QUESTIONSHEETS ENERGETICS II ENTHALPY OF ATOMISATION ENTROPY AND FREE ENERGY CHANGE PRINCIPLES OF ELECTROCHEMICAL CELLS

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1 CHEMISTRY QUESTIONSHEETS A2 Level A2 TOPIC 17 ENERGETICS II Questionsheet 1 Questionsheet 2 Questionsheet 3 Questionsheet 4 Questionsheet 5 Questionsheet 6 Questionsheet 7 Questionsheet 8 Questionsheet 9 Questionsheet 10 Questionsheet 11 Questionsheet 12 Questionsheet 13 Questionsheet 14 Questionsheet 15 Questionsheet 16 Questionsheet 17 ENTHALPY OF ATOMISATION LATTICE ENTHALPY BORN-HABER CYCLES I BORN-HABER CYCLES II ENTROPY AND FREE ENERGY CHANGE ENTROPY AND STATE TEST QUESTION I TEST QUESTION II TEST QUESTION III PRINCIPLES OF ELECTROCHEMICAL CELLS ELECTRODE POTENTIALS AND CELL EMF MEASUREMENT OF ELECTRODE POTENTIALS FACTORS INFLUENCING ELECTRODE POTENTIALS AND CELL EMF PREDICTION OF REDOX CHANGES FROM STANDARD ELECTRODE POTENTIALS FUEL CELLS TEST QUESTION IV TEST QUESTION V 13 marks 16 marks 16 marks 16 marks 17 marks 21 marks 15 marks 15 marks 15 marks 20 marks 17 marks 13 marks 18 marks 18 marks 16 marks 16 marks 17 marks Authors Trevor Birt Donald E Caddy Andrew Jones Adrian Bond Editors John Brockington Stuart Barker John Brockington Kevin Frobisher Andy Shepherd Stuart Barker Curriculum P ress Curriculum Press Licence Agreement: Paper copies of the A-Level Chemistry Questionsheets may be copied free of charge by teaching staff or students for use within their school, provided the Photocopy Masters have been purchased by their school. No part of these Questionsheets may be reproduced or transmitted, in any other form or by any other means, without the prior permission of the publisher. All rights are reserved. This license agreement is covered by the laws of England and Wales Curriculum Press March Curriculum Press Bank House 105 King Street Wellington Shropshire TF1 1NU

2 TOPIC 17 Questionsheet 1 ENTHALPY OF ATOMISATION a) Distinguish between the standard enthalpy of atomisation of an element and the standard enthalpy of atomisation of a covalent compound.... [4] b) What is the relationship between the enthalpy of atomisation of hydrogen and the bond dissociation enthalpy of the H H bond? c) Given that the standard enthalpy of atomisation of water is 920 kj mol -1, calculate the mean bond dissociation enthalpy of the O H bond in water. According to the data books, the mean bond dissociation enthalpy of the O H bond is 463 kj mol -1. How do you account for the difference between this value and the one you have calculated in c)? d) Write an equation for the formation of methane from its elements. Given that the standard enthalpy of formation of methane is 74.9 kj mol -1, and the standard enthalpies of atomisation of carbon and hydrogen are 715 and 218 kj mol-1 respectively, calculate the mean bond dissociation enthalpy of the C H bond in methane.... [4] TOTAL / 13

3 TOPIC 17 Questionsheet 2 LATTICE ENTHALPY a) Define the lattice enthalpy of a simple ionic compound. b) For the chlorides of the metals of Group 1, how does the magnitude of lattice enthalpy change on progressing from LiCl to CsCl? Explain your answer. Change... [1] Explanation [2] From lattice enthalpy considerations only, would you expect the following physical properties of the Group1 chlorides to increase, decrease or remain approximately constant on progressing from LiCl to CsCl? Melting point... [1] Solubility in water... [1] c) Lattice enthalpies and enthalpies of hydration are given in the table for a selection of compounds and ions. Salt Lattice enthalpy Ion Hydration enthalpy /kj mol -1 /kj mol -1 BaF 2 (s) 2352 Ba BaCl 2 (s) 2056 F Cl Use the appropriate data to construct an enthalpy diagram for the dissolving of barium fluoride in water, and calculate the enthalpy of solution of BaF 2 (s). [6] Use your value from c), and the information that the enthalpy of solution of BaCl 2 (s) is 7 kj mol -1, to explain why BaF 2 (s) is less soluble than BaCl 2 (s). TOTAL / 16

