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Question 1
The reaction between A, B and C was studied. 2A + B + C ➔ D The following sequence of steps is a proposed mechanism for the reaction. Step 1: 2A + C ➔ E Step 2: B + E ➔ D The general form of the rate equation for this reaction is: rate = k[A]l[B]m[C]n Step 1 is the slower step in the mechanism. |
a) | Deduce the values of l, m and n in the rate equation.
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b) | The rate equation for this reaction is: rate = k[A]2[C] |
What is the overall order of this reaction? A 0 B 1 C 2 D 3
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c) | Select the correct units for the rate constant of this rate equation? D mol2 dm-6 s-1 B mol dm-3 s-1 A mol-1 dm3 s-1 C mol-2 dm6 s-1
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d) | Draw a line of best fit on the graph below to show that is reaction is first order with respect to C.
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e) | Draw a line of best fit on the graph below to show that is reaction is zero order with respect to B.
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Question 2
Nitrogen dioxide is formed from the oxidation of nitrogen monoxide according to the equation: 2NO(g) + O2(g) ➔ 2NO2(g) The rate of reaction was measured at various concentrations of the two reactants. The results are shown in the table below. |
a) | Use the data in the table to deduce the order of reaction with respect to NO. Show your reasoning.
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b) | Use the data in the table to deduce the order of reaction with respect to O2. Show your reasoning.
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c) | Write the rate equation for the reaction betwen nitrogen monoxide and oxygen.
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d) | Use the data in the table to calculate the initial rate for experiment 4. Give your answer to 2 significant figures.
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e) | Use the results of experiment 2 to calculate the rate constant, k, for this reaction. Include the units of k.
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Question 3
Cyclopentadiene dimerises according to the equation: 2C5H6 ➔ C10H12 The graph below shows how the concentration of cyclopentadiene varies with time. |
a) | Draw a tangent to the curve when the concentration of cyclopentadiene is 0.06 mol dm-3.
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b) | The initial rate of reaction in this experiment has the value 1.92 x 10-5 mol dm-3 s-1. Use this value and the value obtained from the tangent in part a) to justify that the order of this reaction is 2 with respect to cyclopentadiene.
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Question 4
Nitrogen monoxide, NO, reacts with hydrogen, H2 via the equation: 2NO(g) + 2H2(g) ➔ N2(g) + 2H2O(g) The rate equation for this reaction is: rate = k[NO]2[H2] The graph below shows how the initial rate of reaction changes as the concentation of H2 varies. The concentration of NO is 0.50 mol dm-3. |
a) | Using values from the graph, calculate the rate constant, k, for this reaction. Give the units for the rate constant. Give your answer to 2 significant figures.
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b) | State the effect, if any, on the rate constant, k, if the temperature was lowered.
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c) | A three-step mechanism is proposed for the reaction: 2NO(g) + 2H2(g) ➔ N2(g) + 2H2O(g) Step 1: 2NO ➔ N2O2 fast Step 2: H2 + N2O2 ➔ N2O + H2O slow Step 3: .................................................. fast |
Write the equation for step 3.
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d) | Explain why this mechanism is consistent with the rate equation: rate = k[NO]2[H2]
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Question 5
The compound (CH3)3CBr reacts with sodium hydroxide as shown in the following equation. (CH3)3CBr + OH- ➔ (CH3)3COH + Br- This reaction was found to be first order with respect to (CH3)3CBr but zero order with respect to hydroxide ions. The following two-step process was suggested. Step 1: (CH3)3CBr ➔ (CH3)3C+ + Br- Step 2: (CH3)3C+ + OH- ➔ (CH3)3COH
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a) | State what is meant by the term rate-determining step.
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b) | Deduce the rate-determining step in this two-step reaction.
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c) | Write the rate equation for the reaction between (CH3)3CBr and OH-.
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d) | Use your answer to part c) to help you draw the mechanism for the reaction of (CH3)3CBr with hydroxide ions. Include the following:
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Question 6
Hydrogen peroxide reacts with iodide ions in acidic conditions as shown in the equation below. H2O2(aq) + 2I-(aq) + 2H+(aq) ➔ I2(aq) + 2H2O(l) A chemist investigates the rate of this reaction by measuring the time taken for a certain amount of iodine to be produced.
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a) | Outline a series of experiments that the chemist could have carried out using the initial rates method.
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b) | How could the results be used to show that the reaction is first-order with respect to both H2O2 and I-?
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c) | Suggest why initial rates of reaction are used to determine these orders rather than rates of reaction at other times during the experiments.
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d) | The rate equation for this reaction is: rate = k[H2O2(aq)][I-(aq)] In one of the experiments, the chemist reacted together:
The initial rate of this reaction is 1.14 x 10-3 mol dm-3 s-1. |
Calculate the rate constant, k, for this reaction. Give your answer to 3 significant figures and with units.
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e) | The chemist concluded that H+(aq) ions act as a catalyst. Explain why the chemist’s conclusion is not correct.
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f) | Complete the table below with the words 'increase', 'decrease', or 'stays the same' to show how the rate of this reaction and the rate constant, k, are effected by changes in conditions.
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Question 7
Dinitrogen pentoxide decomposes as shown by the equation: 2N2O5(g) ➔ 4NO2(g) + O2(g) |
The table below shows how the concentration of N2O5 varies with time at 40°C. |
a) | Plot a graph of the data in the table.
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b) | Use the graph in part a) to determine the half-life, in s, for the reaction.
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c) | What would be the effect, if any, on the half-life of this reaction of doubling the initial concentration of N2O5? Explain your reasoning.
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d) | How does the graph in part a) show that the reaction is first order with respect to N2O5?
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e) | Use the half-life of N2O5 at 315 K to calculate the rate constant for this reaction. Give the units for the rate constant. Give your answer to 3 significant figures.
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Question 8
The rate equation for a reaction is: rate = k[B] |
a) | In terms of the rate equation, explain why the rate increases with increasing temperature.
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b) | Explain qualitatively why doubling the concentration of B has a much smaller effect on the rate of reaction than doubling the temperature.
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Question 9
NO(g) acts as a catalyst in the oxidation of atmospheric sulfur dioxide. |
a) | State the meaning of the term homogeneous as applied to catalysts.
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b) | Give two equations to describe how NO(g) acts as a catalyst in this process.
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c) | Explain why NO(g) can be described as a catalyst in this reaction.
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d) | Iron catalyses the reaction of hydrogen and nitrogen to form ammonia in the Haber process: N2(g) + 3H2(g) ➔ 2NH3(g) |
Explain what type of catalysis is occuring in the Haber process.
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e) | Describe the stages in the catalytic production of ammonia by iron in the Haber process
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Question 10
W and Y react together to produce Z. 2W + Y ➔ Z The rate equation for a reaction is: rate = k[W]2[Y] |
a) | Deduce the overall effect on the rate of reaction when the concentrations of both W and Y are halved.
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b) | The rate of reaction is 7.35 x 10-6 mol dm-3 s-1 when the concentration of Y is 0.575 mol dm-3. |
Calculate the concentration of W, in mol dm-3, if the rate constant is 1.67 x 10-5 mol-2 dm6 s-1.
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