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L.34

PRE-LEAVING CERTIFICATE EXAMINATION, 2010

CHEMISTRY - HIGHER LEVEL

TIME : 3 HOURS

400 MARKS

Answer eight questions in all

These must include at least two questions from Section A

All questions carry equal marks (50)

Information

Relative atomic masses: H = 1, C = 12, O = 16, Cl = 35.5, Ca = 40

Molar volume at room temperature and pressure = 24.0 litres

Avogadro constant = 6 1023 mol–1

Universal gas constant, R = 8.31 J K mol–1

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Section A Answer at least two questions from this section [see page 1 for full instructions]. 1. An experiment was carried out to determine the concentration of sodium hypochlorite (NaClO) in a sample

of household bleach. A 25 cm3 sample of bleach was diluted to 250 cm3 with deionised water. A 25 cm3 portion of the diluted bleach was placed in a conical flask and an excess of acidified potassium iodide was added. It was then titrated against a 0.21 M solution of sodium thiosulfate a number of times. The average titration volume was 19.6 cm3.

The equations for the reactions are:

ClO– + 2I– + 2H+ Cl– + I2 + H2O 2S2O

23 + I2 S4O

26 + 2I–

(a) Outline how to accurately dilute the bleach. (12)

(b) Why was an excess of potassium iodide used? (5)

(c) Why is potassium iodide used instead of iodine itself? (3)

(d) Name the reagent that is used as a primary standard in redox titrations. (6)

(e) Name the indicator for this experiment and state when to add it. (6)

(f) Calculate the concentration of sodium hypochlorite (NaClO) in (i) the diluted bleach, (ii) the original bleach, in moles per litre. Express the concentration in the original bleach in terms of % (w/v). (18)

2. A group of students prepared ethanal (CH3CHO) in the laboratory by slow addition of an aqueous solution

of ethanol (C2H5OH) and sodium dichromate (VI) (Na2Cr2O7.2H2O) to a hot solution of sulfuric acid (H2SO4). The apparatus was set up as shown in the diagram. The reaction is described by the following equation:

3C2H5OH + Cr2O2

7 + 8H+ 3CH3CHO + 2Cr3+ + 7H2O

(a) Why was the receiving vessel in ice-water? (4)

(b) Why is it not necessary to heat the flask during the preparation? (4)

(c) Describe and explain the colour change that occurred in the flask during the addition. (9)

(d) Describe how to test ethanal with ammoniacal silver nitrate (Tollens’ reagent) and describe the result. (12)

(e) How would you use the same apparatus if you were to prepare ethanoic acid? What change would you make to the reagents? (9)

(f) Calculate the amount (in grams) of ethanal that could be expected from 25 cm3 of ethanol (density 0.8 g cm–3) if a 60% yield was obtained. (12)

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3. A student carried out an experiment to measure the relative molecular mass of a volatile liquid. This involved measuring the mass of a certain amount of the vapour of the volatile liquid, and its temperature, pressure and volume. Calculations were made using the equation of state for an ideal gas, PV = nRT.

(a) Draw a labelled diagram of a suitable apparatus that could be used for this experiment. (8) (b) Describe how to measure the mass of the vapour. (12) (c) Outline how to measure the (i) volume, (ii) pressure and (iii) temperature of the vapour. (12) In an experiment to measure the relative molecular mass of a volatile liquid 0.82 g of the liquid was

vaporised at a temperature of 99 °C. The volume occupied was found to be 255 cm3. The pressure was 100,000 N/m2 (Pa).

(d) Calculate the number of moles of the volatile liquid vaporised. (12) (e) Calculate the relative molecular mass of the volatile liquid. (6)

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Section B [See page 1 for instructions regarding the number of questions to be answered.] 4. Answer eight of the following items (a), (b), (c), etc. (50) (a) Name an aromatic acid-base indicator.

(b) Name the scientist pictured on the right who discovered the neutron by bombarding beryllium with radiation. What kind of radiation did he use?

(c) Write the electronic configuration (s, p, etc.) of a chromium atom.

