L.34 PRE-LEAVING CERTIFICATE EXAMINATION, 2015 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). The information below should be used in your calculations. Relative atomic masses: H = 1, C = 12, O = 16, Fe = 56 Universal gas constant, R = 8.3 J K 1 mol 1 Molar volume at s.t.p. = 22.4 litres Molar volume at room temperature and pressure = 24.0 litres Avogadro constant = 6.0 10 23 mol 1 The use of the Formulae and Tables booklet approved for use in the State Examinations is permitted. A copy may be obtained from the examination superintendent. 2015.1 L.34 1/8 Page 1 of 8
Section A Answer at least two questions from this section [see page 1 for full instructions]. 1. In an experiment to determine the iron (Fe 2+ ) content of an iron tablet, a student dissolved four tablets, of total mass 0.96 g, in sulfuric acid and deionised water to make up a 250 cm 3 solution in a volumetric flask. This solution was then titrated in 25.0 cm 3 portions against a 0.012 M solution of potassium manganate(vii) (KMnO 4 ). The balanced equation for the titration reaction is: 5Fe 2+ + 8H + + MnO 4 5Fe 3+ + Mn 2+ + 4H 2 O (a) Why are iron tablets sometimes medically prescribed? (5) (b) Why was some dilute sulfuric acid used (with deionised water) to dissolve the iron tablets when making up the solution? (3) (c) Why was it necessary to standardise the potassium manganate(vii) solution immediately before use in the titration? What reagent was used for this purpose? (6) (d) Describe in detail the correct procedure for preparing the burette for use in the titrations. (12) (e) What colour change was observed as the solution from the burette flowed into the conical flask? (3) How did the student identify the endpoint of the titration? (3) (g) On average, 5.72 cm 3 of 0.012 M potassium manganate(vii) was required to react with 25.0 cm 3 portions of the iron solution prepared from the four tablets. Calculate (i) the molarity of the Fe 2+ solution, (ii) the total mass of iron in the 250 cm 3 of solution, (iii) the percentage by mass of iron in the tablets. (18) 2015.1 L.34 2/8 Page 2 of 8
2. Ethene gas was prepared from ethanol in a school laboratory. (a) Draw a labelled diagram showing the arrangement of apparatus and the reagents used in the preparation and collection of the ethene. (8) (b) Write a balanced equation for the preparation of ethene from ethanol. (6) (c) Why are the first few test tubes of ethene that were collected discarded? (3) (d) What precaution should be observed when heating is stopped? Why is this necessary? (6) (e) Describe how the ethene gas could have been shown to be unsaturated. (12) When ethanol is converted to ethene by this method, a 60% yield can be expected. Assuming this percentage yield, what is the maximum number of 65 cm 3 test tubes of ethene gas that could be collected at room temperature and pressure when 2.4 cm 3 of ethanol, density 0.8 g cm 3, react? (15) 3. Hydrogen peroxide decomposes in the presence of a manganese(iv) oxide (MnO 2 ) catalyst. (a) Write a balanced equation for the decomposition of hydrogen peroxide. (5) (b) Draw a labelled diagram to show how an apparatus can be assembled to measure the rate of decomposition of hydrogen peroxide in the presence of an MnO 2 catalyst. Indicate clearly on your diagram (i) how the reaction could be started at a time known exactly, (ii) how the gas produced is collected and its volume measured. (12) The table shows the volumes of gas produced at intervals over 8 minutes. Time (min) 0.0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 Volume of gas (cm 3 ) 0 12.0 18.0 21.2 23.0 23.8 24.0 24.0 24.0 (c) Plot a graph of volume of gas produced versus time. (12) (d) What conclusion can be drawn from the graph about the rate of reaction? (3) (e) Use your graph to find the instantaneous rate of reaction after 3 minutes. (6) Describe the change, if any, to the shape of your graph if the same mass of a finely powdered manganese(iv) oxide catalyst is used instead of a coarsely powdered (granulated) manganese(iv) oxide catalyst. Explain your answer. (6) (g) Describe the change, if any, to the shape of your graph if the reaction vessel is surrounded by ice. Explain your answer. (6) 2015.1 L.34 3/8 Page 3 of 8
