each [2]

(ii) dipoles cancel out/molecule is symmetrical [1]

(iii) Sufficient energy at 100°C [1] to break the hydrogen bonds [1]. [2] 15

9453.01 F 7

AVAILABLEMARKS

15 (a) (i) The extent to which an atom attracts the bonding electrons in a covalent bond. [2] (ii) As Group is descended the atomic radius increases Therefore (force of) attraction from nucleus to bonding electrons decreases [2]

(iii) Moving from HCl to HI the molecules increase in RFM/mass/no. of electrons Increased strength of van der Waals forces [2]

(iv) Between molecules of HF there are (van der Waals forces and) H-bonds H-bonds (are much stronger and) require (a lot) more energy to break 2nd mark dependent on first [2]

(v) HF < HCl < HBr < HI, H–I lowest bond energy – both needed [1]

(b) (i) yellow/orange/brown to colourless

(ii) Disproportionation is the simultaneous oxidation and reduction of the same species [2] Oxidation of Br2 to BrO–, 0 → +1 Reduction of Br2 to Br–, 0 → –1 (Both required for second mark) [2]

(iii) 3Br2 + 6OH– → 5Br– + BrO3– + 3H2O [2]

(c) (i) N is potassium iodide [1]

(ii) O is potassium bromide [1] P is bromine gas [1] Q and R are sulfur dioxide and hydrogen bromide [2] 22

Section B 80

Total 100



INFORMATION FOR CANDIDATESThe total mark for this paper is 100.Quality of written communication will be assessed in Question 11(c)(ii).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included in this question paper.9453.05R

For Examiner’suse only

Question Number Marks

Section A 1–10

Section B1112131415

TotalMarks

*ac112*

AC

112

Centre Number

Candidate Number

Chemistry



[AC112]WEDNESDAY 10 JUNE, AFTERNOON


2015

9453.05R 2

Section A



1 Potassium dichromate has the formula K2Cr2O7. Which one of the following lists the oxidation numbers of potassium and chromium in potassium dichromate?

potassium chromiumA 11 13

B 11 16

C 12 13

D 12 16

2 There are three bonding pairs and one lone pair of electrons around the central phosphorus atom in phosphine (PH3). Which one of the following describes the shape of the phosphine molecule?

A Bent

B Pyramidal

C Tetrahedral

D Trigonal planar

3 Which one of the following statements represents how the visible emission line spectrum of atomic hydrogen arises?

A Energy is given out when hydrogen atoms lose electrons to form ions

B Energy is given out when electrons move from higher energy levels to the n51 energy level

C Energy is given out when electrons move from higher energy levels to the n52 energy level

D Energy is given out when electrons move from the n51 energy level to higher energy levels

9453.05R 3 [Turn over

4 The table below shows the first six successive ionisation energies for a Period 2 element.

first second third fourth fifth sixth

Ionisation Energy/ kJ mol21

1090 2350 4610 6220 37800 47000

Which one of the following elements has these ionisation energies?

A Carbon

B Fluorine

C Nitrogen

D Oxygen

5 Which one of the following elements forms an ion with a double negative charge that has the same electronic configuration as argon?

A Calcium

B Chlorine

C Selenium

D Sulfur

6 Boron trichloride reacts with water to form a strongly acidic solution as shown below.

BCl3 1 3H2O → H3BO3 1 3HCl

When 21.6 g of BCl3 is dissolved in 250 cm3 of water the concentration of the hydrochloric acid in this solution is

A 0.55 mol dm23.

B 0.74 mol dm23.

C 2.21 mol dm23.

D 2.94 mol dm23.

9453.05R 4

7 The chlorate(V) ion, ClO32, may be reduced to chlorine.

2ClO32(aq) 1 xH+(aq) 1 ye2 → Cl2(aq) 1 zH2O(l)

Which one of the following represents the correct values of x, y and z?

x y z

A 6 6 3

B 6 4 3

C 12 10 6

D 12 12 6

8 Which one of the following is the most powerful reducing agent?

A Bromine atom

B Chlorine atom

C Fluoride ion

D Iodide ion

9 Which one of the following elements would be expected to form the smallest ion with a noble gas configuration?

A Aluminium

B Chlorine

C Sodium

D Sulfur


10 Which one of the following equations represents the first ionisation energy of fluorine?

A F2(g) 1 2e2 → 2F2(g)

B F(g) 1 e2 → F2(g)

C F(g) → F1(g) + e2

D F2(g) → 2F1(g) 1 2e2

Examiner OnlyMarks Remark

9453.05R 6

Section B


11 Beryllium is a hard silver-white metal which was first isolated by Wöhler in 1828 by the reaction of potassium with beryllium chloride. Potassium being more reactive than beryllium gave a metallic solid in a strongly exothermic process.

(a) Write the equation for the reaction of potassium with beryllium chloride.

[1]

(b) Beryllium chloride can be prepared by the reaction of beryllium with chlorine or hydrogen chloride. Write equations for both of these reactions.

[2]

(c) Beryllium chloride is a covalent molecule with a melting point of 400 °C. Its electronic structure is shown below.

Be Cl•X

X X

X X

XXCl •

X

X X

X X

XX

It reacts vigorously with water.

BeCI2 + 2H2O → Be(OH)2 + 2HCI

(i) Name and explain the shape of the beryllium chloride molecule.

[3]



(ii) Beryllium chloride and sodium chloride are separately added to water. Describe and explain what is observed when Universal Indicator is added to each solution.

[4]


(d) The high melting point of beryllium chloride is explained by its polymeric structure. Part of the polymeric structure is shown below:

CIX X

XX

XX

CIX X

X •

XX

XX

BeX •

CIX X

X •

XX

XX

BeX •

CIX X

XX

XX

(i) Explain whether beryllium, in the polymeric structure, obeys the octet rule.

[1]

(ii) Explain whether chlorine, in the polymeric structure, obeys the octet rule.

[1]


9453.05R 8

(iii) Some of the chlorine atoms in the polymeric structure are forming coordinate bonds. Explain this term.

[2]

(iv) Explain why the polymeric structure has a high melting point.

[2]

BLANK PAGE

(Questions continue overleaf)



9453.05R 10

12 A sample of iron from a meteorite was found to contain the following isotopes: 54Fe, 56Fe and 57Fe.

(a) (i) Complete the table to show the number of protons, neutrons and electrons that are present in each of the isotopes.

isotope protons neutrons electrons54Fe56Fe57Fe

[3]

(ii) From the mass spectrum the relative abundances of the isotopes in this sample of iron were found to be as follows:

m/z ratio 54 56 57

% abundance 5.8 91.6 2.6

Calculate the relative atomic mass of iron to one decimal place.

[2]

(iii) Explain the difference, if any, in the chemical properties of the isotopes of iron.

[1]

(b) (i) Write the electronic configuration of an Fe21 ion.

[1]

(ii) When chlorine gas is bubbled through a solution of Fe21 ions, oxidation to Fe31 ions occurs. Write an equation for this reaction.

[2]



(iii) With reference to s,p,d notation explain the stability of the Fe31 ion relative to the Fe21 ion.

[2]


9453.05R 12

13 The combustion of Group I metals forms their oxides. Depending on the reaction conditions sodium can form the peroxide, Na2O2.

(a) (i) Write an equation for the reaction of sodium with oxygen to form

the peroxide.

[1]

(ii) At higher temperatures and pressures a different oxide Y is formed. One mole of Y contains the Avogadro number of O2- ions and 1.2 3 1024 Na1 ions. Deduce the formula of Y.

[1]

(b) If a large amount of energy is supplied to sodium vapour it ionises. The 1st ionisation energy for sodium is 500 kJ mol21. Calculate the

wavelength of energy absorbed in nm by the sodium vapour.

(1 nm = 1 3 1029 m c 5 3.0 3 108 m s21)

[4]

(c) When strongly heated sodium reacts with ammonia to form sodium amide, NaNH2, and hydrogen.

(i) Write the equation for the reaction between sodium and ammonia.

[1]



(ii) Use the boxes below to give the electronic configuration of the N atom and the N2 ion.

1s 2s 2p

N

N-

[2]

(iii) Draw the shape of an amide ion, NH22, showing any lone pairs of

electrons.

[1]

(iv) Name the shape of the amide ion.

[1] (v) Explain, in terms of electron pair repulsion, why the bond angle in

an amide ion is smaller than the bond angle in an ammonia molecule.

[3]


9453.05R 14

14 Sodium carbonate is manufactured by the Solvay process. This is a two stage process.

STAGE 1 Sodium hydrogencarbonate is formed.

NaCl 1 NH3 1 CO2 1 H2O → NaHCO3 1 NH4Cl

STAGE 2 Sodium hydrogencarbonate is then thermally decomposed.

2NaHCO3 → Na2CO3 1 H2O 1 CO2

(a) (i) Calculate the number of moles of sodium hydrogencarbonate formed from 234 kg of sodium chloride.

[2]

(ii) Calculate the maximum mass of sodium carbonate formed in kg.

[2]

(b) Sodium carbonate can form a number of hydrates of formula Na2CO3.xH2O. A 6.0 g sample of hydrated sodium carbonate was dissolved in water and the solution made up to 250 cm3. A 25.0 cm3 portion of this solution required 24.3 cm3 of 0.2 mol dm23 sulfuric acid for complete reaction.

Na2CO3 1 H2SO4 → Na2SO4 1 H2O 1 CO2

(i) Calculate the number of moles of sulfuric acid required for complete reaction.

[1]

(ii) Deduce the number of moles of sodium carbonate in 25.0 cm3 of the solution.

[1]



(iii) Calculate the number of moles of sodium carbonate in 250 cm3 of solution.

[1]

(iv) Calculate the relative formula mass of the hydrated sodium carbonate.

[1]

(v) Calculate the relative formula mass of anhydrous sodium carbonate.

[1]

(vi) Calculate the value of x.

[1]

(c) Water and carbon dioxide both contain polar bonds.

(i) Show the polarity of the carbon2oxygen bond and the oxygen2hydrogen bond on the bonds drawn below.

C O O H

[2]

(ii) Suggest why the carbon dioxide molecule is non-polar.

[1]

(iii) Explain why water changes to a gas at 100 °C.

[2]


9453.05R 16

15 The table below shows some data about the halogens, Group VII.

element electronegativity boiling point of hydrogen halide/K

bond energy of hydrogen halide/kJ mol21

fluorine 4.0 293 568

chlorine 3.0 188 431

bromine 2.8 206 366

iodine 2.5 238 299

(a) (i) Define the term electronegativity.

[2]

(ii) Explain the trend in electronegativity as the group is descended.

[2]

(iii) Explain the trend in boiling point from hydrogen chloride to hydrogen iodide.