(d) How many molecules of oxygen are present in 250 cm3, measured at s.t.p.?

(e) Define catalyst.

(f) When rain water passes through limestone rock it acquires a calcium compound that causes temporary hardness. Show by an equation how this substance is formed.

(g) Name this ester: CH3COOCH3.

(h) What two measurements can be found from the mass spectrum of an element?

(i) Name two plant nutrients that are removed in the tertiary treatment of sewage.

(j) What is meant by the activation energy of a reaction?

(k) Answer part A or part B.

A Place these atmospheric gases in order of their greenhouse effect, with the least effective first: carbon dioxide methane water

or

B Name two solid substances that are added to iron ore in a blast furnace to extract the iron.

5. (a) Define electronegativity. (5)

(b) Explain why the values of electronegativity increase across a period of the periodic table of the elements. (6)

(c) Use electronegativity values (Mathematical Tables p 46) to predict the kinds of bonding in (i) ammonia, (ii) magnesium oxide. (6)

(d) Use a dot and cross diagram to show the bonding in magnesium oxide. Would you expect magnesium oxide to dissolve in hexane? Explain your answer. (12)

(e) Use a dot and cross diagram to show the bonding in a molecule of ammonia. Explain why the bond angle is 107°. (12)

(f) Explain why ammonia has a much higher boiling point (–33 °C) than methane (–161 °C). (9)

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6. (a) Define octane number of petrol. (5)

(b) Name one of the fractions produced in the oil refinery that contains molecules suitable for petrol. (3)

(c) Name and draw the molecular structure of a hydrocarbon with an octane number of 100. (6)

(d) Explain how (i) isomerisation and (ii) dehydrocyclisation raise the octane number of a fuel. (6)

(e) Apart from the effect on human health, why is lead unsuitable for improving the octane number of petrol? Name a suitable oxygenate that raises the octane number. (6)

(f) Name two components of LPG. What is added to LPG and to natural gas to give it a strong odour? (9)

(g) Write a balanced equation for the combustion of methane (CH4). If the heat of combustion of methane is –890 kJ mol–1 and the heats

of formation of carbon dioxide and of water are –394 kJ mol–1 and –286 kJ mol–1 respectively, calculate the heat of formation of methane. (15)

7. Examine the reaction scheme and answer the questions that follow. X Y Z C2H5OH C2H4 C2H6 C2H5Cl A B C D (a) Which of the compounds is most soluble in water? Explain your answer. (6) (b) For each of the conversions X, Y and Z, classify it as an addition, an elimination or a substitution

reaction. (9) (c) What term describes the mechanism for conversion Z? Name or give the formula for the other

reagent. What provides the energy to initiate the reaction? Name or give the formula for the compound that is formed in trace amounts in the conversion. (14)

(d) Conversion X can be carried out in a school laboratory. Name a catalyst for this reaction.

Describe how to show that the product is unsaturated. (9) (e) Describe what you would see if a small piece of sodium is added to compound A. Write a

balanced equation for the reaction. Explain why compound A has reacted as an acid in this reaction. (12)

8. (a) Define (i) acid, (ii) conjugate pair, according to the Brønsted-Lowry theory. (8) (b) Identify one acid and its conjugate base in this reaction:

NH4 + CO 2

3 HCO3 + NH3 (6)

(c) Calculate the pH of a 0.1 M aqueous solution of NH3 if the Kb for NH3 is 1.8 × 10–5. (12) (d) State two limitations to the pH scale. (6) (e) If an indicator that is a weak acid is represented as HIn, explain how it works as an indicator. (12) (f) Name a suitable indicator for the titration between ammonia solution and hydrochloric acid and

give a reason for your choice. (6)

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9. (a) State Le Chatelier’s principle. (8) (b) Cobalt chloride was dissolved in water to form a pink solution, according

to the following equation:

CoCl 24 + 6H2O Co(H2O) 2

6 + 4Cl–

The forward reaction is exothermic.

