1 In the UK, almost all the sulphuric acid, H 2 SO 4 , is manufactured by the Contact process. One stage in the Contact process involves the reaction between sulphur dioxide and oxygen. 2SO 2 (g) + O 2 (g) 2SO 3 (g) Table 1.1 below shows values of the equilibrium constant, K p , for this equilibrium at different temperatures. (a) Write an expression for the equilibrium constant, K p , of this reaction. (b) In this question, one mark is available for the quality and use of scientific terms. • The conversion of sulphur dioxide and oxygen into sulphur trioxide is carried out at slightly above atmospheric pressure. Comment on this statement. • Explain what happens to the equilibrium amounts of SO 2 , O 2 and SO 3 as temperature increases at constant pressure. Deduce the sign of ΔH for the forward reaction in the equilibrium. Explain your reasoning carefully. ................................................................... ................................................................... ........... 1 Class Reg Number Candidate Name ................................................................... ....
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1 In the UK, almost all the sulphuric acid, H2SO4, is manufactured by the Contact process.One stage in the Contact process involves the reaction between sulphur dioxide and oxygen.
2SO2(g) + O2(g) 2SO3(g)
Table 1.1 below shows values of the equilibrium constant, Kp, for this equilibrium at different temperatures.
(a) Write an expression for the equilibrium constant, Kp, of this reaction.
(b) In this question, one mark is available for the quality and use of scientific terms.
• The conversion of sulphur dioxide and oxygen into sulphur trioxide is carried out at slightly above atmospheric pressure. Comment on this statement.
• Explain what happens to the equilibrium amounts of SO2, O2 and SO3 as temperature increases at constant pressure.
Deduce the sign of ΔH for the forward reaction in the equilibrium. Explain your reasoning carefully.
(c) An equilibrium is set up for the SO2, O2, SO3 equilibrium at 400 °C.At this temperature
• the equilibrium partial pressure of SO2 is 10 kPa• the equilibrium partial pressure of O2 is 50 kPa• Kp = 3.0 × 102 kPa–1.
Calculate the equilibrium partial pressure of SO3 at 400 °C. Hence determine the percentage of SO3 in the equilibrium mixture at this temperature.
(d) In the UK, almost all the sulphuric acid manufactured uses sulphur as a starting material for SO2 production. In some countries, metal ores such as zinc sulphide, ZnS, are used instead to form SO2 by heating with air.
(i) Construct a balanced equation to show the reaction that takes place when zinc sulphide is heated in air.
2 Natural rain is slightly acidic because of the carbon dioxide dissolved in it. Rain which is more acidic than natural rain is called acid rain and this causes damage to the ecosystem in a number of ways. One effect of acid rain is the leaching of aluminium ions from clay soils into water-courses.
(a) Carbon is in the second period of the Periodic Table. Give the formula of a basic oxide of an element in the same period.
(b) The main equilibrium that is set up when carbon dioxide dissolves in water is shown below.
H2O(aq) + CO2(aq) H+(aq) + HCO3–(aq)
(i) Write the equation for the acidity constant, Ka, for this reaction and give its units.
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(ii) The numerical value of Ka in the above units is 4.5 × 10–7 at room temperature.Calculate the pH of unpolluted rain where [CO2(aq)] = 1.2 × 10–5mol dm–3.
(c) A clay material, such as kaolinite, reacts with a solution containing hydrogen ions as shown.
(i) Complete equation 3.2 by writing on the dotted line above. Include the appropriate state symbol.
(ii) An equilibrium constant for this equation is given by the following expression.
Kc =
Show that the value of Kc is approximately 1 × 1016 mol–4 dm12 given that
[Al3+(aq)] = 9 × 10–11 mol dm–3 when the pH = 6.
(iii) This equilibrium is very sensitive to pH.
Calculate the concentration of aluminium ions in solution when[H+] = 1 × 10–4 mol dm–3 (i.e. at pH = 4). Assume Kc = 1 × 1016 mol–4 dm12.
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(d) The ionic radius of an isolated (unhydrated) Al3+ ion is shown below, together withvalues for ions of two other metals in Period 3.
(i) In this question, two marks are available for the quality of the use and organisation of scientific terms.
Explain why the metals form ions with these charges. Say, in terms of their atomic structure, what the ions have in common and explain the trend in ionic radius across the period from Na+ to Al3+.
3 Most of the chemical elements found on Earth were produced in stars. Chemists have arranged the elements into a Periodic Table which allows them to make predictions about the behaviour of the elements and their compounds.