[2]

(iv) Explain why hydrogen fluoride does not follow this trend.

[2]

(v) State and explain the order of increasing acid strength of equimolar solutions of the hydrogen halides.

[1]


9453.05R 17

(b) Bromine water reacts with cold, dilute alkali as shown below:

Br2(aq) 1 2OH2(aq) → Br2(aq) 1 BrO2(aq) 1 H2O(l)

(i) State the colour change observed during this reaction.

[2]

(ii) State the oxidation states of bromine in the reaction and use them to explain why this reaction is an example of disproportionation.

[4]

(iii) Write the ionic equation for the reaction of bromine with hydroxide ions to produce bromate(V), BrO3

2, ions.

[2]

(c) Use the information below to identify N, O, P, Q and R.

(i) When silver nitrate solution is added to a solution of a potassium halide, N, a yellow solid is formed.

N is [1]

(ii) When concentrated sulfuric acid is added to a solid potassium halide O, a red-brown gas P and two colourless gases Q and R are formed.

O is

P is

Q is

R is [4]

THIS IS THE END OF THE QUESTION PAPER

182761

Permission to reproduce all copyright material has been applied for.In some cases, efforts to contact copyright holders may have been unsuccessful and CCEAwill be happy to rectify any omissions of acknowledgement in future if notified.

MARKSCHEME

9453.01 F


2015




[AC112]










9453.01 F 3

AVAILABLEMARKS

Section A

1 B

2 B

3 C

4 A

5 D

6 C

7 C

8 D

9 A


Section A 20


AVAILABLEMARKS

Section B

11 (a) 2K + BeCl2 → 2KCl + Be [1]

(b) Be + Cl2 → BeCl2 [1] Be + 2HCl → BeCl2 + H2 [1]

(c) (i) the electrons in the bonds [1] repel equally [1] linear [1] [3] (ii) BeCl2: red, pH 0–2/strongly acidic. NaCl: green, pH 7/neutral solution. [4]

Quality of written communication [2] (d) (i) no, there are 6 electrons [1] (ii) yes, there are 8 electrons [1] (iii) electron pair donated by one atom to form a bond [2]

(iv) covalent or dipole-dipole bonds need to be broken [1] Needs energy to break bonds [1] 18

12 (a) (i) isotope protons neutrons electrons54Fe 26 28 2656Fe 26 30 2657Fe 26 31 26

[3] error [–1] (ii) (54 × 5.8) + (56 × 91.6) + (57 × 2.6)

100 = 55.9

[1] deducted for an error [2]

(iii) same number/structure of electrons 4 both needed

identical chemical properties [1]

(b) (i) 1s2 2s2 2p6 3s2 3p6 3d6 [1]

(ii) Cl2 + 2Fe2+ → 2Cl– + 2Fe3+

[1] for correct formula [1] for correct ratio [2]

(iii) electronic configuration of Fe3+ contains 3d5 [1] half-filled subshell has increased stability [1] [2] 11

9453.01 F 5

AVAILABLEMARKS

13 (a) (i) 2Na + O2 → Na2O2 [1] (ii) Na+ : O2–

1.2 × 1024 : 6.0 × 1023

2 : 1 Na2O [1]

(b) 500 000 6.02 × 1023 = 8.306 × 10–19 J

8.306 × 10–19

6.63 × 10–34 = 1.253 × 1015

3.0 × 108

1.253 × 1015

2.394 × 10–7 239.4 nm [4]

(c) (i) 2NH3 + 2Na → 2NaNH2 + H2 [1] (ii) 1s 2s 2p N ↑↓ ↑↓ ↑ ↑ ↑ [1]

N– ↑↓ ↑↓ ↑↓ ↑ ↑ [1]

(iii)

N

• •

H

H • •

or or correct dot and cross diagram

[1]

(iv) bent/V-shaped/non-linear [1]

(v) ammonia – 1 lone pair amide ion – 2 lone pairs4 both needed for [1] lone pair–lone pair repulsion on amide [1] > lone pair–bond pair repulsion on ammonia/pushes bond pairs closer/reduces bond angles [1] [3] 14

N

HH

• • • •


AVAILABLEMARKS

AVAILABLEMARKS14 (a) (i) 234 × 103

58.5 = 4 × 103

NaCl : NaHCO3 1 : 1 Moles of NaHCO3 = 4 × 103 [–1] for each error [2]

(ii) NaHCO3 : Na2CO3 2 : 1 Moles of Na2CO3 = 2 × 103

Mass of Na2CO3 = 2 × 103 × 106 = 212 kg [–1] for each error [2]

(b) (i) 24.31000 × 0.2 = 4.86 × 10–3 [1]

(ii) moles of Na2CO3 = 4.86 × 10–3 [1]

(iii) moles of Na2CO3 in 250 cm3 = 4.86 × 10–2 [1] (iv)

4.86 ×106.00

2- = 123.50 [1]

(v) 106 [1]

(vi) 106 + 18x = 123.50 x = 1 [1]

δ+ δ– δ– δ+

(c) (i) C═O O─H

[1] for each [2]

(ii) dipoles cancel out/molecule is symmetrical [1]

(iii) Sufficient energy at 100°C [1] to break the hydrogen bonds [1]. [2] 15

9453.01 F 7

AVAILABLEMARKS

15 (a) (i) The extent to which an atom attracts the bonding electrons in a covalent bond. [2] (ii) As Group is descended the atomic radius increases Therefore (force of) attraction from nucleus to bonding electrons decreases [2]

(iii) Moving from HCl to HI the molecules increase in RFM/mass/no. of electrons Increased strength of van der Waals forces [2]

(iv) Between molecules of HF there are (van der Waals forces and) H-bonds H-bonds (are much stronger and) require (a lot) more energy to break 2nd mark dependent on first [2]

(v) HF < HCl < HBr < HI, H–I lowest bond energy – both needed [1]

(b) (i) yellow/orange/brown to colourless

(ii) Disproportionation is the simultaneous oxidation and reduction of the same species [2] Oxidation of Br2 to BrO–, 0 → +1 Reduction of Br2 to Br–, 0 → –1 (Both required for second mark) [2]

(iii) 3Br2 + 6OH– → 5Br– + BrO3– + 3H2O [2]

(c) (i) N is potassium iodide [1]

(ii) O is potassium bromide [1] P is bromine gas [1] Q and R are sulfur dioxide and hydrogen bromide [2] 22

Section B 80

Total 100


2016

ChemistryAssessment Unit AS 1assessingBasic Concepts in Physicaland Inorganic Chemistry

[AC112]TUESDAY 14 JUNE, AFTERNOON


INSTRUCTIONS TO CANDIDATESWrite your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all eighteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all eight questions in Section B. You must answer the questions in the spaces provided.Do not write outside the boxed area on each page or on blank pages.Complete in blue or black ink only. Do not write with a gel pen.

INFORMATION FOR CANDIDATESThe total mark for this paper is 100.Quality of written communication will be assessed in Question 12(a).In Section A all questions carry equal marks, i.e. two marks for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included in this question paper.

*24AC11201*

*24AC11201*

Centre Number

Candidate Number

10120

*AC112*

*AC112*

10120 AC112 XXX 2016.indd 1 29/04/2016 10:45:12

*24AC11202*

*24AC11202*

10120

Section A



1 Which one of the following shows how many electron pairs can be accommodated in the third energy level, n = 3, of an atom?

A 3

B 6

C 9

D 18

2 Which one of the following molecules contains a total of six bonding electrons?

A C2H4

B CO2

C NH3

D SF6

3 Which one of the following molecules is not polar?

A CHCl3

B CH3OCH3

C CO2

D SO2

10120 AC112 XXX 2016.indd 2 29/04/2016 10:45:13

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10120[Turn over

4 An element X has the following ionisation energies:

1st 2nd 3rd 4th 5th 6th

ionisation energy/kJ mol21

738 1451 7733 10543 13630 18020

Which one of the following is the formula of the chloride of X?

A XCl

B XCl2

C XCl3

D XCl4

5 A salt gives a pink flame in a flame test when observed through cobalt glass. A solution of the salt gives a cream precipitate when added to acidified silver nitrate solution. Which one of the following is the salt?

A Potassium bromide

B Potassium chloride

C Sodium bromide

D Sodium chloride

6 Which one of the following indicators is not suitable for the acid-base titration shown?

0.1 M acid 0.2 M base indicator

A ethanoic acid ammonia solution phenolphthalein

B ethanoic acid sodium hydroxide solution phenolphthalein

C hydrochloric acid ammonia solution methyl orange

D hydrochloric acid sodium hydroxide solution methyl orange

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*24AC11204*

*24AC11204*

10120

7 Iron can be extracted from iron(III) oxide using carbon according to the following equation:

2Fe2O3 1 3C → 4Fe 1 3CO2

Which one of the following is the maximum mass of iron that can be extracted from a mixture of 150.0 tonnes of iron(III) oxide and 15.0 tonnes of carbon?

A 26.3 tonnes

B 52.6 tonnes

C 93.3 tonnes

D 105.3 tonnes

8 The atomic emission spectrum of hydrogen for the visible region is shown below:

X

frequency

Which one of the labelled transitions is responsible for line X in the spectrum?

A B C D

n 5 6n 5 5

n 5 4

n 5 3

n 5 2

n 5 1

10120 AC112 XXX 2016.indd 4 29/04/2016 10:45:14

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*24AC11205*

10120[Turn over

9 A sample of hydrated sodium sulfate contains 56%, by mass, of water. What is the formula of the hydrated sodium sulfate?

A Na2SO4.H2O

B Na2SO4.6H2O

C Na2SO4.10H2O

D Na2SO4.12H2O

10 A cup of coffee contains 500 mg of caffeine which has the chemical formula C8H10N4O2. Which one of the following is the number of nitrogen atoms present in this amount of caffeine?

A 1.55 3 1021

B 6.21 3 1021

C 1.55 3 1024

D 6.21 3 1024

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10120

Section B

Answer all eight questions in this section.

11 Complete the following table for the ions of three elements, A, B and C.

ion atomic number electronic structure of ion

A31 5

B2 1s22s22p63s23p63d104s24p6

C22 16

[3]

12 The 2010 Nobel Prize for Physics was awarded for the discovery of a new material called graphene. It consists of a single layer of carbon atoms obtained from graphite.

(a) Describe the structure and bonding of graphite. Include an explanation why graphite can conduct electricity.

[4]


10120 AC112 XXX 2016.indd 6 29/04/2016 10:45:15

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*24AC11207*

10120[Turn over

(b) Explain why graphene, like graphite, has a high melting point.

[2]

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10120

13 In 1937 the American scientists Taylor and Crist investigated the decomposition of gaseous hydrogen iodide. The hydrogen iodide was heated in a sealed quartz tube.