4 TOPIC 17 Questionsheet 3 BORN-HABER CYCLES I a) Given the thermodynamic data below, construct a Born-Haber cycle for the formation of lithium fluoride and then calculate a value for its lattice enthalpy. kj mol -1 Enthalpy of atomisation of Li(s) Enthalpy of atomisation of F 2 (g) 79.0 First ionisation energy of Li(g) First electron affinity of F(g) Enthalpy of formation of Li + F - (s) Marks are awarded for correct labelling of both enthalpy levels and arrows. Born-Haber cycle Lattice enthalpy of lithium fluoride [12] b) The lattice enthalpy of a compound, determined from a Born-Haber cycle as in a), generally differs from the value that can be calculated from the charge numbers and radii of its ions. Suggest a reason for this. TOTAL / 16

5 TOPIC 17 Questionsheet 4 BORN-HABER CYCLES II a) Magnesium chloride has the formula MgCl 2, i.e. Mg 2+ (Cl - ) 2, and a standard enthalpy of formation of -642 kj mol -1. Use the data below to calculate the standard enthalpy of formation of the hypothetical chloride MgCl, i.e. Mg + Cl -, and hence explain why, when magnesium and chlorine react together, the product is MgCl 2 rather than MgCl. Begin your answer by drawing a Born-Haber cycle. Atomisation enthalpy of magnesium = 150 kj mol -1 Atomisation enthalpy of chlorine = 121 kj mol -1 First ionisation energy of magnesium = 736 kj mol -1 First electron affinity of chlorine = -364 kj mol -1 Lattice enthalpy of MgCl = 756 kj mol -1 Born-Haber cycle for the formation of MgCl [6] TOTAL (Continued...) /

6 TOPIC 17 Questionsheet 4 Continued BORN-HABER CYCLES II Standard enthalpy of formation of MgCl Why MgCl 2 is formed, rather than MgCl b) Which one of the data listed in a) would have been obtained by calculation rather than experiment? c) If you were constructing a Born-Haber cycle for MgCl 2 rather than MgCl, what addtional data would you need? d) Would you expect the lattice enthalpy of MgCl to be greater or less than that of MgCl 2? Explain you answer. Greater or less than Hê lattice (MgCl 2 )... [1] Explanation... TOTAL / 16

7 TOPIC 17 Questionsheet 5 ENTROPY AND FREE ENERGY CHANGE a) Give the name and units of each physical quantity in the equation: Gê = Hê - T. Sê Gê... [1] T... [1] Hê... [1] Sê... [1] b) State, with reasons, whether you would expect the standard entropy change in each of the following reactions to increase, decrease or remain the approximately the same. 2Mg(s) + O 2 (g) 2MgO(s) Mg(s) + 2HCl(aq) MgCl 2 (aq) + H 2 (g) H 2 (g) + Cl 2 (g) 2HCl(g) (iv) 2NH 3 (g) N 2 (g) + 3H 2 (g) c) The combustion of ethanol is represented by the equation: C 2 H 5 OH(l) + 3O 2 (g) 2CO 2 (g) + 3H 2 O(l) Given the following data: Sê/J K -1 mol -1 : C 2 H 5 OH 160.7; O 2 (g) 0; CO 2 (g) 213.6; H 2 O(l) 69.9; calculate the standard entropy change for the reaction. Given the additional information that the standard enthalpy of combustion of ethanol is kj mol -1, calculate the standard free energy change for the reaction. TOTAL / 17

8 TOPIC 17 Questionsheet 6 ENTROPY AND STATE a) Using these standard entropies (J K -1 mol -1 ) : H 2 (g) 130; Cl 2 (g) 223; HCl(g) 187; Na(s) 51; NaCl(s) 72 Explain why two values are considerably lower than the others Calculate the standard entropy change (per mole chloride formed) for the reaction of chlorine with: Sodium Hydrogen [3] [3] Comment on the sign of the value calculated in for NaCl in the light of the reaction being exothermic.... [4] TOTAL (Continued...) /

