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AP®ChemistryCollege BoardExam QuestionsUnit 5: KineticsReaction Mechanisms

Reaction Mechanisms (College Board AP® Chemistry): Exam Questions

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Reaction Mechanisms

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1a1 mark

Nitrogen dioxide, NO2(g) , is produced as a by-product of the combustion of fossil fuels in internal combustion engines. At elevated temperatures NO2(g) decomposes according to the equation below.

2 NO2(g)  →  2 NO(g) + O2(g)

The concentration of a sample of NO2(g) is monitored as it decomposes and is recorded on the graph directly below. The two graphs that follow it are derived from the original data.

q6-june-2019---ap-chemistry

Explain how the graphs indicate that the reaction is second order.

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1b1 mark

Write the rate law for the decomposition of NO2 (g) .

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1c2 marks

Consider two possible mechanisms for the decomposition reaction.

 i) Is the rate law described by mechanism I shown below consistent with the rate law you wrote in part (b)? Justify your answer.

Mechanism I

Step 1:

NO2 (g) + NO2 (g) → NO (g) + NO3 (g)

slow

Step 2:

NO3 (g) → NO (g) + O2 (g)

fast

 ii) Is the rate law described by mechanism II shown below consistent with the rate law you wrote in part (b)? Justify your answer.

 Mechanism II

Step 1:

NO2 (g) + NO2 (g)  ⇌ N2O4 (g)

fast equilibrium

Step 2:

N2O4 (g) → 2 NO (g) + O2 (g)

slow

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2a1 mark

The reaction between nitrogen dioxide and carbon monoxide occurs according to the following equation:

NO2 (g) + CO (g) → NO (g) + CO2 (g) 

Using the following graph, determine the order of reaction with respect to NO2​.

q4a_16-1_ib_hl_medium_sq

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2b2 marks

A student proposes the following single-step mechanism for the reaction:

NO2 + CO → NO + CO2 ​(slow)

The student states that the rate law for the reaction is:

Rate = k [NO2]2 

Justify whether the student’s proposed mechanism is consistent with the given rate law.

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2c1 mark

Another student proposes the following two-step mechanism:

Step 1: NO2 + NO2 → NO + NO3    

Step 2: NO3 + 2CO → NO + 2CO2 

The student states that the rate law for the reaction is:

Rate = k [NO2]2 

Based on the proposed mechanism, identify the rate-determining step and justify your choice.

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3a1 mark

The dimerisation of ethanal (CH3CHO) in dilute alkaline solution to form 3-hydroxybutanal proceeds via the following mechanism:

Step 1: CH3CHO + :OH- → :CH2CHO + H2O (slow)
Step 2: CH3CHO + :CH2CHO → CH3CH(O:-)CH2CHO (fast)
Step 3: CH3CH(O:-)CH2CHO + H2O → CH3CH(OH)CH2CHO + :OH- (fast)

Identify the catalyst in the reaction mechanism. Justify your answer based on its role in the proposed steps.

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3b1 mark

Deduce the rate law based on the proposed mechanism.

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3c1 mark

The following data was collected for the reaction at a constant temperature.

Experiment

Initial [CH3CHO]

(mol dm-3)

Initial [OH-]

(mol dm-3)

Initial rate

(mol dm-3 s-1)

1

0.25

0.20

4.2 x 10-2

2

0.25

0.30

 

Calculate the initial rate for Experiment 2, assuming all conditions remain the same.

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3d1 mark

State the effect, if any, that increasing the concentration of a reactant would have on the value of the rate constant, k.

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4a1 mark

A reaction mechanism for the decomposition of dinitrogen pentoxide (N2O5) in chloroform is proposed:

  1. Step 1: N2O5 → NO2 + NO3 (slow)

  2. Step 2: N2O5 + NO3 → 3NO2​ + O2 (fast)

Identify the rate-determining step. Justify your answer.

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4b1 mark

Write the rate law for the reaction.

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4c1 mark

Determine the overall balanced equation for the reaction.

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4d1 mark

Explain how an increase in temperature affects the rate of this reaction.

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5a1 mark

A reaction occurs according to the following experimentally determined rate law:

Rate = k[A]2[B]

A student proposes the following two-step mechanism:

  1. Step 1: A + A ⇌ A2​ (fast equilibrium)

  2. Step 2: A2 + B → C (slow)

Identify the intermediate in the mechanism.