2HI(g) ? H2(g) 1 I2(g)

(a) Taylor and Crist were able to measure the progress of the decomposition by measuring colour intensity.

(i) State the colour of iodine gas.

[1]

(ii) Suggest what would be observed if this experiment was to be repeated with samples of hydrogen chloride and hydrogen bromide.

[2]

(iii) Explain the difference in observations between hydrogen chloride and hydrogen bromide.

[1]

(b) Hydrogen iodide dissolves in water to form hydriodic acid which is a strong acid.

(i) Explain whether hydriodic acid is a stronger or weaker acid than hydrochloric acid.

[2]

(ii) Suggest an equation for the reaction between sodium carbonate and hydriodic acid.

[2]

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10120[Turn over

(c) The boiling points of the hydrogen halides at atmospheric pressure are shown below:

hydrogen halide boiling point/°C

HF 19.9

HCl 285.0

HBr 266.7

HI 235.4

Explain why hydrogen iodide has a higher boiling point than hydrogen chloride.

[2]

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10120

14 The hydrogen atom contains one electron and is difficult to place in a particular group in the Periodic Table. It could be in either Group I or Group VII.

(a) Suggest reasons, with reference to electron structure, why hydrogen could be placed in Group I or Group VII.

(i) Group I

[1]

(ii) Group VII

[1]

(b) Hydrogen, like the halogens, exists as diatomic molecules. However, it is much less reactive because it has a stronger covalent bond than any of the halogens.

(i) State the trend in bond energy of the halogen molecules.

[2]

(ii) Suggest why hydrogen has a higher bond energy than any of the halogen molecules.

[1]

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(c) Hydrogen reacts with sodium to form sodium hydride. Ions are formed in a similar manner to sodium and chloride ions.

(i) Complete the following diagram to show how the ions are arranged in a sodium chloride lattice.

Sodium chloride

Na1

[1]

(ii) Draw a dot and cross diagram, using outer electrons only, to show the reaction between sodium and hydrogen atoms to form sodium hydride.

[3]

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10120

(iii) Sodium hydride is a powerful reducing agent and will react with water to form sodium hydroxide and hydrogen. Write an equation for this reaction.

[1]

(iv) 0.44 g of sodium hydride is reacted with 75 cm3 of water.

Calculate the number of moles of sodium hydride.

Calculate the number of moles of sodium hydroxide formed.

Calculate the molarity of the sodium hydroxide solution.

[3]

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10120

BLANK PAGE

DO NOT WRITE ON THIS PAGE


[Turn over

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10120

15 Chlorine has many industrial uses, particularly as a bleaching agent. It is used to bleach wood pulp in paper manufacture and to remove ink from paper which is to be recycled.

(a) Chlorine has two stable isotopes 35Cl and 37Cl present in nature in the following proportions.

isotope abundance35Cl 75.78 %37Cl 24.22 %

Calculate the relative atomic mass of chlorine to two decimal places.

[2]

(b) Household bleach contains sodium chlorate(I) rather than molecular chlorine. Sodium chlorate(I) can be made by reacting sodium hydroxide with chlorine gas in a disproportionation reaction.

(i) Explain what is meant by a disproportionation reaction.

[1]

(ii) Write an equation for the reaction between chlorine and sodium hydroxide to form sodium chlorate(I) and state the conditions for the formation of sodium chlorate(I).

equation [2]

conditions [1]

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(iii) Explain, in terms of bonding, why sodium chlorate(I) has a higher boiling point than chlorine.

[2]

(c) Chlorine can form a number of chlorine oxides. Complete the table below giving the oxidation number of chlorine in each chlorine oxide.

formula of chlorine oxide oxidation number of chlorine

Cl2O

ClO2

Cl2O7

[3]

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10120

16 Strontium carbonate is commonly used in fireworks and flares as it gives a red flame colour. It contains strontium ions which are isoelectronic with krypton atoms.

(a) (i) Suggest the formula and electronic configuration for the strontium ion.

[2]

(ii) Suggest the meaning of the term isoelectronic.

[1]

(b) The red light emitted by one mole of strontium ions has an energy of 171.09 kJ.

(i) Calculate the energy, in joules, emitted by one ion of strontium.

[2]

(ii) Calculate the frequency of this light.

[1]

(iii) Explain, using electronic transitions, why strontium ions give a red colour in fireworks.

[3]

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*24AC11217*

10120[Turn over

(c) 60 cm3 of 2.0 mol dm23 hydrochloric acid was added to 2.56 g of a sample from the firework. The resultant solution was filtered and made up to 500 cm3 with deionised water. 25.0 cm3 of this solution was titrated against 0.2 mol dm23 sodium hydroxide. The following results were obtained:

initial burette reading/cm3

final burette reading/cm3

titre/cm3

rough 0.0 24.9 24.9

1st accurate 24.9 49.5 24.6

2nd accurate 0.0 24.5 24.5

The reactions which occur are:

SrCO3 1 2HCl → SrCl2 1 CO2 1 H2O

NaOH 1 HCl → NaCl 1 H2O

(i) Calculate the total number of moles of hydrochloric acid added.

(ii) Calculate the number of moles of sodium hydroxide reacted.

(iii) How many moles of hydrochloric acid are there in 500 cm3 of the solution?

(iv) Calculate the number of moles of hydrochloric acid that reacted with the strontium carbonate.

10120 AC112 XXX 2016.indd 17 29/04/2016 10:45:23

*24AC11218*

*24AC11218*

10120

(v) Calculate the mass of strontium carbonate in the sample.

(vi) Calculate the percentage by mass of strontium carbonate in the sample.

[6]

10120 AC112 XXX 2016.indd 18 29/04/2016 10:45:24

*24AC11219*

*24AC11219*

10120

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[Turn over

10120 AC112 XXX 2016.indd 19 29/04/2016 10:45:25

*24AC11220*

*24AC11220*

10120

17 A typical “lithium ion battery” consists of a lithium cobalt oxide (LiCoO2) electrode and a graphite electrode separated by lithium fluorophosphate (LiPF6).

(a) (i) The dot and cross diagram for the fluorophosphate ion is shown below. State the octet rule and explain whether or not the atoms in the ion obey this rule.

• •• •

X •F• •• •

• •

X

•

F• •

• •• X

F • •

• •• • FP

• •

• •

• •FX •

• • • •

• •X • F

• •

• •

• •

X • F

2

[3]

(ii) Draw and name the shape of the PF62 ion.

[2]

10120 AC112 XXX 2016.indd 20 29/04/2016 10:45:26

*24AC11221*

*24AC11221*

10120[Turn over

(iii) Explain why PF62 has the shape selected.

[2]

(b) Redox reactions will occur in a working battery.

(i) Define a redox reaction.

[1]

(ii) What is the oxidation state of cobalt in LiCoO2?

[1]

(iii) The lead–acid battery, common in many motor vehicles, relies on the following redox processes.

Balance the half-equations shown below.

PbO2 1 SO422 1 H1 → PbSO4 1 H2O

Pb 1 HSO42 → PbSO4 1 H1

[2]

(iv) Combine the half-equations into an equation showing the overall reaction.

[1]

10120 AC112 XXX 2016.indd 21 29/04/2016 10:45:26

*24AC11222*

*24AC11222*

10120

18 Electronic cigarettes have been developed as an alternative to tobacco smoking. They are controversial as some studies have suggested that they release very small amounts of metal ions, such as silver, into the air.

(a) Suggest how the vapour produced by an electronic cigarette could be tested for silver ions. Indicate the result that would be expected if silver ions were present.

[3]

(b) Silver ions can be used to sterilise water, 0.001 g of silver ions being required for 1000 dm3 of water.

(i) What is the concentration of silver ions in mol dm23?

[2]

(ii) What mass of silver ions is required to sterilise an Olympic sized swimming pool which contains 2.5 3 106 dm3 of water?

[1]

10120 AC112 XXX 2016.indd 22 29/04/2016 10:45:27

*24AC11223*

*24AC11223*

10120

(c) Silver has also been used to dispose of chemical weapons such as mustard gas (C4H8SCl2), which will react with silver(II) ions. The silver(II) ion is a powerful oxidising agent.

(i) Write the formula of a silver(II) ion.

[1]

(ii) An alternative method of disposing of mustard gas is through reaction with sodium hydroxide, which produces C4H8S(OH)2 and sodium chloride. Write an equation for this reaction.

[1]


10120 AC112 XXX 2016.indd 23 29/04/2016 10:45:28

208717



*24AC11224*

*24AC11224*


QuestionNumber Marks

Section A 1–10

Section B1112131415161718

TotalMarks

10120 AC112 XXX 2016.indd 24 29/04/2016 10:45:28

Periodic Table of the ElementsFor the use of candidates taking

Advanced Subsidiary and Advanced Level Chemistry Examinations

Copies must be free from notes or additions of any kind. No other type of data booklet or information

sheet is authorised for use in the examinations.

gce A/AS examinations

chemistry(advanced)

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* 58–71 Lanthanum series† 90–103 Actinium series

a = relative atomic mass (approx.)x = atomic symbolb = atomic number

THE PERIODIC TABLE OF ELEMENTSGroup

ab x

0VIIVIVIVIIIIII

*

†

One mole of any gas at 20 °C and a pressure of 1 atmosphere (105 Pa) occupies a volume of 24 dm3.Planck Constant = 6.63 × 10–34 J sGas Constant = 8.31 J mol–1 K–1

Avogadro Constant = 6.02 × 1023 mol–1

MARKSCHEME

10120.01 F


2016




[AC112]

TUESDAY 14 JUNE, AFTERNOON

10120.01 F 2 [Turn over








10120.01 F 3

AVAILABLEMARKS

Section A

1 C

2 C

3 C

4 B

5 A

6 A

7 C

8 D

9 C

10 B [2] for each correct answer [20] 20

Section A 20

10120.01 F 4 [Turn over

AVAILABLEMARKS

Section B

11 Ion Atomic number Electronic structureA3+ 5 1s2

B– 35 1s22s22p63s23p63d104s24p6

C2– 16 1s22s22p63s23p6

[3] 3

12 (a) Each carbon covalently bonded to three others Hexagonal layers with weak forces between them Each carbon has a delocalised electron Electron moves and carries charge [4]


(b) Many strong covalent bonds which take a lot of energy to break [2] 8

13 (a) (i) violet/purple [1]

(ii) no change with HCl [1] red-brown gas forms with HBr [1] [2] (iii) HCl has higher bond enthalpy than HBr [1]

(b) (i) Hydriodic is stronger Hydrogen iodide has a weaker covalent bond [2] (ii) Na2CO3 + 2HI → 2NaI + CO2 + H2O [2]