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) Give two possible shapes for molecules of general formula AB 2. (b) How could you confirm the presence of phosphate ions ( PO 3 4 ) in an aqueous solution? (c) How many (i) electrons, (ii) neutrons are present in the 23 11Na ion? (d) How many atoms of hydrogen are there in 560 cm 3 of hydrogen at s.t.p.? (e) Name the analytical technique that is based on the promotion of electrons from their ground state energy levels to higher energy states. Name two metals which act as catalysts in the catalytic converters of modern cars. (g) Account for the difference in bond angle between water (104.5 ) and methane (109.5 ). (h) Define atomic orbital. (i) (j) When carbon was heated with 5.8 g of iron oxide, 4.2 g of iron was formed. Find by calculation the empirical formula of the iron oxide. Under what circumstances can ionic compounds conduct electricity? (k) Answer part A or part B. A B State two uses of nitrogen gas based on its chemical stability. or Name the main ore that is used commercially in the extraction of (i) aluminium, (ii) iron. 5. (a) Write the electron configuration (s, p) of an atom of zinc showing the distribution of electrons in atomic orbitals in the ground state. (6) (b) Use a dot and cross diagram to show the bonding in a magnesium oxide (MgO) molecule. (6) (c) Define electronegativity. (5) (d) Using electronegativity values predict the types of bonding in (i) methane (CH 4 ), (ii) water (H 2 O), (iii) aluminium fluoride (AlF 3 ). (9) (e) Outline the contribution of J.J. Thomson to the discovery of the atom. (6) In 1909 Ernest Rutherford carried out an experiment in which he bombarded thin sheets of gold foil with alpha particles. State and explain the three observations he made. (18) J.J. Thomson Ernest Rutherford 2015.1 L.34 4/8 Page 4 of 8
6. (a) Identify the two hydrocarbons used as references when establishing the octane number of a fuel. (6) (b) Fossil fuels are a non-renewable resource and they will eventually run out. Hydrogen is predicted by many as the fuel of the future. State two ways of making hydrogen on a large scale and write an equation to illustrate one of these reactions. (12) (c) Give one advantage and one disadvantage of using hydrogen as a fuel. (6) (d) The combustion of cyclohexane (C 6 H 12 ) is described by the following balanced equation: C 6 H 12 + 9O 2 6CO 2 + 6H 2 O Calculate the heat of combustion of cyclohexane given that the heats of formation of cyclohexane, carbon dioxide and water are 156, 394 and 286 kj mol 1, respectively. (13) (e) Upon catalytic cracking of a molecule of C 17 H 36, a C 5 H 10 molecule, a C 8 H 18 molecule and a four-carbon alkene molecule are obtained. Identify the four-carbon alkene obtained. Give the systematic IUPAC names of the three isomers of this four-carbon alkene. (13) 7. (a) What happens during the secondary treatment of sewage? Name one substance which is removed by tertiary treatment of sewage. (9) (b) Give one reason why heavy metal ions may be found in rivers and lakes. Give an example of a heavy metal ion and name a method of removing these ions from a water supply. (9) (c) A number of tests were carried out on a sample of swimming pool water to test its quality. (i) A colorimetric experiment was used to estimate the concentration of free chlorine in the swimming pool water sample. What is the general principle of all colorimetric experiments? (8) (ii) Describe how you would estimate the concentration of free chlorine in a sample of swimming pool water using either a comparator or a colorimeter. (12) (iii) When 1500 cm 3 of swimming pool water was filtered, the mass of the filter paper, upon drying, had increased by 0.84 g. When 300 cm 3 of the filtered sample was evaporated to dryness, the mass of the residue obtained was 0.38 g. Calculate the concentration of: (i) suspended solids in ppm, (ii) dissolved solids in ppm. (12) 8. The production of ammonia from nitrogen and hydrogen is normally carried out at a pressure of 200 atmospheres and a temperature of 450 C. (a) Name the process by which ammonia is produced commercially. (5) (b) What is the effect of a catalyst on a reversible reaction? (6) (c) The equilibrium reaction for the production of ammonia is: N 2 + 3H Fe 2 2NH 3 ΔH = 92.4 kj State the effect of each of the following on (i) the equilibrium reaction and (ii) the yield of ammonia (NH 3 ): (i) increasing the concentration of hydrogen, (ii) increasing the temperature, (iii) decreasing the concentration of ammonia. (18) (d) Write the equilibrium constant (K c ) expression for this reaction. (6) (e) Ten moles of nitrogen and thirty moles of hydrogen were mixed and allowed to come to equilibrium in a sealed 5 litre vessel at a temperature T. There were fifteen moles of ammonia in the vessel at equilibrium. Find the number of moles of nitrogen and hydrogen present in the mixture at equilibrium. Calculate the value of K c for this reaction at this temperature. (15) 2015.1 L.34 5/8 Page 5 of 8
