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Question 1 (12 marks) (a) There are several different types of chemical reactions and more than one way of classifying them. The following table lists some chemical equations representing different reactions. Chemical equation A. 2HCl(aq) + Na2CO3(aq) →2NaCl(aq) + H2O(l) + CO2(g) B. Cu(NO3)2(aq) + Mg(s) → Mg(NO3)2(aq) + Cu(s) C. Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq) D. BaCO3(s) → BaO(s) + CO2(g) i. The signs for recognising when a chemical reaction has occurred are listed in the table below. Sometimes only one of these things happens, at other times several things happen. Identify one chemical equation in the previous table that could produce the observed change. (A chemical equation could be used more than once). (3 marks) Observations Chemical equation A precipitate formed Gas bubbling in solution A new solid formed ii. Select the chemical equation as being a precipitation reaction. (2 marks) Write the net ionic equation for this reaction. List the spectator ions. (b) Three beakers, labelled A, B and C, are placed on a student’s bench. The student adds 1.0 M MgSO4 solution to beaker A and 1.0 M ZnSO4 to beaker B. She then drops a piece of zinc into beaker A and a piece of magnesium into beaker B. The beakers are examined after an hour and the student’s observations are included in the table shown. Next for beaker C, she drops a piece of zinc in 0.10 M H2SO4 solution and records her observation immediately. Reactants Observations Beaker A Zinc metal in 1.0 M MgSO4 solution No change Beaker B Magnesium metal in 1.0 M ZnSO4 solution Magnesium looks to be disintegrating; pieces of solid are forming Beaker C Zinc metal in 0.10 M H2SO4 solution Zinc dissolves in the acid very quickly. Bubbles form, and after they are collected in a large test tube and ignited, a loud “pop” can be heard. i. Look at beaker A and B. Which metal is the more reactive? Explain your answer with reference to the experiment. (2 marks) ii. What is the limiting reactant in beaker A? (1 mark) iii. Write a balanced equation for the reaction occurring in beaker B. (1 mark) iv. What gas was produced in beaker C? (1 mark) v. What is the excess reactant in beaker C? (1 mark) vi. List one specific safety precaution the student should follow when performing the experiment in beaker C. (1 mark) Question 2 (15 marks) The concentration of sodium hydroxide in a sample of drain cleaner was analysed by titration. A student took 25.00 mL aliquot of the commercial product and added it into a 500 mL volumetric flask. Some distilled water is added to mix the sample and more distilled water is continuously added to the 500 mL calibration line. The flask was then closed with a stopper and inverted multiple times to make sure that the solution is homogeneous. This is the diluted drain cleaner solution. The student was provided with a 0.0109 M standardized hydrochloric acid solution. She rinses a clean burette with some distilled water and then filled it up with hydrochloric acid. She then takes a 20.00 mL volumetric pipette, washed it with some distilled water and then used it to measure the diluted drain cleaner. Before she started the titration her teacher asked her to explain what she did. Upon listening, the teacher told her that she made two errors in her technique. (a) List the errors and explain how these would have affected the calculated concentration of NaOH in the drain cleaner. (4 marks) Error 1: Error 2: (b) The errors were rectified, and the student followed the correct procedure. The diluted drain cleaner was accurately transferred to a conical flak and three drops of indicator were added to it. i. The student had three acid-base indicators available for use. Explain, with proper reasoning, which indicator is the most suitable to use for this analysis.(2 marks) Indicator pH range Colour in acid Colour at end point of titration Colour in base Methyl orange 3.1-4.4 red orange yellow Phenolpthalein 8.0-10.0 colourless Light pink Bright pink Bromothymol blue 6.0-7.6 yellow green blue ii. What colour change will she observe is she overshoots the titration end point? Explain why she will see this colour.