(c) HI higher Mr /more electrons [1] Stronger/more van der Waals [1] 10

10120.01 F 5

AVAILABLEMARKS

14 (a) (i) contains 1 electron in outer subshell [1]

(ii) needs one more electron to fill shell/subshell [1]

(b) (i) bond energy increases from fluorine to chlorine [1] it then decreases down the group [1]

(ii) bond is shorter, (shorter) bonds are stronger [1]

Cl– Na+

Na+

Na+

Na+

Cl–

Cl–

Cl–

Sodium chloride

(c) (i)

[1]

(ii) Na• Hx

[Na]+ [Hx]–↓

•

(error [-1] mark) [3]

(iii) NaH + H2O → NaOH + H2 [1]

(iv) 0.0183 [1] 0.0183 [1] 0.244 (mol dm–3) [1] 13

15 (a) 35.48 [2]

(b) (i) oxidation and reduction of the same species in the same reaction [1]

(ii) 1. 2NaOH + Cl2 → NaCl + NaOCl + H2O [2]

2. conditions – cold, dilute [1]

(iii) sodium chlorate(l) is ionic – strong ionic bonds chlorine is molecular – weak van der Waals [2]

(c) formula of chlorine oxide oxidation number of chlorine

Cl2O +1ClO2 +4Cl2O7 +7

[3] 11

10120.01 F 6 [Turn over

AVAILABLEMARKS

AVAILABLEMARKS

16 (a) (i) 1s22s22p63s23p63d104s24p6 /Sr2+ [2] (ii) Same number of electrons/same electronic structure [1]

(b) (i) 2.84 × 10–19 (J) [2] (ii) 4.28 × 1014 (Hz) [1]

(iii) Electrons excited to higher energy levels. Electrons fall to lower levels. Energy emitted as red light [3]

(c) (i) 0.12 (ii) 0.0049 (iii) 0.098

(iv) 0.022

(v) 1.63 g

(vi) 63.7% [6] 15

17 (a) (i) (When forming a compound,) an atom (tends to) gain, lose or share electrons to achieve eight in its outer shell. [1] Phosphorus does not as it has 12 electrons in the outer shell [1] fluorine does [1] [3] (ii)

F

P

F FF

F F

Octahedral [2]

(iii) six bonding pairs [1] repel equally to maximise separation [1] [2]

(b) (i) oxidation and reduction occur in the same reaction [1]

(ii) +3 [1]

(iii) PbO2 + SO42– + 4H+ + 2e– → PbSO4 + 2H2O [1]

Pb + HSO4– → PbSO4 + H+ + 2e– [1]

(iv) PbO2 + Pb + SO42– + HSO4

– + 3H+ → 2PbSO4 + 2H2O [1] 12

10120.01 F 7

AVAILABLEMARKS

18 (a) Bubble vapour through A solution of sodium chloride or hydrochloric acid White precipitate formed [3]

(b) (i) 9.26 × 10–6 moles , forming 9.26 × 10–9 (mol dm–3) [2]

(ii) 0.02315 moles, 2.5 g [1]

(c) (i) Ag2+ [1]

(ii) C4H8SCl2 + 2NaOH → C4H8S(OH)2 + 2NaCl [1] 8

Section B 80

Total 100


2017


[SCH12]FRIDAY 26 MAY, MORNING


INSTRUCTIONS TO CANDIDATESWrite your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all fifteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all five questions in Section B. You must answer the questions in the spaces provided.Do not write outside the boxed area on each page or on blank pages.Complete in black ink only. Do not write with a gel pen.

INFORMATION FOR CANDIDATESThe total mark for this paper is 90.Quality of written communication will be assessed in Question 13(c).In Section A all questions carry equal marks, i.e. one mark for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included with this question paper.

Centre Number

Candidate Number

10675

*SCH12*

*SCH12*

*20SCH1201*

*20SCH1201*

New

Specif

icatio

n

*20SCH1202*

*20SCH1202*

10675

Section A – Multiple Choice


Each multiple choice question is worth 1 mark.

1 Bromine is formed in the reaction below.

Cl2 + 2NaBr → 2NaCl + Br2

Which statement about the reaction is correct?

A Bromide ions lose electrons

B Bromine is reduced by chlorine

C Chloride ions are reduced

D Chlorine is a weaker oxidising agent than bromide

2 Which trend in the Periodic Table is correct?

A Boiling point decreases from fluorine to bromine

B First ionisation energy decreases from lithium to caesium

C First ionisation energy increases from nitrogen to oxygen

D Melting point decreases from sodium to silicon

3 Which of the following is the structure of 55Mn2+ ?

protons neutrons electrons

A 25 30 23

B 25 30 27

C 27 30 25

D 30 25 28

*20SCH1203*

*20SCH1203*

10675[Turn over

4 Potassium iodide is formed when potassium is warmed in iodine vapour. Which of the following shows the bonding in the three species?

potassium iodine potassium iodide

A ionic covalent ionic

B metallic ionic covalent

C covalent covalent ionic

D metallic covalent ionic

5 The element astatine lies below iodine in the Periodic Table and is likely to

A be black.

B be a volatile liquid at room temperature and pressure.

C form an astatide ion, At2-.

D oxidise iodide ions to iodine.

6 Which molecule is non-polar?

A H2S

B NH3

C PF3

D SF6

*20SCH1204*

*20SCH1204*

10675

7 The element boron has a relative atomic mass of 10.8. In this sample, boron exists as two isotopes, 10B and 11B. The percentage abundance of 10B in this sample of boron is

A 10.8%.

B 20.0%.

C 80.0%.

D 89.2%.

8 When burned in oxygen magnesium forms magnesium oxide.

2Mg + O2 → 2MgO

What is the number of molecules of oxygen required for the complete oxidation of 1.2 g of magnesium?

A 1.5 × 1022

B 3.0 × 1022

C 3.0 × 1023

D 6.0 × 1023

9 Which statement describes the trends in electronegativity values in the Periodic Table?

A Decrease across a Period and increase down a Group

B Decrease across a Period and decrease down a Group

C Increase across a Period and increase down a Group

D Increase across a Period and decrease down a Group

*20SCH1205*

*20SCH1205*

10675[Turn over

10 Which of the following would exactly neutralise 10.0 cm3 of 1.00 mol dm-3 NaOH(aq)?

A 2.50 cm3 of 1.00 mol dm-3 CH3COOH

B 5.00 cm3 of 1.00 mol dm-3 HCl

C 5.00 cm3 of 1.00 mol dm-3 H2SO4

D 3.00 cm3 of 1.00 mol dm-3 H3PO4

*20SCH1206*

*20SCH1206*

10675

Section B

Answer all five questions in the spaces provided.

11 Sulfate, hydrogensulfate and thiosulfate ions are formed when sulfuric and thiosulfuric acids ionise.

(a) (i) Write the equation for the complete ionisation of thiosulfuric acid.

[2]

(ii) Write the formula for the hydrogensulfate ion.

[1]

(b) (i) Write the formula for ammonium sulfate.

[1]

(ii) Describe the bonding in ammonium sulfate.

[2]

(c) Describe how you could use chemical tests on an aqueous solution of ammonium sulfate to prove that it contains ammonium ions and sulfate ions.

[4]

*20SCH1207*

*20SCH1207*

10675[Turn over

12 Some properties of the metals sodium and aluminium are shown in the table below.

metal charge on metal ion electronic structure of the atom

melting point /°C

sodium 1+ 1s2 2s2 2p6 3s1 98

aluminium 3+ 1s2 2s2 2p6 3s2 3p1 660

(a) Describe, without using a diagram, the bonding in sodium metal.

[2]

(b) Explain why aluminium has a higher melting point than sodium.

[2]

(c) (i) Write the equation, including state symbols, for the first ionisation energy of sodium.

[2]

(ii) The first six ionisation energies, in kJ mol-1, of sodium are 496, 4563, 6913, 9544, 13352 and 16611. Explain which of these values can be used to identify sodium as belonging to Group I of the Periodic Table.

[2]

*20SCH1208*

*20SCH1208*

10675

(iii) The outer electron in the sodium atom is located in the 3s orbital. Explain what is meant by the term orbital.

[2]

(d) Aluminium forms covalent bonds with chlorine.

(i) Explain what is meant by the term covalent bond.

[2]

(ii) Write the equation for the reaction of aluminium with chlorine to form aluminium chloride, AlCl3.

[1]

(iii) State the octet rule and explain whether the atoms in aluminium chloride obey the rule.

[3]

*20SCH1209*

*20SCH1209*

10675[Turn over

13 (a) Zinc reacts with chlorine to form the ionic compound zinc chloride. Draw a dot and cross diagram, using outer electrons only, to show how zinc chloride, ZnCl2, is formed from zinc and chlorine atoms.

[2]

(b) Zinc is an essential trace element. People who have a zinc deficiency can take hydrated zinc sulfate, ZnSO4.xH2O, as a dietary supplement.

The value of x can be determined by heating hydrated zinc sulfate to constant mass.

A student heated 5.65 g of hydrated zinc sulfate and obtained 3.85 g of anhydrous zinc sulfate.

(i) Calculate the number of moles of anhydrous zinc sulfate obtained.

[1]

(ii) Calculate the mass of water present in the hydrated zinc sulfate.

[1]

(iii) Calculate the number of moles of water present in the hydrated zinc sulfate.

[1]

(iv) Calculate the value of x in ZnSO4.xH2O

[1]

*20SCH1210*

*20SCH1210*

10675

(c) Describe how you would prepare 250.0 cm3 of a 28.7 g dm-3 zinc sulfate solution from the anhydrous solid.

In this question you will be assessed on using your written communication skills including the use of specialist scientific terms.

[6]

*20SCH1211*

*20SCH1211*

10675[Turn over

14 Nitrogen and phosphorus are Group V elements. They form the toxic hydrides ammonia and phosphine.

(a) Ammonia is formed by the reversible reaction of nitrogen with hydrogen. Write the equation for this reaction.

[2]

(b) Phosphine is formed by the reaction of phosphorus with aqueous sodium hydroxide.

(i) Balance the equation for the formation of phosphine.

P4 + NaOH + H2O → NaH2PO2 + PH3 [1]

(ii) Deduce the oxidation number of phosphorus in:

P4

NaH2PO2

PH3 [3]

(iii) Explain, using the oxidation numbers of phosphorus, why the reaction is described as disproportionation.

[3]

*20SCH1212*

*20SCH1212*

10675

(c) The boiling point of ammonia is -33 °C while that of phosphine is -88 °C. Explain why the boiling point of ammonia is higher than that of phosphine.

[3]

(d) Both ammonia and phosphine molecules react with H+ ions.

PH3 + H+ → PH4+

(i) Name the type of bond formed between a phosphine molecule and the H+ ion.