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5b1 mark

Explain whether the mechanism is consistent with the given rate law.

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5c1 mark

If the temperature is increased, describe how the rate constant k changes.

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5d1 mark

Predict the effect of doubling [B] while keeping [A] constant.

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1a1 mark

The decomposition of ozone (O3​) in the presence of chlorine radicals (Cl.) occurs via the following mechanism:

  1. Cl. + O3 rightwards harpoon over leftwards harpoon with k subscript 1 on top ClO. + O2​ (fast)

  2. ClO. + O → Cl. + O2​ (slow)

Identify the intermediate in the reaction mechanism and explain your reasoning.

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1b2 marks

i) Write the rate law for the overall reaction based on the rate-determining step.

ii) Explain why the concentration of Cl. can be considered constant in this reaction.

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1c1 mark

Predict how the rate would be affected if the concentration of O3​ were doubled. Justify your answer.

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2a2 marks

The reaction between hydrogen (H2​) and iodine (I2​) to form hydrogen iodide (HI) was investigated at 298 K, and the following data was obtained:

Experiment

[H2] (mol dm-3)

[I2] (mol dm-3)

Initial rate(mol dm-3 s-1)

1

0.0258

0.0137

6.43 x 10-22

2

0.0258

0.0274

1.29 x 10-21

3

0.0516

0.0137

1.29 x 10-21

The following three-step mechanism has been proposed for the reaction:

Step 1: I2 (g) ⇋ 2I (g) (fast equilibrium)

Step 2: H2 (g) + I (g) ⇋ H2I (g) (fast equilibrium)

Step 3: H2I (g) + I (g) → 2HI (g) (slow)

Determine the rate law using the experimental data.

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2b2 marks

Justify how the proposed mechanism supports the experimentally determined rate law.

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3a2 marks

The reaction between iodide ions (I-) and persulfate ions (S2O82-) is often studied using a clock reaction to investigate kinetics. The reaction is given by:

2I- (aq) + S2O82- (aq) → I2 (aq) + 2SO42- (aq)

i) Identify the species being oxidized. Justify your answer using oxidation states.

ii) Identify the species being reduced. Justify your answer in terms of electron transfer.

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3b3 marks

A study of the reaction produced the following experimental data:

Experiment

[S2O82-] (mol dm-3)

[I-] (mol dm-3)

Initial rate(mol dm-3 s-1)

1

0.25

0.10

8.0 x 10-3

2

0.10

0.10

3.2 x 10-3

3

0.20

0.30

1.92 x 10-2

i) Use the experimental data to determine the order of reaction with respect to S2O82-.

ii) Use the experimental data to determine the order of reaction with respect to I−.

iii) Write the overall rate law for this reaction based on your answers to (i) and (ii).

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3c2 marks

Determine the rate constant k using data from Experiment 1. Include units in your answer.

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3d3 marks

The following mechanisms have been proposed for the reaction:

Mechanism 1:

  1. I- (aq) + I- (aq) → I22- (aq) (slow)

  2.  I22- (aq) + S2O82- (aq) → I2 (aq) + 2SO42- (aq) (fast)

Mechanism 2:

  1. I- (aq) + S2O82- (aq) → S2O8I3- (aq) (slow)

  2.  S2O8I3- (aq) + I- (aq) → I2 (aq) + 2SO42- (aq) (fast)

Mechanism 3:

  1. I- (aq) + S2O82- (aq)  → S2O8I3- (aq) (fast)

  2.  S2O8I3- (aq) + I- (aq) → I2 (aq) + 2SO42- (aq) (slow)

For each mechanism, evaluate whether it is consistent with the experimentally determined rate law. Justify your answers.

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4a1 mark

A reaction mechanism is proposed:

Step 1: A + B ⇌ C (fast equilibrium)

Step 2: C + D → E (slow)

Identify the intermediate in this reaction mechanism.

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4b1 mark

Write an expression for [C] in terms of [A] and [B] using the pre-equilibrium assumption.

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4c1 mark

Deduce the overall rate law for the reaction using the pre-equilibrium approximation. Assume that D is in excess.

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4d1 mark

Explain why Step 2 is the rate-determining step in this reaction mechanism.

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