CAMOSUN COLLEGE Chemistry 121 Section 03 Instructor: H. M. Cartwright Page 1 Midterm 2 Examination, March 25 th, 2015, 6.30 p.m. Time: 1 ½ hours Short answers Name Student Number Answer all questions on the examination paper itself. If you need extra space, write on the reverse of the paper, indicating clearly within your answer to the question that you have done so. Attempt all the questions. The numbers in square brackets indicate the approximate number of marks available for each part of the question. Do not open the examination paper until you are instructed to do so.
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1. (a) What is an exothermic reaction? [1] One in which heat is released. (b) What sign does H have for an exothermic reaction? [1] Page 3 Minus 2. Explain the difference in thermodynamics between the system and the surroundings. System is the part on which we focus, generally the part where the reaction or process takes place; surroundings is everything else. 3. (a) What is a thermochemical equation? [1] A balanced chemical equation that also includes a statement of the enthalpy change during the reaction. (b) What is Hess's law? [2] Enthalpy change during a cycle is zero. Or enthalpy change in a reaction is independent of the course of the reaction. Or anything similar. (c) Calculate H for the formation of acetylene, C 2 H 2, as shown in the reaction 2C(s) + H 2 (g) C 2 H 2 (g) using the following data: [5] [3] C 2 H 2 (g) + 2 ½ O 2 (g) 2CO 2 (g) + H 2 O(l) C(s) + O 2 (g) CO 2 (g) H 2 (g) + ½ O 2 (g) H 2 O(l) H = 1299.6kJ H = 393.5 kj H = 285.8 kj - (1 st reaction) + 2( 2 nd reaction) + (3 rd ) = +226.8 kj (d) Acetylene is burnt in oxy-acetylene torches to give a very high temperature. Taking note of your answer to part (c) above, and the equations given in the question, give reasons why an oxy-acetylene flame reaches such a high temperature. [3]
(a) formation of acetylene is endothermic (just calculated), so the reverse reaction, its breakdown into the constituent elements, is exothermic; and (b) combustion of those elements is also exothermic, as the 2 nd and 3 rd equations show. Thus overall we'd expect a strongly exothermic reaction. Page 4 4. Explain what is meant by the following terms which are used in thermochemistry: (i) Open system Both matter and energy can enter or leave the system (ii) Closed system Energy can leave or enter the system, but matter cannot. (iii) Isolated system Neither energy nor matter can enter or leave the system 5. (a) What is a state function? [2] One whose value depends only upon the current state of the system, not on how it was prepared. (b) For each of the following quantities place a tick to indicate whether the quantity is, or is not, a State Function: [3] Quantity Is a State Function Is not a State Function Temperature Work Internal energy Entropy Heat [3]
Enthalpy 6. (a) Define the standard enthalpy of formation, H o f, of a compound. [2] Enthalpy change when 1 mole of a material in its standard state is formed from the stable form of the constituent elements in their standard states. Page 5 (b) Carbon nanotubes are an elemental form of carbon produced by an electrical discharge between carbon electrodes. Which ONE of the following statements about the standard enthalpy of formation of carbon nanotubes is true? [2] (i) H o f is zero because nanotubes are just another elemental form of carbon. (ii) H o f is zero because changing the subscripts of an elemental formula does not change the standard heat of formation. (iii) H o f is not zero because carbon nanotubes are not the most stable form of the element under standard conditions. (iv) H o f is not zero because there is a temperature change when nanotubes are formed. 7. For EITHER the 1 st law OR the 2 nd law of thermodynamics: (a) State the chosen law, indicating which law it is. [2] 1 st law: energy is conserved 2 nd law: entropy of the universe increases during any spontaneous process (b) Write the relevant equation for the law. [1] U = q + w S univ positive (c) Explain the meaning of the symbols in the equation. [2] U: change in internal energy q: heat absorbed by the system w: work done on the system S univ : change in entropy in the universe