[1]

(ii) Draw and name the shapes of the molecule PH3 and the ion PH4+.

PH3

Shape _________________

PH4+

Shape _________________

[4]

(iii) Explain why the bond angle in PH3 is different from the bond angle in PH4+.

[3]

*20SCH1213*

*20SCH1213*

10675[Turn over

(e) Ammonia is very soluble in water. Draw diagrams to show the two ways in which a molecule of ammonia can be attracted to a molecule of water. Include all partial charges and lone pairs in your diagram.

[4]

*20SCH1214*

*20SCH1214*

10675

15 Ammonia is used to make nitric acid by the Ostwald Process outlined below.

Reaction 1: 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g)

Reaction 2: 2NO(g) + O2(g) → 2NO2(g)

Reaction 3: 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g)

(a) (i) Calculate the number of moles of oxygen needed to react with 6.8 kg of ammonia.

[3]

(ii) Calculate the number of moles of nitrogen(IV) oxide which can be obtained from 6.8 kg of ammonia.

[2]

(iii) Calculate the concentration of nitric acid, in g dm-3, produced on reacting the nitrogen(IV) oxide obtained in part (ii) with 50 dm3 of water.

[3]

*20SCH1215*

*20SCH1215*

10675[Turn over

(b) Ammonia reacts with nitric acid according to the equation below.

NH3 + HNO3 → NH4NO3

The following results were obtained by diluting 25.0 cm3 of a concentrated ammonia solution to 250.0 cm3 in a volumetric flask and then titrating 25.0 cm3 portions of the diluted ammonia solution using 0.100 mol dm-3 nitric acid.

titration initial burette reading /cm3

final burette reading /cm3 titre /cm3

rough 0.00 22.00 22.00

first accurate 0.10 21.40 21.30

second accurate 0.20 21.60 21.40

(i) Name a suitable indicator for the titration and state the colour change at the end point.

[3]

(ii) Calculate the mean titre.

[1]

*20SCH1216*

*20SCH1216*

10675

(iii) A burette has an uncertainty of ±0.05 cm3. Calculate the uncertainty when two burette readings are used to calculate a titre value.

[1]

(iv) Calculate the concentration of the concentrated ammonia solution in mol dm-3.

[5]

*20SCH1217*

*20SCH1217*

10675

BLANK PAGE



*20SCH1218*

*20SCH1218*

10675

BLANK PAGE


*20SCH1219*

*20SCH1219*

10675

BLANK PAGE


*20SCH1220*

*20SCH1220*


221023


QuestionNumber Marks

Section A1–10

Section B1112131415

TotalMarks


CHEMISTRY DATA SHEETGCE A/AS EXAMINATIONS CHEMISTRY

Including the Periodic Table of the Elements

For the use of candidates takingAdvanced Subsidiary and Advanced LevelChemistry Examinations

Copies must be free from notes or additions of any kind.

No other type of data booklet or information sheet is

authorised for use in the examinations.

For first teaching from September 2016For first award of AS Level in Summer 2017For first award of A Level in Summer 2018Subject Code: 1110

GCE

NewSpe

cifi ca

tionGeneral Information

1 tonne = 106 g1 metre = 109 nmOne mole of any gas at 293 K and a pressure of 1 atmosphere (105 Pa) occupies a volume of 24 dm3

Avogadro Constant = 6.02 1023 mol–1

Planck Constant = 6.63 10–34 J sSpecifi c Heat Capacity of water = 4.2 J g–1 K–1

Speed of Light = 3 108 m s–1

Characteristic absorptions in IR spectroscopy

Wavenumber/cm–1 Bond Compound550–850 C–X (X = Cl, Br, I) Haloalkanes750–1100 C–C Alkanes, alkyl groups1000–1300 C–O Alcohols, esters, carboxylic acids 1450–1650 C ̿ C Arenes1600–1700 C ̿ C Alkenes1650–1800 C ̿ O Carboxylic acids, esters, aldehydes, ketones, amides, acyl chlorides2200–2300 C N Nitriles2500–3200 O–H Carboxylic acids2750–2850 C–H Aldehydes2850–3000 C–H Alkanes, alkyl groups, alkenes, arenes3200–3600 O–H Alcohols3300–3500 N–H Amines, amides

Proton Chemical Shifts in Nuclear Magnetic Resonance Spectroscopy (relative to TMS)

Chemical Shi Structure

0.5–2.0 –CH Saturated alkanes0.5–5.5 –OH Alcohols1.0–3.0 –NH Amines2.0–3.0 –CO–CH Ketones –N–CH Amines C6H5–CH Arene (alipha c on ring)2.0–4.0 X–CH X = Cl or Br (3.0–4.0) X = I (2.0–3.0)4.5–6.0 –C ̿ CH Alkenes5.5–8.5 RCONH Amides6.0–8.0 –C6H5 Arenes (on ring)9.0–10.0 –CHO Aldehydes10.0–12.0 –COOH Carboxylic acids

These chemical shi s are concentra on and temperature dependent and may be outside the ranges indicated above.

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* 58–71 Lanthanum series† 90–103 Actinium series

a = relative atomic mass (approx)x = atomic symbolb = atomic number

THE PERIODIC TABLE OF ELEMENTSGroup

ab

x

*

†

1

1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

III III IV VI VII 0V

285

112Copernicium

MARKSCHEME

10675.01 F


2017


assessing


[SCH12]

FRIDAY 26 MAY, MORNING

New

Specifi

catio

n

10675.01 F 2 [Turn over








10675.01 F 3

AVAILABLEMARKS

Section A

1 A

2 B

3 A

4 D

5 A

6 D

7 B

8 A

9 D


Section A 10

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Section B

11 (a) (i) H2S2O3 + 2H2O → 2H3O+ + S2O32–

or

H2S2O3 → 2H+ + S2O32– [2]

(ii) HSO4

– [1] (b) (i) (NH4)2SO4 [1] (ii) oppositely charged ions are held together by electrostatic forces of attraction error [–1] [2]

(c) NH4+: warm with (dilute) sodium hydroxide solution [1]

(the gas produced) gives white fumes with (a glass rod dipped in) concentrated HCl [1] SO4

2–: add a solution of barium chloride [1] a white precipitate forms [1] [4] 10

12 (a) (attraction) between positive sodium ions [1] and delocalised electrons [1] [2]

(b) increase in charge of metal ion/aluminium ion has 3+ compared to 1+ in sodium [1] increase in the number of electrons in delocalised cloud/sea [1] [2]

(c) (i) Na(g) → Na+(g) + e– [2]

(ii) 1st and 2nd values [1] large gap between [1] [2] (iii) A region within an atom that can hold up to two electrons [1] with opposite spin [1]. [2]

(d) (i) (Electrostatic) attraction between a shared pair of electrons and the nuclei of bonded atoms. [2]

(ii) 2Al + 3Cl2 → 2AlCl3 [1] (iii) When forming a compound, an atom tends to gain, lose or share electrons [1] to achieve eight in its outer shell. [1] Aluminium has six electrons in the outer shell [1]

[3] 16 Chlorine has eight

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13 (a)

Zn +

Cl

Cl

→ Zn2+

Cl

–

Cl

–

Outer shell of zinc could show ten electrons [2]

(b) (i) 3.85 ÷ 161 = 0.024/0.0239 [1]

(ii) 5.65 – 3.85 = 1.80g [1]

(iii) 1.80 ÷ 18 = 0.1/0.10 [1]

(iv) 0.0239 : 0.10 1 : 4.18/4.2/4.0/4 X = 4.18/4.2/4.0/4 [1]

(c) Indicative content • (Weigh out) 7.18g of zinc sulfate * • in a beaker/suitable container • Dissolve the solid in a small amount (50–100cm3) of distilled/deionised water • Transfer the solution/with washings • to the 250.0 cm3 volumetric flask * • and make up to the mark • Stopper and invert the flask

* essential for [6] in Band A If either * missing max [5]

Band Response Mark

A

Candidates must use appropriate specialist terms to fully explain the preparation of the standard solution using 6 points of indicative content. They must use good spelling, punctuation and grammar and the form and style are of an excellent standard.

[5]–[6]

B

Candidates must use appropriate specialist terms to explain the preparation of the standard solution using a minimum of 4 points of indicative content. They must use satisfactory spelling, punctuation and grammar and the form and style are of a good standard.

[3]–[4]

C

Candidates must partially explain the preparation of the standard solution using a minimum of 2 points of indicative content. They use limited correct spelling, punctuation and grammar and the form and style are of a basic standard.

[1]–[2]

D Response not worthy of credit. [0]

[6] 12

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14 (a) N2 + 3H2 2NH3 [2] ([–1] each error)

(b) (i) P4 + 3NaOH + 3H2O → 3NaH2PO2 + PH3 [1] (ii) P4 = 0 [1] NaH2PO2 = +1 [1] PH3 = –3 [1] [3]

(iii) 0 to +1 is oxidation; 0 to –3 is reduction [1] oxidation and reduction of the same element [1] in the same reaction [1] [3]

(c) Between molecules of NH3 there are van der Waals’ forces and H-bonds [1] Between molecules of PH3 there are only van der Waals’ forces [1] Hydrogen bonds are stronger and require more energy to break. [1] [3]

(d) (i) dative covalent/co-ordinate bond [1]

(ii) H

P

HH H

+

P

HH H

••

[2]

pyramidal tetrahedral [2]

(iii) PH3 has one lone pair/3 bond pairs [1] greater repulsion between lone pair–bond pair [1] pushes bond pairs closer together/reduces bond angle [1] [3] (e)

O O

N••H+

H+

H+

–H+

H+

•• ••• •

••

–

N••–

H+

H+

H+

H+

H+

–and

[2] [2] [4] 24

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15 (a) (i) 6.8 × 1000 = 6800g 6800 ÷ 17 = 400 NH3 : O2 4 : 5 Moles of O2 = 500 ([–1] each error) [3]

(ii) NO : NO2 1 : 1 Moles of NO2 = 400 [2]

(iii) 400 × (32 ) = 266.6667

Concentration = 266.667 ÷ 50 = 5.33 mol dm–3

5.33 × 63 = 336 g dm–3

([–1] each error) [3] (b) (i) Methyl orange [1] Yellow [1] to red [1] [3]

(ii) 2

21.30 21.40+] g = 21.35 cm3 [1] (iii) 0.05 + 0.05 = 0.1 cm3 [1]

(iv) 1000

21.35 0.100#^ h = 2.135 × 10–3

NH3 : HNO3 1 : 1 moles of diluted ammonia in 25 cm3 = 0.002135 moles of diluted ammonia in 250 cm3 = 0.02135 concentration of diluted ammonia = 0.02135 ÷ 0.25 = 0.0854 concentration of undiluted ammonia = 0.0854 × 10 = 0.854 mol dm–3

([–1] each error) [5] 18

Section B 80

Total 90


2018


[SCH12]TUESDAY 22 MAY, MORNING


INSTRUCTIONS TO CANDIDATESWrite your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all fourteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all four questions in Section B. You must answer the questions in the spaces provided.Do not write outside the boxed area on each page or on blank pages.Complete in black ink only. Do not write with a gel pen.