9. Answer the questions below with reference to the compounds A D in the table on the right. (a) Give the systematic (IUPAC) names for A D. (8) A C 2 H 4 (b) Alcohols can be obtained by the reduction of aldehydes and ketones using hydrogen and a suitable catalyst. Name the catalyst used in these reduction reactions. (4) B C CH 3 CHO C 2 H 5 OH (c) Compound A can be synthesised from compound C in the school laboratory. D CH 3 COCH 3 (i) Write an equation for this reaction. (ii) What term is used to describe this type of reaction? (9) (d) Account for the difference in boiling points between compounds C and D. (8) (e) Describe how you would carry out a Fehling s test on compound B. What observation is made during this test? (12) Name the aromatic compound, found in almond kernels, that belongs to the same homologous series as compound B. (5) (g) Give one use of the family of compounds to which D belongs. (4) 10. Answer any two of the parts (a), (b) and (c). (2 25) (a) (i) Use the data below to plot, on graph paper, the ph curve for the titration between 25 cm 3 of ethanoic acid (CH 3 COOH), and a sodium hydroxide (NaOH) solution added from the burette. (18) Volume of NaOH added (cm 3 ) 0 5 10 15 17 19 21 23 25 30 35 40 ph 3.0 4.1 4.7 5.2 5.5 5.9 11.2 11.6 12.0 12.1 12.3 12.4 (ii) Name a suitable indicator for this titration. Use your graph to justify your choice. (7) (b) (i) Define oxidation number. (5) (ii) Use oxidation numbers to identify the species oxidised in the following reaction: ClO + I + H + Cl + I 2 + H 2 O (3) (iii) Use oxidation numbers to identify the oxidising reagent in the following reaction: I 2 + 2 S 2O 3 I + 2 S 4O 6 (4) (iv) What is the oxidation number of chromium in (i) HCrO 4, (ii) Cr 2 O 3? (6) (v) What is the oxidation number of oxygen in OF 2? Explain why this is the case for oxygen in this compound. (7) (c) Eutrophication in water may result from the addition of large quantities of nitrate fertilizers to it. Describe the processes occurring in the water leading to eutrophication. (7) A sample of river water was diluted from 20 cm 3 to 1 litre with well-aerated pure water. The dissolved oxygen concentration of half the sample was measured immediately; the other half was stored under suitable conditions and its dissolved oxygen concentration was measured later. Concentrations of dissolved oxygen of 9.7 ppm and 2.5 ppm, respectively, were recorded. (i) Under what conditions, and for how long, was the second sample stored? (9) (ii) Calculate the biochemical oxygen demand (BOD) of the undiluted river water. (9) 2015.1 L.34 6/8 Page 6 of 8
11. Answer any two of the parts (a), (b) and (c). (2 25) (a) Give the conjugate base of (i) sulfuric acid (H 2 SO 4 ), (ii) the weak acid HA. Which of these bases is stronger? Justify your answer. (13) In acting as an acid-base indicator bromothymol blue behaves like a weak base. Letting XOH represent bromothymol blue, it dissociates as follows: XOH X + + OH Colour A Colour B State and explain the colour observed when a few drops of the bromothymol blue solution are added to (i) a 0.1 M solution of HCl, (ii) a 0.1 M solution of NaOH. (12) (b) State Avogadro s law. (5) What is an ideal gas? (4) Under what conditions of temperature and pressure do real gases deviate most from ideal gas behaviour? (4) Calculate, using the ideal gas equation, the volume that 20 g of oxygen gas will occupy at 25 C and pressure 20 kpa. (12) (c) Answer part A or part B. A (i) Give the chemical formula for ozone. State one beneficial effect of the ozone layer. (7) (ii) What are CFCs? What use was made of CFCs before their production was restricted? (6) (iii) Explain, using equations, how CFCs contribute to ozone depletion. (12) B or (i) Write a brief note of the contribution made to our understanding of crystal structures by (i) Lawrence and William Bragg and (ii) Dorothy Hodgkin. (7) (ii) What is a crystal? (6) (iii) State the binding forces that exist in (i) a molecular crystal and (ii) a metallic crystal. (6) (iv) Give an example of (i) a covalent macromolecular crystal and (ii) an ionic crystal. (6) 2015.1 L.34 7/8 Page 7 of 8
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