9 TOPIC 17 Questionsheet 6 Continued ENTROPY AND STATE b) Consider the standard enthalpy of vaporisation and boiling point for these hydrogen halides: Hydrogen halide Boiling point (K) Hê vaporisation (kj mol -1 ) HF HCl HBr HI Sê vap = - Hê vap T (T = boiling point in K) Use these values to calculate the standard entropies of vaporisation for HF HCl [2] Explain why the values in can be calculated using the function given. [2] Suggest why HF has a lower vaporisation entropy than the other hydrogen halides which are similar in value. TOTAL / 21

10 TOPIC 17 Questionsheet 7 TEST QUESTION I The reaction between sodium and water can be represented by the following equation: Na(s) + H 2 O(l) Na + (aq) + OH - (aq) + ½H 2 (g); Hê = kj mol -1 a) The reaction can be broken down into the following seven steps. Name the enthalpy change for each step. Step 1 Na(s) Na(g); Hê = kj mol -1 Name... [1] Step 2 Na(g) Na + (g) + e - ; Hê = kj mol -1 Name... [1] Step 3 Na + (g) + aq Na + (aq); Hê = -406 kj mol -1 Name... [1] Step 4 H 2 O(l) + aq H + (aq) + OH - (aq); Hê x Name... [1] Step 5 H + (aq) H + (g) + aq ; Hê = kj mol -1 Name... [1] Step 6 H + (g) + e- H(g); Hê = kj mol -1 Name... [1] Step 7 H(g) ½H 2 (g); Hê = -218 kj mol-1 Name... [1] b) Sketch an enthalpy diagram for this reaction. [4] TOTAL (Continued...) /

11 TOPIC 17 Questionsheet 7 Continued TEST QUESTION I b) Apply Hess s law to your diagram and then calculate the standard enthalpy change ( Hê x ) for Step 4. c) From your answer to b), predict a value for the enthalpy of neutralisation of a strong acid by a strong base and explain your answer. Predicted value... [1] Explanation... TOTAL / 15

12 TOPIC 17 Questionsheet 8 TEST QUESTION II Magnesium oxide is an ionic compound that can easily by prepared by burning magnesium in excess oxygen. It has a melting point of 2827 C and so is commonly used for manufacturing linings for open-hearth steel furnaces. Use the information in the table below to answer the following questions. Hê / kj mol -1 O 2 (g) 2O(g) +498 O(g) + e- O - (g) -142 O - (g) + e- O 2- (g) +844 Mg(s) Mg(g) +147 Mg(g) Mg + (g) + e Mg + (g) Mg 2+ (g) + e Mg 2+ (g) + O 2- (g) MgO(s) a) Write a chemical equation which represents the standard enthalpy of formation of magnesium oxide. b) Construct a Born-Haber cycle which could be used to calculate the standard enthalpy of formation of magnesium oxide. Write, by each arrow, the value of the enthalpy change it represents. [4] TOTAL (Continued...) /

13 TOPIC 17 Questionsheet 8 Continued TEST QUESTION II Calculate the standard enthalpy of formation of magnesium oxide. c) The standard molar entropy, Sê, for each species involved in the standard enthalpy of formation of magnesium oxide is shown below. Sê / J K -1 mol -1 Mg(s) 32.7 O 2 (g) MgO(s) 27.0 Use the data to calculate the standard entropy change for the formation of magnesium oxide. d) Calculate the standard free energy change at 298 K for the formation of magnesium oxide. Show the equation on which your calculation is based. Why does this reaction not happen spontaneously at 298 K? TOTAL / 15

14 TOPIC 17 Questionsheet 9 TEST QUESTION III Magnesium hydroxide, Mg(OH) 2, and calcium hydroxide, Ca(OH) 2, are both hydroxides of Group 2 metals but have different solubilities in water. a) Explain, with a suitable example, the meaning of the following terms. Enthalpy of hydration (of an ion) Enthalpy of solution (of an ionic compound) b) Use the following data to answer parts and of this question. Enthalpy change Hê /kj mol -1 Lattice enthalpy of MgCl 2 (s) Enthalpy of solution of MgCl 2 (s) -155 Enthalpy of hydration of Cl - (g) -364 Construct an enthalpy diagram for the dissolving of MgCl2(s) in water. Your diagram should include the enthalpy of hydration of Mg2+(g), as well as the enthalpy changes shown in the table. Apply Hess s law to your enthalpy diagram to write an equation which relates the enthalpy of solution of a compound to its lattice enthalpy and the hydration enthalpies of its ions. [3] Calculate a value for the enthalpy of hydration of Mg2+(g) TOTAL (Continued...) / 20