INFORMATION FOR CANDIDATESThe total mark for this paper is 90.Quality of written communication will be assessed in Question 13(a).In Section A all questions carry equal marks, i.e. one mark for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included with this question paper.

*SCH12*

*SCH12*

*20SCH1201*

*20SCH1201*

Centre Number

Candidate Number

11282

*20SCH1202*

*20SCH1202*

11282

Section A



1 A solution of barium chloride was added to sodium sulfate solution.

heat

P Q R S

Which combination of methods should be used to obtain the precipitate and the other product as a solid?

A P + Q

B P + R

C Q + S

D R + S

2 Which species has the same electronic arrangement as a lithium ion, Li+?

A Be−

B B2+

C H+

D He

*20SCH1203*

*20SCH1203*

11282[Turn over

3 Sodium azide decomposes, in an airbag, to form sodium and nitrogen.

2NaN3 → 2Na + 3N2

The sodium then reacts with potassium nitrate to form more nitrogen gas.

10Na + 2KNO3 → 5Na2O + N2 + K2O

0.50 mol of sodium azide produces

A 0.50 mol of nitrogen.

B 0.75 mol of nitrogen.

C 0.80 mol of nitrogen.

D 2.00 mol of nitrogen.

4 Chlorine has two isotopes. How many peaks are there in the mass spectrum of chlorine?

A 2

B 3

C 4

D 5

5 Which molecule is not planar?

A BF3

B BeCl2

C HCHO

D NCI3

*20SCH1204*

*20SCH1204*

11282

6 The curves shown below are for 25 cm3 of acids HX and HY when they are reacted with 0.1 M sodium hydroxide solution.

pH

20

HX

HY

30

volume of 0.1 M NaOH added/cm3

3

1

Compared to acid HY, the acid HX is

A more concentrated and stronger.

B more concentrated and weaker.

C less concentrated and stronger.

D less concentrated and weaker.

7 The largest mass of silver chloride precipitated is when excess silver ions are added to

A 25.0 cm3 of 0.80 M hydrochloric acid. B 30.0 cm3 of 0.30 M iron(III) chloride solution.

C 50.0 cm3 of 0.20 M magnesium chloride solution. D 50.0 cm3 of 0.50 M sodium chloride solution.

*20SCH1205*

*20SCH1205*

11282[Turn over

8 On melting, covalent bonds are broken in

A bromine.

B diamond.

C sodium chloride.

D sulfur(IV) oxide.

9 Which of the following equations represents a redox reaction?

A CaCO3 + SiO2 → CaSiO3 + CO2

B 3Cl2 + 6OH− → 5Cl− + ClO3− + 3H2O

C 2CrO42− + 2H+ → Cr2O7

2− + H2O

D HNO3 + 2H2SO4 → NO2+ + H3O+ + 2HSO4

−

10 Which halide has the most covalent character?

A AIBr3

B AIF3

C MgBr2

D MgF2

*20SCH1206*

*20SCH1206*

11282

Section B

Answer all four questions in this section

11 Chlorine monoxide is a brown-yellow gas with a boiling point of 4 °C while chlorine has a boiling point of −34 °C. The monoxide is formed when excess chlorine is passed over mercury(II) oxide.

2Cl2(g) + HgO(s) → HgCl2(s) + Cl2O(g)

Cl2

HgO

The escaping gases are passed through a U-tube which is cooled to −30 °C. The chlorine monoxide condenses in the U-tube.

(a) (i) How could you test to show that chlorine is passing into the reaction tube?

[2]

(ii) What is the colour of chlorine?

[1]

(iii) Why is it important to limit the temperature of the U-tube to −30 °C and not to have it lower than this temperature?

[1]

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*20SCH1207*

11282[Turn over

(iv) How could you show that it is a chloride which remains in the reaction tube?

[3]

(v) Mercury(ll) oxide decomposes when heated to form oxygen and mercury. How could you show that there was no mercury(ll) oxide left in the reaction tube at the end of the experiment?

[3]

(vi) Chlorine monoxide cannot be collected over water as it is very soluble in water, with a solubility of 143 g in 100 cm3 at room temperature and pressure. Explain how you could show that chlorine monoxide is very soluble in water.

[4]

*20SCH1208*

*20SCH1208*

11282

(b) Chlorine monoxide slowly reacts with water to form hypochlorous acid

Cl2O + H2O → 2HOCI

(i) Hypochlorous acid has a systematic name based on chloric acid. State the systematic name for hypochlorous acid.

[1]

(ii) Hypochlorous acid is a weak acid. Explain what is meant by the term weak acid.

[2]

(iii) Hypochlorous acid decomposes to give hydrochloric acid and oxygen. Write the equation for this reaction.

[2]

(c) Chlorine monoxide obeys the octet rule.

(i) State the octet rule.

[2]

(ii) Draw the electronic structure of chlorine monoxide showing the outer electrons only.

[2]

*20SCH1209*

*20SCH1209*

11282[Turn over

(d) Chlorine monoxide has the following shape.

O

Cl Cl

(i) Name the shape.

[1]

(ii) How many lone pairs are there in a chlorine monoxide molecule?

[1]

(iii) Explain why chlorine monoxide forms this shape.

[2]

(e) Fluorine also forms an oxide but this oxide is known as oxygen fluoride because fluorine has a greater electronegativity than oxygen. State how electronegativity changes across a Period and down a Group.

[2]

(f) It has been suggested that chlorine monoxide is the active ingredient in the treatment of water for drinking purposes. Name two substances that are used to treat water for drinking.

[2]

*20SCH1210*

*20SCH1210*

11282

12 Ammonium dichromate is used in the “volcano” experiment. When heated, it decomposes to produce a vast amount of green chromium oxide and gases which push out the green “ash” to form a pile of “lava”.

(NH4)2Cr2O7 → Cr2O3 + N2 + 4H2O

The water forms steam because of the heat of the reaction.

(a) Write the equation for the reaction, with state symbols, for the reactants and products.

[1]

(b) Ammonium dichromate is very soluble in water. At room temperature 10.0 g of ammonium dichromate dissolve in 25.0 cm3 of water. The orange solution can be tested for the presence of ammonium ions.

(i) Calculate the solubility of the ammonium dichromate in g dm−3 to 3 significant figures.

[1]

(ii) Calculate the solubility of the ammonium dichromate in mol dm−3 to 3 significant figures.

[1]

(iii) Explain how you would show that the orange solution contains ammonium ions.

[3]

*20SCH1211*

*20SCH1211*

11282[Turn over

(c) The nitrogen given off in the reaction consists of two isotopes, nitrogen-14 and nitrogen-15. The percentage abundance of nitrogen-14 is 99.632%.

(i) Explain what is meant by the term isotopes.

[2]

(ii) Calculate the percentage abundance of nitrogen-15 given off.

[1]

(iii) Calculate the relative atomic mass of nitrogen to three decimal places.

[2]

(iv) Explain why there is a difference between the calculated relative atomic mass and the one provided in the data sheet.

[1]

.

*20SCH1212*

*20SCH1212*

11282

(d) The dichromate ion is a very strong oxidising agent. The half-equation which shows its oxidising ability is:

Cr2O72− + 14H+ + 6e− → 2Cr3+ + 7H2O

(i) Use this equation to explain, in terms of oxidation numbers, why the dichromate ion is an oxidising agent.

[2]

(ii) Use this equation to explain, in terms of electrons, why the dichromate ion is an oxidising agent.

[1]

(e) Dichromates react with chlorides in the presence of concentrated sulfuric acid to produce chromyl chloride, CrO2Cl2, which is a deep red liquid with a boiling point of 117 °C. Using this information, explain whether chromyl chloride is ionic or covalent.

[2]

*20SCH1214*

*20SCH1214*

11282

13 (a) There are several types of structure which apply to chemical formulae. The species present may be atoms, molecules or ions. In each of the following examples describe which type of structure it is and which type of species is present.

sodium chloride

diamond

bromine

In this question you will be assessed on using your written communication skills including the use of specialist scientific terms.

[6]

(b) The different types of structure have different physical properties. State four physical properties that depend upon structure.

[3]

*20SCH1215*

*20SCH1215*

11282[Turn over

(c) The structure of sodium is shown below.

+ + +

+ + +

+ + +

(i) Attach words to the labels shown. [2]

(ii) Use this diagram to explain whether magnesium has a greater or lower conductivity than sodium.

[2]

(iii) Explain, using a labelled diagram, how you could compare the electrical conductivities of sodium and magnesium in the laboratory.

[3]

*20SCH1216*

*20SCH1216*

11282

14 Ethene is a gas at room temperature and has a boiling point of −104 °C at atmospheric pressure. It has a relative molecular mass of 28 which is approximately the same as the average relative molecular mass of air. It is a planar molecule which has the following structure:

H H C C

H H

(a) The ethene molecule contains single bonds and a double bond which are formed from s- and p-orbitals.

(i) Draw the shape of an s-orbital.

[1]

(ii) Draw the shape of a p-orbital.

[1]

(iii) Explain what is meant by the term orbital.

[2]

*20SCH1217*

*20SCH1217*

11282[Turn over

(b) Ethene is a non-polar molecule. There are two reasons why ethene can be considered to be non-polar. One is based on electronegativity and the other is based on shape.

(i) What is meant by the term electronegativity?

[2]

(ii) Explain why ethene is considered non-polar based on electronegativity.

[1]

(iii) Explain why ethene is considered non-polar based on shape.

[1]

(c) Ethene contains a double bond. Other molecules can contain triple bonds.

(i) Draw the structure of the hydrocarbon ethyne which contains two carbon atoms and a triple bond.

[1]

(ii) Name an element which contains a triple bond.

[1]

*20SCH1218*

*20SCH1218*

11282

(d) When ethene burns, carbon dioxide and water are produced. Describe how you would carry out a test for carbon dioxide and the result expected for a positive test.

[2]

(e) Gases can be collected by two different methods A or B depending on their relative molecular masses compared to air.

gas lighter than air gas heavier than air

A B

(i) Explain which method could be used to collect methane, CH4.

[1]

(ii) Explain which method could be used to collect chlorine.

[1]

*20SCH1219*

*20SCH1219*

11282

(f) The boiling point of methane is −161 °C. Explain why the boiling point of methane is lower than that of ethene.