(i) What change, if any, would you observe if concentrated hydrochloric acid was added to the mixture? Explain your answer. (9)

(ii) Would you heat or cool the pink solution to change the colour to blue? Explain your answer. (6)

(c) Hydrogen iodide decomposes to form hydrogen and iodine gases according to the following

equation:

2HI(g) H2(g) + I2(g)

(i) Write the equilibrium constant (Kc) expression for the reaction. (6)

(ii) Two moles of hydrogen iodide were placed in a vessel at 683 K and allowed to come to equilibrium. The Kc at 683 K is 0.0156. Calculate the number of moles of hydrogen iodide that is present at equilibrium. (15)

(iii) Why is it not necessary to know the volume of the flask in this calculation? (6)

10. Answer any two of the parts (a), (b) and (c). (2 × 25) (a) (i) Define the rate of a chemical reaction. (4)

(ii) Why does the rate of a chemical reaction generally decrease with time? (3)

(iii) Explain why the reaction between ionic compounds in solution is generally instantaneous, while reactions between covalent compounds are slower. (6)

(iv) Explain why an increase in temperature generally increases the rate of a chemical reaction. (6)

(v) Why can explosions occur in flour storage silos? (6) (b) In 1922 Neils Bohr was awarded the Nobel Prize in physics for his theory of

atomic orbits, which he developed from a consideration of the emission spectrum of hydrogen.

(i) Distinguish between an atomic orbit and an atomic orbital. (4)

(ii) How is one atomic orbital different to another orbital? (3)

(iii) Name the series of lines in the visible region of the hydrogen spectrum. (3)

(iv) Explain why these lines are emitted from hydrogen gas that is conducting electricity. (15) (c) (i) Define oxidation in terms of electron transfer. (4)

(ii) What is the oxidation number of S in (a) SO 23 , (b) SO 2

4 ? (6)

(iii) Outline how to use a solution of sodium sulfite to demonstrate that chlorine water is an oxidising agent. (12)

(iv) Explain why the oxidising strength of halogens decreases down the group. (3)

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11. Answer any two of the parts (a), (b) and (c). (2 × 25) (a) Calcium carbonate reacts with hydrochloric acid according to the balanced equation:

CaCO3 + 2HCl CaCl2 + H2O + CO2

Carbon dioxide reacts with calcium hydroxide according to the balanced equation:

2CO2 + Ca(OH)2 Ca(HCO3)2

(i) What mass of carbon dioxide is formed when 10 g of calcium carbonate reacts completely with hydrochloric acid? (9)

(ii) What is the volume of that amount of gas at room temperature and pressure? (6)

(iii) What mass of calcium hydrogencarbonate would be formed if that amount of carbon dioxide reacted completely? (6)

(iv) Give the common name for a solution of Ca(OH)2. (4) (b) (i) Explain the principle upon which chromatography is based. (9)

(ii) Outline an experiment to separate the components of a mixture of dyes using paper or thin-layer chromatography. (12)

(iii) Name the instrument that may be attached to a gas chromatograph to identify the molecular mass of the components. (4)

(c) Answer either part A or part B. A

(i) What is meant by the term the greenhouse effect of the atmosphere? (6)

CO2 is a greenhouse gas that is produced by human activity. (ii) Apart from the combustion of fossil fuels name an industrial process that releases large

quantities of CO2 into the atmosphere. Name a natural process that removes CO2 from the atmosphere. (6)

(iii) After CO2 dissolves in water it may be found in its free state or as two ions. Show by equations how these ions are formed. (9)

(iv) Name the group of greenhouse gases that also cause ozone depletion. (4)

or B

(i) Name the father and son team, pictured right, who pioneered a method of determining crystal structure. What method did they develop? (6)

Diamond and graphite are different forms (allotropes) of carbon. (ii) What type of crystal do they possess?

Name the form that conducts electricity and explain how its crystal structure gives it this physical property. (12)

(iii) Name a substance that forms molecular crystals with Van der Waals forces as the binding force between the lattice points. Give one physical property based on this crystal structure. (7)

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