15 TOPIC 17 Questionsheet 9 Continued TEST QUESTION III c) Describe how the value of Hê solution changes as solubility increases. Explain why, on descending Group 2, the lattice enthalpy of the metal hydroxides decreases. Explain why calcium hydroxide is more soluble in water than magnesium hydroxide under standard conditions, even though the hydration enthalpy of cations decreases down Group 2. TOTAL / 15

16 TOPIC 17 Questionsheet 10 PRINCIPLES OF ELECTROCHEMICAL CELLS a) A redox reaction can occur between metallic zinc and aqueous copper(ii) sulfate. Write down the two ionic half-equations, one for the oxidation half-reaction and the other for the reduction half-reaction. Oxidation... [1] Reduction... [1] Combine these so as to give an ionic equation for the overall reaction. The reaction can be carried out in a beaker, by adding powdered zinc to an aqueous solution of copper(ii) sulfate. Describe briefly what you would observe when this reaction occurs. b) The reaction can also be conducted in an electrochemical cell, so that the two half-reactions occur in separate half-cells as shown below. Zn SALT BRIDGE Cu ZnSO 4 (aq) CuSO 4 (aq) What would you observe while the cell was in use? What is the essential difference between the energy released here and when the experiment is carried out in a beaker? TOTAL (Continued...) / 17

17 TOPIC 17 Questionsheet 10 Continued PRINCIPLES OF ELECTROCHEMICAL CELLS c) Why must every electrochemical cell comprise two half-cells? What is the purpose of the salt bridge and what might it contain? What particles flow through the external circuit, and in which direction do they flow? (iv) What are the principal particles which flow through the salt bridge, and in which direction do they flow? TOTAL / 20

18 TOPIC 17 Questionsheet 11 ELECTRODE POTENTIALS AND CELL EMF a) Explain how an electrode potential can develop when a metal rod is partially immersed in an aqueous solution of one of its salts. What three energy changes contribute to the electrode potential of a metal? For comparison purposes, electrode potentials must be measured under standard conditions of temperature and concentration. What are these standard conditions? (iv) Under the same conditions of temperature and concentration, different metals give different electrode potentials. Why is this? b) In some redox electrodes, the reductant is not a metal rod: instead, the reductant and oxidant are two chemically related species in solution, e.g. Fe 2+ (aq) and Fe 3+ (aq). When setting up a standard electrode of this sort: how is electrical contact made with the rest of the circuit? what should be the concentration of each species in solution? c) What is meant by the electromotive force (e.m.f.) of an electrochemical cell? How could it be measured in the laboratory? Write down a mathematical equation which relates the e.m.f. of a cell to the potentials of its electrodes. TOTAL / 17

19 TOPIC 17 Questionsheet 12 MEASUREMENT OF ELECTRODE POTENTIALS a) Why must all electrode potentials be related to that of a reference electrode? By international convention, the standard hydrogen electrode (SHE) is used as a primary reference electrode and assigned a potential of 0.00 V. Label the diagram of this electrode, making clear the necessary conditions. Temperature = [4] b) The following cell was used to determine the standard electrode potential of the Fe 3+ (aq)/fe 2+ (aq) couple. Pt + V _ SALT BRIDGE Cu 1M FeCl 3 (aq) & 1M FeCl 2 (aq) 1M CuSO 4 (aq) Write ionic half-equations for the two half-reactions. Reduction half-equation... [1] Oxidation half-equation... [1] Hence write an ionic equation for the complete cell reaction. Calculate the standard electrode potential of the Fe 3+ (aq)/fe 2+ (aq) couple. Given that Eê (Cu 2+ (aq)/ Cu(s)) = V and the e.m.f. of the cell was found to be +0.43V. (iv) If in the diagram Fe 3+ (aq)/fe 2+ (aq) was replaced by Ti 3+ (aq)/ti 2+ (aq) (Eê= -0.37V), calculate the e.m.f of the cell. TOTAL / 13