[2]


MARKSCHEME

11282.01 F


2018




[SCH12]

TUESDAY 22 MAY, MORNING

11282.01 F 2 [Turn over








11282.01 F 3

AVAILABLEMARKS

Section A

1 D

2 D

3 C

4 D

5 D

6 D

7 B

8 B

9 B


Section A 10

11282.01 F 4 [Turn over

AVAILABLEMARKS

Section B

11 (a) (i) Damp indicator (paper) [1] is bleached [1] [2]

(ii) Chlorine is green/green-yellow/yellow-green [1] [1]

(iii) Chlorine would condense [1]

(iv) Dissolve the chloride in water [1] add silver nitrate solution [1] white precipitate formed [1] [3]

(v) Heat the solid left in the tube in a test tube [1] test with a glowing splint [1] it does not relight [1] [3]

(vi) Either Fill a test tube with chlorine monoxide [1] seal the end of the test tube [1] invert in a beaker of water [1] open the test tube to see how far the water rises [1]

Or Bubble through H2O [1] stated volume of H2O/until bubbles appear [1] measured using a physical property, e.g. colour, pH, density, mass, volume, forming precipitate with AgNO3(aq) [2] [4]

(b) (i) Chloric(I) acid [1]

(ii) Partial dissociation in solution (to form hydrogen ions) [2]

(iii) 2HOCI → 2HCI + O2 [2]

(c) (i) When reacting an atom tends to gain, lose or share electrons to achieve 8 in its outer shell [2]

(ii) ×CI ×O××

××CI

[2]

(d) (i) Bent [1]

(ii) 8 [1]

(iii) The lone pairs repel the bonded pairs more [2]

(e) Electronegativity increases across a period [1]

Electronegativity decreases down a group [1] [2]

(f) Chlorine [1] ozone [1] [2] 31

11282.01 F 5

AVAILABLEMARKS

12 (a) (s) (s) (g) (g) [1]

(b) (i) 10.0 g in 25.0 cm3; 40 × 10.0 g = 400 g in 1 dm3 [1]

(ii) (NH4)2Cr2O7 = 2 × 18 + 2 × 52 + 7 × 16 =

36 + 104 + 112 = 252 400.0/252 = 1.59 [1]

(iii) Heat with sodium hydroxide solution [1] moist indicator paper/litmus/red litmus/UI paper [1] turns blue [1] [3]

(c) (i) Atoms which have the same atomic number but a different mass number or contain the same number of protons but a different number of neutrons [2]

(ii) 100 – 99.632 = 0.368 [1]

(iii) 99.632 × 14 = 1394.848 0.368 × 15 = 5.52 = 1400.368 = 14.004 [2]

(iv) The RAMs in the table are listed as whole numbers (exception of chlorine) [1]

(d) (i) Oxidation number in dichromate is +6; oxidation number in Cr3+ is +3; the oxidation number goes down when an oxidant reacts [2]

(ii) Oxidising agents gain electrons (which are supplied by the reducing agent) [1]

(e) Low boiling point hence covalent [1]; liquid at room temperature hence covalent [1] [2] 17

11282.01 F 6 [Turn over

AVAILABLEMARKS

AVAILABLEMARKS

13 (a) Sodium chloride has ions, it has a giant ionic lattice structure [2] Diamond has atoms, it has a giant covalent structure [2] Bromine has molecules in it, it has a molecular covalent structure [2] [6]

Band Response Mark

A Candidates use 5 or more indicative points above. They use appropriate specialist terms and the spelling, punctuation and grammar and form and style are of a good standard.

[5]–[6]

B Candidates use 3–4 indicative points above. They use appropriate specialist terms and the spelling, punctuation and grammar and form and style are of a satisfactory standard.

[3]–[4]

C Candidates make reference to 1–2 indicative points above using limited spelling, punctuation and grammar and the form and style are of limited standard and they have made no use of specialist terms.

[1]–[2]

D Not worthy of credit. [0]

(b) Melting point; boiling point; hardness; electrical conductivity [3]

(c) (i) Delocalised electrons [1]; positive ions [1] [2]

(ii) Magnesium has a greater conductivity [1]; there are more delocalised electrons [1] or magnesium produces 2 delocalised electrons compared to 1 with sodium [2]

(iii) A

cell

metal

ammeter (or light bulb)

Labels [2] Compare reading on ammeter (brightness of light bulb) [1] [3] 16

11282.01 F 7

AVAILABLEMARKS

14 (a) (i) [1]

(ii) [1]

(iii) A region within an atom that can hold up to two electrons with opposite spins [2]

(b) (i) The extent to which an atom attracts the bonding electrons in a covalent bond [2]

(ii) There is little difference in the electronegativities of carbon and hydrogen [1]

(iii) The molecule is symmetrical and the polarities cancel out [1] (c) (i) H C C H [1]

(ii) nitrogen [1]

(d) Bubble into limewater [1] goes milky [1] [2]

(e) (i) A CH4 = 16 hence lighter than air [1]

(ii) B Cl2 = 71 hence heavier than air [1]

(f) Lower mass [1] less/weaker van der Waals’ forces [1] [2] 16

Section B 80

Total 90


2019


[SCH12]MONDAY 20 MAY, MORNING


INSTRUCTIONS TO CANDIDATESWrite your Centre Number and Candidate Number in the spaces provided at the top of this page.Answer all sixteen questions.Answer all ten questions in Section A. Record your answers by marking the appropriate letter on the answer sheet provided. Use only the spaces numbered 1 to 10. Keep in sequence when answering.Answer all six questions in Section B. You must answer the questions in the spaces provided.Do not write outside the boxed area on each page or on blank pages.Complete in black ink only. Do not write with a gel pen.

INFORMATION FOR CANDIDATESThe total mark for this paper is 90.Quality of written communication will be assessed in Question 15(c).In Section A all questions carry equal marks, i.e. one mark for each question.In Section B the figures in brackets printed down the right-hand side of pages indicate the marks awarded to each question or part question.A Periodic Table of Elements, containing some data, is included with this question paper.

Centre Number

Candidate Number

11852

*SCH12*

*SCH12*

*20SCH1201*

*20SCH1201*

*20SCH1202*

*20SCH1202*

11852

Section A – Multiple Choice


Each multiple choice question is worth 1 mark.

1 In which of the following does chromium not have an oxidation state of +6?

A CrO3

B CrO42−

C Cr2O72−

D Cr2O3

2 Which bonding type is described as intermolecular?

A Covalent

B Ionic

C Metallic

D van der Waals’ forces

3 Which of the following is the formula of the nitrite ion?

A N3−

B NH4+

C NO2−

D NO3−

*20SCH1203*

*20SCH1203*

11852[Turn over

4 25.0 cm3 of 0.10 M sodium hydroxide solution is exactly neutralised by

A 12.5 cm3 of 0.05 M sulfuric acid.

B 25.0 cm3 of 0.05 M sulfuric acid.

C 12.5 cm3 of 0.20 M sulfuric acid.

D 25.0 cm3 of 0.10 M sulfuric acid.

5 The electronic configuration of a Group III element is

A 1s2 2s2 2p6 3s2 3p6 3d3 4s2.

B 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1.

C 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2.

D 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p3.

6 The bond angle in ammonia is

A 104.5°.

B 107°.

C 109.5°.

D 120°.

*20SCH1204*

*20SCH1204*

11852

7 The sulfate(VI) ion can be reduced to sulfur dioxide.

SO42− + xH+ + ye− → SO2 + zH2O

Which of the following represents the correct values of x, y and z?

x y z

A 2 2 4

B 2 4 2

C 4 2 2

D 4 4 2

8 Which type of titration can use phenolphthalein as a suitable indicator?

A Strong acid/strong base only

B Strong acid/strong base and weak acid/strong base

C Strong acid/strong base and strong acid/weak base

D Strong acid/weak base and weak acid/strong base

9 In which of the following molecules does the central atom obey the octet rule?

A BF3

B BeCl2

C ClF3

D PH3

*20SCH1205*

*20SCH1205*

11852[Turn over

10 Which species is the most powerful oxidising agent?

A Bromide

B Bromine

C Chloride

D Chlorine

*20SCH1206*

*20SCH1206*

11852

Section B


11 Water can act as an acid or as a base. It can either lose or gain hydrogen ions.

(a) State and explain the shape of a water molecule.

[3]

(b) Water can react with hydrogen ions forming hydroxonium ions, H3O+.

(i) Draw a dot and cross diagram to show the bonding in a hydroxonium ion, showing all the outer shell electrons.

[2]

*20SCH1207*

*20SCH1207*

11852[Turn over

(ii) Suggest why the bond angle in the hydroxonium ion is greater than the bond angle in water.

[2]

(iii) Suggest why the hydroxonium ion does not react with a hydrogen ion to form H4O2+.

[1]

*20SCH1208*

*20SCH1208*

11852

12 Four new elements have recently been added to the Periodic Table. The four elements are given below, along with their atomic numbers and the mass numbers of their most common isotope.

element atomic number mass number

nihonium 113 286

moscovium 115 289

tennessine 117 294

oganesson 118 294

(a) What is the meaning of the following terms?

(i) Atomic number

[1]

(ii) Mass number

[1]

(iii) Isotopes

[1]

(b) State and explain which element has the most neutrons.

[2]

*20SCH1209*

*20SCH1209*

11852[Turn over

(c) Suggest why tennessine is placed in Group VII of the Periodic Table.

[1]

(d) Erbium is a soft, silvery solid that tarnishes slowly in air. It is used in fibre optic cables. There are six known isotopes of erbium and its relative atomic mass is 167.26.

(i) Define the term relative isotopic mass.

[2]

(ii) The table below gives the percentage abundances of six isotopes in the mass spectrum of erbium.

relative isotopic mass 161.93 163.93 165.93 167.93 169.94

% abundance 0.14 1.60 33.50 22.87 26.98 14.91

Calculate the missing relative isotopic mass.

[3]

*20SCH1210*

*20SCH1210*

11852

13 Chloroauric acid, HAuCl4, is an orange solid that is used widely in gold refining. During World War II, the Hungarian chemist George de Hevesy dissolved two gold Nobel Prize medals in a mixture of concentrated nitric and hydrochloric acids to prevent the Germans from confiscating them. Later the medals were reconstructed from the dissolved chloroauric acid and returned.

(a) The reaction between gold, concentrated hydrochloric acid and concentrated nitric acid produces chloroauric acid, nitrogen(IV) oxide and water. Write the equation for this reaction.

[2]

(b) Gold is extracted from recycled electronic materials by reaction with chlorine and hydrochloric acid, forming chloroauric acid. Elemental gold is recovered by electrolysis of chloroauric acid.

