SIR MICHELANGELO REFALO CENTRE FOR FURTHER STUDIES VICTORIA GOZO

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1 SIR MICHELANGELO REFALO CENTRE FOR FURTHER STUDIES VICTORIA GOZO Page 1 of 7 Half Yearly Exam 2013 Subject: Chemistry 1 st Year Level: Advanced Time: 3 hrs Answer SEVEN (7) questions. All questions carry equal marks. You are reminded of the importance of clear presentation in your answers, and the use of good English. Important information: A Periodic Table is attached on the last page. 1 mole of gas at s.t.p. occupies a volume of cm 3 ; R = 8.31 J K -1 mol This question is about energetics. a) Use the following information to draw the Born-Haber cycle for potassium bromide, KBr, including state symbols and the enthalpy values for each step. Hence calculate the lattice enthalpy of KBr. (6 marks) Process Enthalpy Change / kj mol -1 Enthalpy of sublimation of potassium Enthalpy of vaporisation of bromine (Br (l) Br (g)) 29.8 Enthalpy of ionisation of potassium Enthalpy of atomisation of bromine First electron affinity of bromine -324 Enthalpy of formation of KBr b) 200 cm 3 of mol dm -3 sodium hydroxide were mixed with 200 cm 3 of mol dm -3 hydrochloric acid in a calorimeter. The initial temperature of both solutions was 22.0 o C. The temperature after mixing the two solutions reached a maximum of 25.4 o C. Given that the specific heat capacity of the final solution was 4.18 J cm -3 K -1, calculate the enthalpy of neutralisation of sodium hydroxide with hydrochloric acid. c) The standard enthalpy of solution of sodium sulfate, Na 2 SO 4, is s kj mol -1 and its lattice enthalpy is l kj mol -1. The standard enthalpies of hydration of sodium and sulfate ions are a and b kj mol -1 respectively. Draw a Hess diagram which relates together all the mentioned enthalpy changes, and hence give an equation relating a, b, s and l. d) Use the following data to calculate the standard enthalpy of formation of propanone (CH 3 COCH 3 ): H θ c for CH 3 COCH 3 (l) = kj/mol ; H θ c for C (graphite) = kj/mol H θ c for H 2 (g) = kj/mol

2 2. This question is about chemical equilibrium. a) State what is meant by dynamic equilibrium. b) The equation for the main reaction in the Haber process is N 2 (g) + 3H 2 (g) 2NH 3 (g) H = -92 kj/mol Use this information to explain the effect, if any, on the equilibrium position and yield of ammonia of: (i) Increasing the pressure at constant temperature. (1.5 marks) (ii) Increasing the temperature at constant pressure. (1.5 marks) c) Lead(II) iodide, PbI 2, dissociates in water as follows: PbI 2 (s) + aq Pb 2+ (aq) + 2I - (aq) (i) Write an expression for the solubility product of PbI 2, and state its units. (ii) Write an expression which relates the solubility product of PbI 2 to its molar solubility (denote the molar solubility of PbI 2 as s). d) When a sample of NO 2 is heated to 1000 o C, it partly decomposes to form NO and O 2, according to the following equation: 2NO 2 (g) 2NO (g) + O 2 (g) The equilibrium constant K p at 1000 o C has a value of 160 atm. If the partial pressure of O 2 at equilibrium is atm, calculate the equilibrium pressures of NO and NO 2. e) Consider the following equilibrium: N 2 O 4 (g) 2NO 2 (g) (i) 0.37 moles of N 2 O 4 were introduced in a gas reaction vessel of volume 3.0 dm 3 and heated to a temperature of 70 o C. Calculate the value of the equilibrium constant K c at this temperature if 50% of the N 2 O 4 was found to be dissociated at equilibrium. (3.5 marks) (ii) State, giving a reason, how you would expect the value of K c for the above equilibrium to change, if at all, on increasing the pressure inside the reaction vessel. (1.5 marks) 3. This question is about nuclear chemistry. a) Define the term relative atomic mass. (1.5 marks) Page 2 of 7