20 TOPIC 17 Questionsheet 13 FACTORS INFLUENCING ELECTRODE POTENTIALS AND CELL EMF An electrochemical cell was set up comprising a standard hydrogen electrode and a standard silver electrode Ag + (aq)/ag(s) Eê = V H + (aq)/ ½ H 2 (g) Eê = 0.00 V a) Give the conventional cell representation. Calculate the cell e.m.f. In which direction do electrons flow through the external circuit? (iv) Write an ionic equation for the cell reaction. (v) If the experiment is left for a long time, the ph of the hydrogen electrode changes. State whether it increases or decreases, and why. b) What would be the effect on the potential of the electrode concerned, and hence on the cell e.m.f., of each of the following modifications? Increasing the surface area of the platinum foil of the hydrogen electrode. Effect on electrode potential... [1] Effect on e.m.f.... [1] Increasing the hydrogen ion concentration of the hydrogen electrode. Effect on electrode potential... [1] Effect on e.m.f.... [1] Raising the temperature of the hydrogen electrode. Effect on electrode potential... [1] Effect on e.m.f.... [1]. (iv) Increasing the gas pressure of the hydrogen electrode. Effect on electrode potential... [1]. (v) Effect on e.m.f.... [1] Increasing the silver ion concentration of the silver electrode. Effect on electrode potential... [1] Effect on e.m.f.... [1] TOTAL / 18

21 TOPIC 17 Questionsheet 14 PREDICTION OF REDOX CHANGES FROM STANDARD ELECTRODE POTENTIALS For all parts of this Questionsheet you should refer to the following data. ½ S 2 O 8 2- (aq) + e - ¾ SO 4 2- (aq) Eê = V ½ Cl 2 (g) + e - ¾ Cl - (aq) Eê = V ½Cr 2 O 7 2- (aq) + 7H + (aq) + 3e - ¾ Cr 3+ (aq) + 7 / 2 H 2 O(l) Eê = V ½ O 2 (g) + 2H + (aq) + 2e - ¾ H 2 O(l) Eê = V ClO - (aq) + H 2 O(l) + e - ¾ ½ Cl 2 (g) + 2OH- (aq) Eê = V Fe 3+ (aq) + e - ¾ Fe 2+ (aq) Eê = V 2SO 4 2- (aq) + 10H+ (aq) + 5e- ¾ S 2 O 3 2- (aq) + 5H 2 O(l) Eê = V ½I 2 (aq) + e - ¾ I - (aq) Eê = V ½ S 4 O 6 2- (aq) + e - ¾ S 2 O 3 2- (aq) Eê = V a) Give the formulae of all the species produced when the following pairs of substances are mixed together in aqueous solution under standard conditions. (If there are no products, write none.) Sodium thiosulfate and iodine. Sodium thiosulfate and chlorine. Chromium(III) sulfate and potassium chloride. (iv) Iron(II) iodide and potassium dichromate(vi) in acidic solution. TOTAL (Continued...) /

22 TOPIC 17 Questionsheet 14 Continued PREDICTION OF REDOX CHANGES FROM STANDARD ELECTRODE POTENTIALS b) Imagine that each of the following reactions were to be carried out in an electrochemical cell, with oxidation in one half-cell and reduction in the other. Calculate the e.m.f. of the cell and then decide the outcome if the reaction were to be attempted in a beaker: would it go to completion, exist in dynamic eqilibrium, or not occur at all under standard conditions? 4Fe 3+ (aq) + 2H 2 O(l) 4Fe 2+ (aq) + O 2 (g) + 4H + (aq) Eê cell... [1] Outcome... [1] Cr 2 O 7 2- (aq) + 6I - (aq) + 14H + (aq) 2Cr 3+ (aq) + 3I 2 (aq) + 7H 2 O(l) E ê cell... [1] Outcome... [1] 2Cl 2 (g) + 2H 2 O(l) O 2 (g) + 4Cl - (aq) + 4H + (aq) E ê cell... [1] Outcome... [1] (iv) Cl 2 (g) + 2OH - (aq) ClO - (aq) + Cl - (aq) + H 2 O(l) E ê cell... [1] Outcome... [1] c) Comment on each of the following observations. On mixing together solutions of sodium peroxodisulfate, Na 2 S 2 O 8, and potassium iodide at room temperature, there is no immediate colour change. When concentrated hydrochloric acid is added to potassium dichromate(vi) solution, chlorine gas is evolved. TOTAL / 18