(a)(i) Calcium in Group 2 reacts with water to produce a solution of calcium hydroxide and bubbles of hydrogen gas.
Predict a balanced equation for the reaction of radium, Ra, with water. Include state symbols. Write your equation in the space below.
(c) Atoms react together to form molecules in the dense gas clouds in interstellar space.Molecules of H2S, NH3 and OCS (similar in structure to CO2) have been detected.
(i) Complete the dot-and-cross diagram for each molecule in the boxes below.
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(ii) Use the theory of ‘electron pair repulsion’ to decide which of the possible shapes below represents the shape of each molecule.
Write the formula of each of the molecules H2S, NH3 and OCS underneath its shape.
4 The question below relates to chlorides of some of the elements in Period 3 of the PeriodicTable.
(a) Draw ‘dot and cross’ diagrams to show the bonding in magnesium chloride and silicon(IV) chloride. Only draw the outer shell electrons.
(b) Describe the difference in behaviour when magnesium chloride and silicon(IV)chloride are added separately to cold water.
You may include in your answer• the pH of any resulting solution,• relevant chemical equations,• experimental observations,• the name of the process taking place.
(iii) Explain why solid aluminium chloride does not conduct electricity, but when aluminium chloride is added to water, the resulting solution will conduct electricity.
(ii) Calculate the enthalpy change of reaction, Hr, in kJ mol–1, for the thermal decomposition of barium nitrate using the enthalpy changes of formation, Hf, given in the table.
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(c) A student investigates the volume of gas formed when barium nitrate is heated.The diagram shows the apparatus the student uses.
(i) A 1.31 g sample of barium nitrate is completely decomposed.
Calculate the volume, in cm3, of gas formed at room temperature and pressure.1 mol of gas molecules occupies 24 000 cm3 at room temperature and pressure.
(ii) Suggest one problem that the student may encounter when carrying out the investigation.
(d) Barium nitrate has a higher decomposition temperature than calcium nitrate. One of the reasons for this is the difference between the lattice enthalpy of barium oxide and that of calcium oxide.
(i) Explain what is meant by the term lattice enthalpy.
6 The element titanium, Ti, atomic number 22, is a metal that is used in the aerospace industry for both airframes and engines.
A sample of titanium for aircraft construction was analysed using a mass spectrometer and was found to contain three isotopes, 46Ti, 47Ti and 48Ti. The results of the analysis are shown in Table 1.1 below.
7 The Group 2 element radium, Ra, is used in medicine for the treatment of cancer. Radium was discovered in 1898 by Pierre and Marie Curie by extracting radium chloride from its main ore pitchblende.
(d) Reactions of the Group 2 metals involve removal of electrons. The electrons are removed more easily as the group is descended and this helps to explain the increasing trend in reactivity.
(i) The removal of one electron from each atom in 1 mole of gaseous radium atoms is
called the ...............................................................................................................
8 A student had a stomach-ache and needed to take something to neutralise excess stomach acid. He decided to take some Milk of Magnesia, which is an aqueous suspension of magnesium hydroxide, Mg(OH)2.
(a) The main acid in the stomach is hydrochloric acid, HCl(aq), and the unbalanced equation for the reaction that takes place with Milk of Magnesia is shown below.
Balance the equation by adding numbers where necessary in the unbalanced equation above.
(b) The student’s stomach contained 500 cm3 of stomach fluid with an acid concentration of 0.108moldm–3. The student swallowed some Milk of Magnesia containing 2.42 g Mg(OH)2. He wondered whether this dose was sufficient to neutralise the stomach acid. Assume that all the acid in the stomach fluid was 0.108moldm–3 hydrochloric acid.
(i) Calculate the mass of Mg(OH)2 necessary to neutralise this stomach fluid.
(ii) Determine whether the student swallowed too much, too little, or just the right amount of Milk of Magnesia to neutralise the stomach acid.
(c) Chewing chalk has been used for many years to combat excess stomach acid and indigestion tablets often contain calcium carbonate, CaCO3. Suggest, with the aid of an equation, how these tablets work.
10 One of the largest uses of phosphorus is in boxes of safety matches. A safety match ignites when it is rubbed against the striking surface of the match box.The friction between the match head and the striking surface generates enough heat for the phosphorus to burn. This in turn provides enough energy for the decomposition of potassium chlorate(V), KClO3, on the match head.