8. (a) Explain what is meant by the specific heat of a material. [2] Amount of heat needed to raise the temperature of 1g of material by 1K. (b) 26 ml of 0.2 M HCl reacted with 8 ml of 0.6 M NaOH in a calorimeter. When the reaction was complete it was found that the temperature had risen by 1.83 K. You may assume that the calorimeter and liquids were all at the same temperature initially. Page 6 (i) Determine which, if either, of the reagents is present in excess. [2] Moles HCl = 0.2 x (26/1000) = 0.0052 moles Moles NaOH = 0.6 x (8/1000) = 0.0048 moles Thus HCl is in excess; NaOH is limiting reagent (ii) Assuming that the density of the solutions is 1 g ml -1, the specific heat of the solutions is 4.18 J K -1 g -1 and the heat capacity of the calorimeter is 41 J K -1, calculate the total heat released by the neutralisation. [2] 34mL of liquid in total = 34g Total heat capacity of liquid is 34 x 4.18 = 132.12 J / K Since the temperature change is 1.83K, total heart absorbed by liquid = 142.12 x 1.83 = 260.1 J Heat absorbed by calorimeter = 41 x 1.83 = 75 J Total heat absorbed = 260.1 + 75 = 335 J (iii) Calculate the enthalpy of the neutralisation reaction per mole of acid. [2] Only 0.0048 moles of acid is neutralised (see part (i)), so molar enthalpy is 335 J / 0.0048 moles = 69.8 kj / mole (iv) In a second experiment the HCl was replaced by 26 ml of 0.2M HNO 3 and the experiment repeated. What should the temperature rise be? Justify your answer. [2] 1.83K. Identity of the acid is immaterial, since the actual reaction is H + + OH - water (v) In a final experiment the HCl was replaced by 26 ml of 0.2M CH 3 COOH (acetic acid). Should the temperature rise be the same as that in part (iv) or different? Justify your answer. [2]
Temperature rise is still 1.83K. Acetic acid is weak, so produces only a low concentration of H +. However, these ions react completely with the available OH -, so more of the acetic acid dissociates to try to replenish the supply (think Le Chatelier). This continues until all the acid has dissociated. Since the amount of acid in this part of the question is the same as before, the total amount of reaction is the same and so the temperature rise will be the same also. Page 7 9. For each of the following processes predict whether the entropy of the system will increase or decrease; briefly explain your reasoning for each one: [3 x 2] O 3 (g) 1 ½ O 2 (g) Increase; more moles of gas in the product. Liquid sulphur (108 o C) Solid sulphur (108 o C) Decrease; solids have lower entropy than liquids 2HBr(g) Br 2 (g) + H 2 (g) Increase; the number of moles of gas is the same on both sides, but a mixture of gases is more random, and therefore has more entropy, than the same amount of a pure gas. 10. (a) What is a spontaneous process in thermodynamics? [2] One that proceeds without the need for an ongoing supply of energy (b) What thermodynamic parameter is used in thermodynamics to predict the spontaneity of the reaction? [1] Gibbs' Free Energy, G (c) Write down an equation for the change in Gibb's Free Energy, G, in terms of changes in entropy and enthalpy. [1] G = H T S (d) Write a balanced equation for the combustion of liquid dimethyl ether (C 2 H 6 O) in oxygen. [2] C 2 H 6 O (l) + 3O 2 (g) 2CO 2 (g) + 3H 2 O (g) (water formed in the gas phase since this is a combustion)
(e) What sign should S have for the combustion in part (d)? Justify your answer. [2] Positive; increase in the number of moles of gas (f) What sign should H have for the combustion in part (d)? Justify your answer. [2] Page 8 Negative; combustion is exothermic. (g) Taking note of your answers to parts (c), (e) and (f), over what temperature range would this reaction be spontaneous? [2] Change in Gibbs' energy is negative at all temperatures, so reaction is always spontaneous 11. Halogens can react with one another to form "interhalogen" compounds; the reaction between chlorine and brome is typical: Cl 2 (g) + Br 2 (g) 2ClBr(g) (a) Which ONE of the following statements is true when the reaction has reached equilibrium at a temperature of 80 o C when all species are gases? [2] (i) All reaction has come to a halt. (ii) The amounts of reactant and products are equal. (iii) The amount of reactant is zero. (iv) The concentration of chlorine and bromine is equal. (v) The rates of the forward and back reaction are equal. (b) The normal boiling point of ClBr is 5 o C, of Cl 2 is -34 o C, and of Br 2 is 59 o C. Suppose that the reaction above is carried out at 25 o C; would your conclusion about which of options (i)-(v) is true be different? Explain. No. At equilibrium the rates of forward and reverse reactions must be equal; if that were not the case, we wouldn't be at equilibrium. 12. (a) State Le Chatelier's Principle. [3] If a constraint, such as a change in concentration, temperature or pressure, is applied to a system in equilibrium, the position of equilibrium will move, if possible to try to nullify the constraint.