3 b) A sample of lithium was found to have the following composition: Isotope 6 Li 7 Li Relative abundance (%) Calculate the relative atomic mass of this sample of lithium, giving your answer to two decimal places. c) The first three steps of the uranium-238 decay series are: U α + Th β + Pa β+ X (i) What is the atomic number of Thorium (Th)? (ii) What is the atomic number and mass number of Protactinium (Pa)? (iii) Identify and name element X. (0.5 marks) (0.5 marks) d) Explain the meaning of the term half-life of a radioactive decay process. e) Describe fully the principle of the carbon dating technique for certain archaeological objects. Include nuclear equations where relevant. f) Give a brief description, including a fully labelled diagram, to show the principle of operation of a mass spectrometer. (5 marks) g) The following are the structure of a compound and its associated mass spectrum: Suggest the identity of the chemical species responsible for the peaks at m/z = 29, 57 and 86. (1.5 marks) 4. This question is about bonding. Discuss briefly each of the following statements, drawing structures if they aid in the explanation: a) The crystalline structure and co-ordination number of NaCl are different from those of CsCl. Page 3 of 7

4 b) The electronegativity of elements decreases in going down a group of the Periodic Table. c) Silicon(IV) oxide (SiO 2 ) is a solid at room temperature, while CO 2 is a gas at the same temperature. d) Both graphite and diamond have macromolecular structures. However, graphite is a good conductor of electricity, whereas diamond is not. e) Phosphorus can form PCl 5, while nitrogen cannot form NCl 5. f) The structure of NH 3 is different from that of BH 3. As a result, there is a difference in the angular separation between bonds in both compounds. 5. This question is about gases. a) Ammonium nitrate, NH 4 NO 3, decomposes explosively on heating to form N 2 O and H 2 O, according to the following equation: NH 4 NO 3 (s) N 2 O (g) + 2H 2 O (g) What is the total volume of gaseous products at 110 o C and 1.01 x 10 5 Pa when 10.0 g of ammonium nitrate decomposes completely? Give your answer in units of dm 3. (4 marks) b) Discuss two assumptions which are made in deducing the ideal gas equation and which do not apply for real gases. Explain why under certain conditions, the assumptions break down and cause gases to deviate from ideal behaviour. (4 marks) c) When a gas syringe is used to measure the relative molecular mass of ethanoic acid, CH 3 COOH, it is found that the value varies from 60 to 120, depending on the temperature used in the experiment. Account for this finding. d) Sketch two graphs on the same set of axes to show the distribution of speed of gas molecules at two different temperatures T 1 and T 2, where T 2 > T 1. e) State Dalton s Law of Partial Pressures. f) A mixture of gases at a total pressure of 9.60 x 10 4 Pa contains 0.40 moles nitrogen, 0.35 moles carbon dioxide and 0.25 moles oxygen. Calculate the partial pressure of carbon dioxide in the mixture. 6. This question is about acids and bases. a) Explain the terms acid and conjugate base as used in the Bronsted-Lowry theory of acids and bases. Give an example of an acid and its conjugate base. Page 4 of 7