2KClO3(s) 2KCl(s) + 3O2(g)
Sulphur on the match stick ignites and sufficient heat is generated to ignite paraffin wax and then the wood in the match.
(a) When the phosphorus burns, phosphorus(V) oxide forms. What is the formula of phosphorus(V) oxide?
(b) Calculate the volume of oxygen, measured at room temperature and pressure, that forms when 0.368 g of potassium chlorate(V) is decomposed.One mole of any gas occupies 24 000 cm3 at room temperature and pressure.
(c) Suggest why the match head contains potassium chlorate(V).
(b) The percentage purity of an impure sample of FeSO4.7H2O can be determined by titration against potassium dichromate(VI), K2Cr2O7, under acid conditions, using a suitable indicator.
During the titration, Fe2+(aq) ions are oxidised to Fe3+(aq) ions.
• Stage 1 – A sample of known mass of the impure FeSO4.7H2O is added to a conical flask.
• Stage 2 – The sample is dissolved in an excess of dilute sulphuric acid.• Stage 3 - The contents of the flask are titrated against K2Cr2O7(aq).
(i) The reduction half equation for acidified dichromate(VI) ions, Cr2O72–, is as follows.
Cr2O72–(aq) + 14H+(aq) + 6e– 2Cr3+(aq) + 7H2O(l)
Construct the balanced equation for the redox reaction between Fe2+(aq), Cr2O72–(aq) and
(b) A 1.50 × 10–2 mol dm–3 solution of HCOOH has [H+] = 1.55 × 10–3 mol dm–3.
(i) Calculate the pH of this solution and give one reason why the pH scale is a more convenient measurement for measuring acid concentrations than [H+].
13 The preparation of hydrogen iodide, HI(g), from hydrogen and iodine gases is a reversible reaction which reaches equilibrium at constant temperature.
H2(g) + I2(g) 2HI(g)
(a) Write the expression for Kc for this equilibrium.
(b) A student mixed together 0.30 mol H2(g) with 0.20 mol I2(g) and the mixture was allowed to reach equilibrium. At equilibrium, 0.14 mol H2(g) was present.
(i) Complete the table below to show the amount of each component in the equilibriummixture.
(ii) Calculate Kc to an appropriate number of significant figures. State the units, if any.
(c) The student compressed the equilibrium mixture so that its volume was reduced. The temperature was kept constant.
Comment on the value of Kc and the composition of the equilibrium mixture under these new conditions.
(e) Hydroiodic acid, HI(aq), is a strong acid that is an aqueous solution of hydrogen iodide.In the laboratory, hydroiodic acid can be prepared by the method below.
(i) Construct a balanced equation, with state symbols, for the preparation of hydroiodic acid.
(ii) Determine the percentage yield of hydroiodic acid.
(iii) Calculate the pH of the hydroiodic acid fraction.
14 This question looks at some chemicals found in food.1 mole of gas molecules occupies 24.0 dm3 at room temperature and pressure.The Avogadro constant = 6.02 × 1023 mol–1.
(a) A chemistry student bought a bar of chocolate. The student looked on the label and found that the main ingredient listed was ‘sugars’, making up 47.0% by mass of the 43.0 g chocolate bar. Throughout this question you can assume that all the sugars are present as sucrose, C12H22O11.
(i) How many sucrose molecules were in the bar of chocolate?
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A mixture of 480 g of iodine and 600 cm3 of water was put into a flask. The mixture was stirred and hydrogen sulphide gas, H2S(g), was bubbled through for several hours.
The mixture became yellow as sulphur separated out. The sulphur was filtered off and the solution was purified by fractional distillation. A fraction of HI(aq) was collected containing 440 g of HI in a total volume of 750 cm3.
(ii) The student ate the bar of chocolate. The standard enthalpy change of combustion of sucrose is –5640 kJ mol–1. On food labels, the energy content is measured in Calories.
1 Calorie = 4.18 kJ.
• Write an equation for the chemical change involved in the standard enthalpy change of combustion of sucrose.
• How much energy, in Calories, is available to the student from the sugars in the chocolate bar?
(b) An oxide of nitrogen is used as the propellant in whipped cream. This oxide contains 63.64% N by mass, and has a density of 1.833 g dm–3 at room temperature and pressure.
What is the molecular formula of this gas? Show your working.
(c) Chocolate mousse contains gelatine and a compound to promote fast setting of the mousse.
Compound A is such a setting agent. It has two acidic hydrogen atoms per molecule and is one of the six acids listed below.