(b) Write down the expression for the equilibrium constant K c for each of the following reactions: [4] 2NH 3 (g) N 2 (g) + 3 H 2 (g) [N 2 ] [H 2 ] 3 / [NH 3 ] 2 Page 9 CH 3 COOH (aq) CH 3 COO (aq) + H + (aq) [ CH 3 COO ] [ H + ] / [ CH 3 COOH ] CH 3 COOH (aq) + H 2 O (l) CH 3 COO (aq) + H 3 O + (aq) [ CH 3 COO ] [ H 3 O + ] / [CH 3 COOH ] [CO 2 ] CaCO 3 (s) CaO (s) + CO 2 (g) (c) Predict whether the reaction given below is more likely to be exothermic or more likely to be endothermic (hint: consider the reverse reaction). Justify your choice. [2] 2CO 2 (g) + 3H 2 O(l) C 2 H 5 OH (l) + 3 O 2 (g) Endothermic; the reverse reaction is a combustion, which we can expect to be exothermic. To the right (d) The reaction shown in part (c) above is allowed to run to equilibrium. In which direction would the equilibrium shift if each of the following changes were made: [4 x 2] (i) An increase in the concentration of CO 2? (ii) An increase in the amount of liquid methanol? (Actually we mean ethanol ) No change, since the methanol is a liquid. However, some credit given for saying equilibrium would move to the left. To the right To the right (iii) An increase in the temperature? (iv) A decrease in the total pressure?
(v) The addition of a catalyst? No change; catalysts don't affect the position of equilibrium, just the rate at which it is reached. 13. Phosphorus trichloride reacts with chlorine in an equilibrium reaction to generate phosphorus pentachloride: Page 10 PCl 3 (g) + Cl 2 (g) PCl 5 (g) (a) Write down the expression for the equilibrium constant for the reaction, K c. [1] Prods/reactants as usual (b) At a certain temperature when the reaction had reached equilibrium, the concentrations of the three species were: [PCl 3 ] = 0.42 M [Cl 2 ] = 0.37 M [PCl 5 ] = 0.39M Calculate the value of the equilibrium constant. [1] Plug the values in to get K = 2.51 (c) In a second experiment at the same temperature, the initial concentrations were: [PCl 3 ] = 0.53M [Cl 2 ] = 0.53M [PCl 5 ] = 1.1M. Determine the concentrations of all three species in this second experiment once equilibrium has been achieved. [4] Note: the solutions of the quadratic equation ax 2 + bx + c = 0 are x = b± (b2 4ac) 2a Use ICE as normal. Note that if the change in concentration of phosphorus trichloride is x, that of the pentachloride is +x, not +2x. x is +0.1011 or -0.1011, depending upon whether you have chosen x to be the amount by which the concentration of product increases or decreases. Final concentrations are [PCl 3 ] = [Cl 2 ] = 0.6311M, [PCl 5 ] = 0.9989 M. 14. Explain the difference between Arrhenius and Bronsted-Lowry acids. Give one example of each type. [4] Arrhenius generates H + in water, B-L is proton donor. Examples of either HCl, ethanoic acid, etc.