5 b) Calculate the ph of a 2.00 mol dm -3 solution of ethanoic acid, CH 3 COOH, given that its acid dissociation constant is 1.75 x 10-5 mol dm -3. c) Carbonic acid, H 2 CO 3, is a weak acid having a pk a value of Using the appropriate calculation(s), state whether ethanoic acid or carbonic acid is the weaker acid. d) Calculate the ph of a solution containing 1.00 mol dm -3 of ammonia and mol dm -3 of ammonium chloride. K b for ammonia is 1.80 x 10-5 mol dm -3. (4 marks) e) Sketch a ph vs. volume of alkali titration curve for the titration of aqueous ammonia and hydrochloric acid, and give (with a reason) an estimate of the ph at the equivalence point of this titration. f) Use the titration curve in (e) to explain whether phenolphthalein or methyl orange would be a suitable indicator for the titration of aqueous ammonia and hydrochloric acid. State the colour change which would occur at the end point of this titration when the alkali is placed in the burette. 7. This question is about quantum chemistry and atomic structure. a) Define the term first ionisation energy. Support your definition with a suitable equation. b) Sketch a graph showing the variation in first ionisation energy for the elements Li to Ne. (Note that the succession of elements in period 2 is Li, Be, B, C, N, O, F, Ne) c) Describe, giving suitable reasons, the general trend in the above graph and any exceptions to this trend. d) State, giving a reason, whether you would expect the first ionisation energy of sodium to be more or less endothermic than that of lithium. e) From the above mentioned elements in period 2 of the Periodic table, which one would you expect to have the most endothermic second ionisation energy? Give TWO reasons to support your answer. (1.5 marks) f) Define the term orbital. g) The chemical species F -, O 2- and Na + are isoelectronic. (i) Explain the meaning of isoelectronic species. (ii) Write down the electronic configuration, in terms of orbitals, for these species. h) The successive ionisation energies of element X are shown in the following table: Ionisation 1 st 2 nd 3 rd 4 th 5 th energy/kj mol Page 5 of 7

6 (i) State, with a complete explanation, in which group of the Periodic table element X is found. (1.5 marks) (ii) Given that X is found in period 3 of the Periodic table, write the electronic configuration, using the electrons-in-boxes method, of the outer shell electrons of element X. 8. This question is about redox reactions. a) Write down the oxidation number of the underlined atom(s) in each of the following chemical species: (i) HBr (ii) CrCl 3 (iii) NH 4 NO 3 2- (iv) SO 3 - (v) NO 2 b) Balance the following redox equations: (i) MnO H + + H 2 S S + Mn 2+ + H 2 O (ii) Cr 2 O H + + SO 3 2- SO Cr 3+ + H 2 O (iii) H 2 O 2 + Cr 3+ + OH - CrO H 2 O (5 marks) c) An experiment was carried out to determine the value of n in the formula Na 2 S 2 O 3.nH 2 O (hydrated sodium thiosulfate). A redox titration was performed. 4.22g of Na 2 S 2 O 3.nH 2 O were dissolved in cm 3 of water in a volumetric flask cm 3 of mol dm -3 iodine solution were pipetted into a conical flask, and this solution was titrated against the thiosulfate solution from a burette cm 3 of the thiosulfate solution were required. Determine the value of n. (7 marks) 9. This question is about redox equilibria. a) Values of electrode potentials, E, can vary with the experimental conditions used. State two specific conditions that must be fixed when determining standard electrode potentials. b) Two standard electrode potentials are given below: Br 2 (aq) + 2e - 2Br - (aq) E θ = V Co 2+ (aq) + 2e - Co (s) E θ = V (i) State the name of the reference electrode against which the above standard electrode potentials were measured. Give the value of E θ for this electrode. (ii) Draw a fully labelled diagram of a galvanic cell based on the above half-reactions, labelling the anode and cathode and direction of electron flow when the cell is operating. (iii) Write the cell statement for the galvanic cell. (iv) Calculate the e.m.f. of the cell. Page 6 of 7

7 (v) Write the complete equation, including state symbols, for the redox reaction which occurs when the cell is in use. (vi) Use the E θ values to explain whether Br - or Co is the stronger reducing agent. (vii) Use the E θ values to explain whether Br 2 or Co 2+ is the stronger oxidising agent. c) A student suggests that an acidified solution of iron(ii) chloride is unstable in the presence of atmospheric oxygen since both the cation and anion present may be oxidised. Use the standard electrode potentials below to fully discuss whether the student s statement is correct (include equations if relevant): Half reaction E θ (Volts) Cl 2 (g) + 2e - 2Cl - (aq) O 2 (g) + 4H + (aq) + 4e - 2H 2 O (l) Fe 3+ (aq) + e - Fe 2+ (aq) (4 marks) Page 7 of 7

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