(2 marks) (c) The student performed five titrations and recorded the following information in a data table. Complete the table by calculating the volume of each titre of hydrochloric acid. (1 mark) Titration 1 2 3 4 5 Initial burette reading (mL) 0.00 1.70 24.90 1.05 5.80 Final burette reading (mL) 23.80 24.90 48.00 24.80 29.10 Volume of HCl used in titration (mL) (d) Use the concordant results to calculate the average titre used. (1 mark) (e) The reaction between sodium hydroxide and hydrochloric acid is: NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l) i. Calculate the moles of hydrochloric acid used in the titration. (1 mark) ii. Calculate the moles of sodium hydroxide in the 20.00 mL diluted drain cleaner solution. (1 mark) iii. Calculate the concentration of sodium hydroxide in the diluted drain cleaner solution, in mol/L. (1 mark) iv. Calculate the concentration of sodium hydroxide in the original undiluted drain cleaner solution, in mol/L. (1 mark) (f) If 0.02 M HCl was used instead of 0.0109 M, would you need more or less of HCl for the titration? Explain with proper reasoning. (1 mark) Question 3 (12 marks) First year Chemistry students were studying factors which can disturb an equilibrium system. As a laboratory activity, they were asked to observe the effect of disturbances (changes) made to two equilibrium systems. System 1 The first system they observed consisted of a bromothymol blue solution. Bromothymol blue is a weak organic acid with a complex formula. Its formula can be abbreviated to HBb. Chemical Equation for the Equilibrium HBb(aq) {yellow} H+(aq) {colourless} + Bb- (aq) {blue} The following tests were carried out on separate samples of the solution. Test 1: Two drops of 0.1 M hydrochloric acid, HCl, is added to the solution and mixed. Test 2: Two drops of 0.1 M sodium hydroxide, NaOH, is added to the solution and mixed. Test 3: Several mL of deionized water is added to the solution and mixed. Complete the following table: Observed change (predict the colour change to the system) Explain what happened chemically and the position of equilibrium shift (left or right) due to the change Test 1 Test 2 Test 3 (6 marks) System 2: The second system consists of a saturated copper(II) chloride solution which is green, while its diluted aqueous solution has a pale blue colour. Note it is an endothermic reaction. Chemical Equation for the Equilibrium CuCl42-(aq) {green} + 4 H2O(l) + heat Cu(H2O)42+(aq) {pale blue} + 4 Cl-(aq) {colourless} The following tests were carried out on separate samples of the solution. Test 1: A small quantity of solid calcium chloride, CaCl2 is added to the solution and mixed to dissolve the solid. Test 2: A tube of the solution is placed in a hot-water bath. i. Complete the following table: Observed change (predict the colour change to the system) Explain what happened chemically and the position of equilibrium shift (left or right) due to the change Test 1 Test 2 (4 marks) ii. For test 2 explain how change in temperature affect the rate of reaction in terms of collision theory. (2 marks) Question 4 (8 marks) (a) The following list contains a number of different chemical substances: CH3COOH H2PO4- NH3 CH4 Cl2 NaCl CO2 LiOH NH4+ H2O Choose a substance from the above list to fill in the table (a substance from the list can only be used once). (2 marks) a strong base a conjugate acid/base pair An amphiprotic substance A polyprotic acid (b) A large beaker contains 40 mL of hydrochloric acid with a pH of 0. i. What is the hydrogen ion concentration of the solution? ((1 mark) ii. A further 160 mL of water is added to the beaker. Write a balanced chemical equation for the reaction occurring between the acid and water. (1 mark) iii. Is water acting as an acid, a base or a neutral substance in this reaction? Use Lowry–Bronsted definitions to explain your answer. (1 mark) iv. Calculate the concentration of the diluted acid. (1 mark) v. Calculate the pH of the diluted acid. Give answer correct to three decimal places. (1 mark) vi. Consider the original HCl solution of pH 0. How many times do you need to dilute the solution to increase it’s pH to 2? (1 mark) Question 5 (8 marks) (a) Carbonic acid (H2CO3) and bicarbonate ion (HCO3-) are a conjugate acid-base pair. For each of the following compounds/ions, would the carbonic acid or bicarbonate be more likely to react with them and serve as a buffer to prevent pH change? (2 marks) HBr NaOH NH4+ HCl (b) An experiment was conducted to observe the pH reading of a series of solutions as shown below. Solution Observed pH 0.01 M HCl 1.00 0.01 MCH3COOH 2.85 0.01 M CH3COONa 7.50 0.01 M NaOH 13.01 CH3COOH/CH3COO- buffer 4.58 5 mL buffer + 1 ml 0.01 M HCl 4.09 5 mL buffer + 1 mL 0.01 M NaOH 5.06 i. In the experiment HCl, CH3COOH and CH3COONa all have the same concentration but their pH values are different. Explain, with proper reasoning, which of these three is the strongest acid which makes its pH value different from the others. (2 marks) ii. The experimental results show CH3COOH/CH3COO- buffer in action. Explain with appropriate equations how the buffer has stabilized pH of the acid, HCl, and the base, NaOH. (4 marks)

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