2Au + 3Cl2 + 2HCl → 2HAuCl4

(i) Deduce the oxidation state of gold in chloroauric acid.

[1]

(ii) With reference to oxidation numbers, explain why this is a redox reaction.

[3]

*20SCH1211*

*20SCH1211*

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(c) When heated, chloroauric acid forms gold(III) chloride and hydrogen chloride gas. The bonding in gold(III) chloride is considered to be covalent.

(i) Suggest, in terms of electronegativity, why the bonding in gold(III) chloride is covalent.

[1]

(ii) Describe the chemical test for hydrogen chloride gas.

[2]

*20SCH1212*

*20SCH1212*

11852

14 The recommended daily allowance for salt, sodium chloride, is 6.0 g. Eating too much salt can lead to high blood pressure, potentially causing heart disease and strokes.

(a) State the electronic configuration of a sodium atom and use it to explain why sodium is regarded as an s-block element.

[2]

(b) (i) Define the term Avogadro’s constant.

[1]

(ii) Calculate the number of sodium ions in the recommended daily allowance of sodium chloride.

[2]

(c) A solid sample of salt was analysed to confirm the identity of the ions present. A flame test was first conducted on the sample using nichrome wire and concentrated hydrochloric acid to identify sodium ions. The presence of chloride ions was subsequently confirmed.

(i) State two reasons why nichrome wire was used.

[2]

*20SCH1213*

*20SCH1213*

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(ii) State two reasons why concentrated hydrochloric acid was used.

[2]

(iii) State the colour observed in the flame test.

[1]

(d) Describe how the presence of chloride ions could be confirmed in the solid salt.

[4]

(e) A second salt sample was thought to be contaminated with sodium carbonate. Describe a chemical test to confirm the presence of carbonate ions.

[3]

*20SCH1214*

*20SCH1214*

11852

(f) The salt sample, of mass 6.0 g, contaminated with sodium carbonate was dissolved in water. A solution of magnesium chloride was added, forming a precipitate of magnesium carbonate. The precipitate was filtered off and dried to give 1.4 g of magnesium carbonate.

(i) Draw a dot and cross diagram to show the bonding in magnesium chloride showing all the outer electrons.

[2]

(ii) Write the equation for the reaction between sodium carbonate and magnesium chloride.

[2]

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(iii) Use the following headings to calculate the percentage of sodium carbonate in the salt sample.

Relative formula mass of magnesium carbonate

Number of moles of magnesium carbonate

Number of moles of sodium carbonate

Relative formula mass of sodium carbonate

Mass of sodium carbonate in the sample

Percentage of sodium carbonate in the sample

[6]

*20SCH1216*

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11852

15 The third Period in the Periodic Table from sodium to argon displays a number of periodic trends.

(a) State and explain the general trend in first ionisation energy across Period three.

[3]

(b) (i) Write an equation, including state symbols, for the first ionisation energy of phosphorus.

[2]

(ii) Explain why the first ionisation energy of phosphorus is higher than that of sulfur.

[2]

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(c) The graph below shows the melting points of the elements in the third Period.


1800

1600

1400

1200

1000

800

600

400

200

0

melting point/K

With reference to the structure and bonding of the elements, explain the change in melting point from silicon to argon.

In this question you will be assessed on your written communication skills including the use of specialist scientific terms.

[6]

*20SCH1218*

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11852

16 Lead(II) iodide is yellow and was once used as a pigment in paint until concerns over its toxicity led to its use being discontinued. It has a low solubility in water.

(a) Lead(II) iodide can be prepared by reaction between solutions of potassium iodide and lead(II) nitrate. Write the equation for this reaction.

[2]

(b) 75.6 mg of lead(II) iodide dissolve in 100 cm3 of water at 20°C. Calculate the molarity of iodide ions in a saturated solution of lead(II) iodide at 20°C.

[4]

(c) Chlorine water was added to potassium iodide solution in a test tube.

(i) State the colour observed.

[1]

(ii) A solution of starch was then added to the test tube. State the colour observed.

[1]

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*20SCH1219*

*20SCH1219*

(d) (i) State three observations made when concentrated sulfuric acid is added to solid potassium iodide.

[3]

(ii) Explain why concentrated phosphoric acid does not give iodine when added to solid potassium iodide.

[1]


*20SCH1220*

*20SCH1220*


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MARKSCHEME

11852.01


2019




[SCH12]

MONDAY 20 MAY, MORNING

Standardising Meeting Version

Not to be circulated beyond the Examining Team

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IntroductionMark schemes are published to assist teachers and students in their preparation for examinations. Through the mark schemes teachers and students will be able to see what the examiners are looking for in response to questions and exactly where the marks have been awarded. The publishing of the mark schemes may help to show that examiners are not concerned about finding out what a student does not know but rather with rewarding students for what they do know.


The job of the examiners is to set the questions and the mark schemes; and the job of the revisers is to review the questions and mark schemes commenting on a large range of issues about which they must be satisfied before the question papers and mark schemes are finalised.


The main purpose of the mark scheme is to provide a uniform basis for the marking process so that all the markers are following exactly the same instructions and making the same judgements in so far as this is possible. Before marking begins a standardising meeting is held where all the markers are briefed using the mark scheme and samples of the students’ work in the form of scripts. Consideration is also given at this stage to any comments on the operational papers received from teachers and their organisations. During this meeting, and up to and including the end of the marking, there is provision for amendments to be made to the mark scheme. What is published represents the final form of the mark scheme.


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AVAILABLEMARKS

Section A

1 D

2 D

3 C

4 B

5 B

6 B

7 C

8 B

9 D

10 D [1] for each correct answer [10] 10

Section A 10

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Section B

11 (a) Bent [1]

Two lone pairs and two bonding pairs [1]

Increased repulsion [1] [3] (b) (i)

× H

××O××

×H

H

+

[2]

(ii) As there is only one lone pair [1]

Greater repulsion between bonded pairs of electrons [1] [2]

(iii) Repulsion between positively charged ions [1] 8

12 (a) (i) The number of protons in (the nucleus of) an atom [1]

(ii) The total number of protons and neutrons in (the nucleus of) an atom [1]

(iii) Atoms which have the same atomic number but a different mass number (contain the same number of protons but a different number

of neutrons) [1]

(b) Tennessine [1]

177 neutrons [1] [2] (c) Seven electrons in the outer energy level [1]

(d) (i) The mass of an atom of an isotope of an element relative to one-twelfth of the mass of an atom of carbon-12 [2]

(ii) 167.26 = [(161.93 × 0.14) + (163.93 × 1.60) + (165.93 × 33.50) + (RIM × 22.87) + (167.93 × 26.98) + (169.94 × 14.91)]/100 RIM = 166.93 [3] 11

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13 (a) Au + 3HNO3 + 4HCl → HAuCl4 + 3NO2 + 3H2O [2] (b) (i) +3 [1] (ii) Au 0 → +3 [1]

Cl 0 → –1 [1]

0 to +3 oxidation and 0 to –1 reduction [1] [3] (c) (i) Small difference/similar in electronegativity [1]

(ii) white fumes/smoke/solid with stopper from bottle of concentrated ammonia solution/glass rod dipped in concentrated ammonia solution [2] 9

14 (a) 1s22s22p63s1 [1]

Outer electron in an s orbital/sub-shell [1] [2] (b) (i) Number of atoms in 12.000 g of carbon-12 [1] (ii) 6/58.5 = 0.103 0.103 × (6.02 × 1023) = 6.20 × 1022 [2] (c) (i) High melting point [1]

Unreactive [1] [2] (ii) Cleans the wire [1]

Helps the solid stick to the wire [1] Forms (volatile) chlorides [1] to a maximum of [2] [2] (iii) Yellow/orange [1] (d) Dissolve the solid in water/dilute nitric acid [1]

Add silver nitrate solution [1]

White [1] precipitate [1] [4] (e) Add solid to a named dilute acid [1]

Effervescence [1]

Gas turns limewater milky [1] [3]

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AVAILABLEMARKS

AVAILABLEMARKS

(f) (I)

Cl

Cl

–

–Mg

2+

[2]

(ii) Na2CO3 + MgCl2 → 2NaCl + MgCO3 [2] (iii) Relative formula mass of magnesium carbonate = 84

Number of moles of magnesium carbonate = 1.4/84 = 0.017

Number of moles of sodium carbonate = 0.017

Relative formula mass of sodium carbonate = 106 Mass of sodium carbonate in the sample = 0.017 × 106 = 1.8 g

Percentage of sodium carbonate in the sample

1.8/6 × 100 = 30% error [–1] [6] 27

15 (a) Increase across the period [1]

Increasing nuclear charge/Atomic radius decreases [1]

Outer electron in same shell more strongly attracted to the nucleus [1] [3]

(b) (i) P(g) → P+(g) + e– [2]

(ii) P has half-filled p sub-shell [1]

Half-filled p sub-shells have stability increased [1] [2]

(c) Indicative Content – comments related to melting points

• Silicon – giant covalent • Many strong covalent bonds to be broken • P, S, Cl (Ar) molecular • Comment on van der Waals • van der Waals depend on RMM • S8 versus P4

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Response Mark

Candidates must use appropriate specialist terms to describe fully the trend in melting point from silicon to argon (using a minimum of 6 points of indicative content including reference to all three distinct parts of the trend). They use good spelling, punctuation and grammar and the form and style are of a high standard.

[5]–[6]

Candidates must use appropriate specialist terms to describe the trend in melting point from silicon to argon (using a minimum of 4 points of indicative content). They use satisfactory spelling, punctuation and grammar and the form and style are of a satisfactory standard.

[3]–[4]

Candidates briefly and partially describe the trend in melting point from silicon to argon (using a minimum of 2 points of indicative content). They use limited spelling, punctuation and grammar and they have made little use of specialist terms. The form and style are of a limited standard.

[1]–[2]

[6] 13

16 (a) Pb(NO3)2 + 2KI → PbI2 + 2KNO3 [2]

(b) 75.6 mg = 0.0756 g

Moles lead iodide in 100 cm3 = 0.0756/461 = 1.64 × 10–4

Moles of iodide ions in 100 cm3 = 3.28 × 10–4

Molarity of iodide ions = 3.28 × 10–4/0.1 = 3.28 × 10–3 M [4]

(c) (i) yellow/brown [1] (ii) blue-black [1] (d) (i) Any three from: steamy/misty fumes violet/purple vapour smell of rotten eggs yellow solid grey-black solid (on the sides of the test-tube) [3]

(ii) Phosphoric acid is not an oxidising agent [1] 12

Section B 80

Total 90

AS1 Chemistry Papers: 2006 - 2014 & Mark …...10 The enthalpy of neutralisation when an acid reacts with an alkali is, by definition, the number of kilojoules released by A the formation

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