23 TOPIC 17 Questionsheet 15 FUEL CELLS a) What are the essential differences between a fuel cell and an ordinary electrochemical cell? b) The following diagram depicts an oxygen-hydrogen fuel cell. Several of these cells can be used together to provide power, for example, for public transport vehicles. + _ O 2 (g) H 2 (g) Hot concentrated KOH(aq) Unreacted O 2 (g) H 2 O(g) + some unreacted H 2 (g) Porous graphite electrodes Write an ionic half-equation for the reaction at the electrode involving hydrogen. Write an ionic half-equation for the reaction at the electrode involving oxygen. Write an overall equation for the cell reaction. (iv) Suggest why the graphite electrodes need to be porous. TOTAL (Continued...) 16/

24 TOPIC 17 Questionsheet 15 Continued FUEL CELLS c) Calculate the volume of hydrogen, at 1 atm and 298 K, needed to generate 1.00 kwh of electricity. 1 kilowatt hour = C 1 Faraday = C mol -1 1 mole of a gas at 1 atmosphere pressure and 298 K occupies 24 dm 3. If the cell s efficiency is rated at 72%, what is the mass of hydrogen required? d) Suggest two advantages of using fuel cells for powering city buses. Suggest two ways in which the problem of storing hydrogen gas on a bus could be reduced. TOTAL / 16

25 TOPIC 17 Questionsheet 16 TEST QUESTION IV a) Explain the term standard electrode potential, Eê, as applied to the reduction of aqueous halogens.... [5] b) Eê/ volts I Fe 3+ (aq) + e - ¾ Fe 2+ (aq) II Cl 2 (aq) + 2e - ¾ 2Cl - (aq) III Br 2 (aq) + 2e - ¾ 2Br - (aq) IV I 2 (aq) + 2e - ¾ 2I - (aq) By referring to the electrode potential data, state how the oxidising power of the elemental halogens varies down Group 7. c) Write full ionic equations and calculate the e.m.f. of possible reactions between the following pairs of species. Comment on whether or not each reaction is feasible. Fe 2+ (aq) and Cl 2 (aq) Equation... [1] e.m.f.... [1] Feasibility... [1] Fe 2+ (aq) and I 2 (aq) Equation... [1] e.m.f.... [1] Feasibility... [1] Fe 3+ (aq) and I - (aq) Equation... [1] e.m.f.... [1] Feasibility... [1] d) In the light of your answer to c), comment on the stability of iron(iii) iodide. TOTAL / 16

26 TOPIC 17 Questionsheet 17 TEST QUESTION V a) Hydrogen is a common fuel for fuel cells. Suggest the missing words in this statement from these possibilities: chemical positive catalyst electrolyte redox electrochemical oxidant biomass carrier negative catalysis combustion reduction electrolysis renewable A fuel cell is an (A) energy conversion device. Fuel goes in on the (B) side and the (C) side takes the electron acceptor (the (D)). These two materials react in the presence of an (E). A fuel works by using a (F), such as a platinum alloy. Hydrogen acts as an energy (G) not an energy source. A=... B=... C=... D=... E=... F=... G=... [7] b) Hydrogen can be used as a fuel gas. Suggest two advantages and disadvantages if the gas is used for this purpose. Advantages... Disadvantages... c) The standard hydrogen electrode forms a cell with a half-cell of copper(ii) sulfate and copper under standard conditions. The standard electrode potential for Cu 2+ /Cu = +0.34V. Give an ionic half-equation to show the reaction at the reducing electrode The copper(ii) sulfate concentration is increased. State and explain the effect on the cell e.m.f. What is the standard cell e.m.f? TOTAL / 17