Reaction Mechanisms (College Board AP® Chemistry): Free Response Questions

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.

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

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

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

 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

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​.

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.

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.

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.

Deduce the rate law based on the proposed mechanism.

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

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

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

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 ⇌ A_2​ (fast equilibrium)

2. Step 2: A₂ + B → C (slow)

Identify the intermediate in the mechanism.

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

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

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

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

1. Cl^. + O₃ $\stackrel{k_1}{\rightleftharpoons }$ ClO^. + O₂​ (fast)

2. ClO^. + O → Cl^. + O₂​ (slow)

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

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.

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

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

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.

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

The reaction between iodide ions (I^-) and persulfate ions (S₂O₈^2-) is often studied using a clock reaction to investigate kinetics. The reaction is given by:

2I^- (aq) + S₂O₈^2- (aq) → I₂ (aq) + 2SO₄^2- (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.

A study of the reaction produced the following experimental data:

i) Use the experimental data to determine the order of reaction with respect to S₂O₈^2-.

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).

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

The following mechanisms have been proposed for the reaction:

Mechanism 1:

1. I^- (aq) + I^- (aq) → I₂^2- (aq) (slow)

2.  I₂^2- (aq) + S₂O₈^2- (aq) → I₂ (aq) + 2SO₄^2- (aq) (fast)

Mechanism 2:

1. I^- (aq) + S₂O₈^2- (aq) → S₂O₈I^3- (aq) (slow)

2.  S₂O₈I^3- (aq) + I^- (aq) → I₂ (aq) + 2SO₄^2- (aq) (fast)

Mechanism 3:

1. I^- (aq) + S₂O₈^2- (aq)  → S₂O₈I^3- (aq) (fast)

2.  S₂O₈I^3- (aq) + I^- (aq) → I₂ (aq) + 2SO₄^2- (aq) (slow)

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

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.

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

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

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