Reaction Kinetics: Chemistry 9701 (Cambridge International AS & A Level)

A chemist studying a gas-phase reaction adds a small amount of a solid catalyst to the reaction mixture. The rate of reaction increases substantially, and the mass of the solid catalyst is unchanged at the end of the reaction.

Which statement correctly explains why the catalyst increases the rate of reaction?

A student investigates the reaction between small pieces of zinc metal and excess dilute sulfuric acid: $\text{Zn(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{H}_2\text{(g)}$ The flask is placed on a balance, and the loss in mass (due to escaping hydrogen gas) is recorded at regular time intervals.

(a) Explain, in terms of collision theory, why using a more concentrated solution of sulfuric acid (with the same volume of acid and the same mass of zinc) increases the initial rate of reaction. [2]

(b) The student repeats the experiment using the original concentration of acid, but at a temperature $20\text{ °C}$ higher than before. Using ideas about the Maxwell–Boltzmann distribution of molecular energies, explain why this temperature increase produces a much larger increase in rate than doubling the acid concentration would. [4]

(c) Describe how the shape of the Maxwell–Boltzmann distribution curve for the acid particles changes when the temperature is raised, and explain how this change relates to the activation energy, $E_A$, of the reaction. [3]

(a) State what is meant by a homogeneous catalyst and by a heterogeneous catalyst, in terms of the physical state (phase) of the catalyst relative to the reactants. [2]

(b) A gas-phase reaction between two colourless gases, $\text{P(g)}$ and $\text{Q(g)}$, occurs extremely slowly at $400\text{ K}$. When a small amount of a transition metal is added as a solid, gas particles adsorb onto its surface and the rate increases dramatically; the mass of the metal is unchanged at the end of the reaction. State whether this catalyst is acting homogeneously or heterogeneously, and explain, in terms of activation energy and reaction pathway, why the rate increases. [4]

(c) The same solid catalyst is used in two further experiments, each using the same total mass of metal: in one, the metal is a single, large block; in the other, the same mass is used as a very fine powder. State and explain which of these two experiments gives the greater initial rate of reaction. [2]

In an investigation into the kinetics of the reaction between two aqueous reactants, A and B, $\text{A(aq)} + 2\text{B(aq)} \rightarrow \text{C(aq)} + \text{D(aq)}$ the initial rate of reaction was measured for three different combinations of starting concentrations, at the same constant temperature. The results are shown below.

(a) Using Experiments 1 and 2, deduce the order of reaction with respect to A. Using Experiments 2 and 3, deduce the order of reaction with respect to B. Show your reasoning in each case. [3]

(b) Write the rate equation for this reaction, and state the overall order of reaction. [2]

(c) Calculate the rate constant, $k$, for this reaction at this temperature, including its units. [3]

(d) Calculate the initial rate of reaction, in mol dm⁻³ s⁻¹, for an experiment in which $[\text{A}] = 0.30\text{ mol dm}^{-3}$ and $[\text{B}] = 0.15\text{ mol dm}^{-3}$. [2]

A colourless compound, X, decomposes in aqueous solution. A student follows the concentration of X over time, obtaining the results below.

(a) Use the data to show that this reaction is first order with respect to X, explaining how the data supports this conclusion. [3]

(b) Calculate the rate constant, $k$, for this reaction, in $\text{s}^{-1}$. [3]

(c) The reaction is believed to occur by the following two-step mechanism:

Step 1: $\text{X} \rightarrow \text{Y} + \text{Z}$ (slow)

Step 2: $\text{Y} + \text{X} \rightarrow 2\text{W}$ (fast)

Deduce, with a reason, which of these two steps is the rate-determining step, and explain how this is consistent with the reaction being first order overall in X. [3]

The concentration of a reactant, R, is followed over time for three separate reactions, each involving only R (all other reactants, if any, are in large excess so their concentration does not change).

• Reaction 1: a graph of $[\text{R}]$ against time is a straight line with a constant negative gradient. $[\text{R}]$ falls by the same amount in every equal time interval.
• Reaction 2: a graph of $[\text{R}]$ against time is a curve for which the time taken for $[\text{R}]$ to halve is the same at every stage, regardless of the concentration at the start of that stage.
• Reaction 3: a graph of $[\text{R}]$ against time is a curve for which the time taken for $[\text{R}]$ to halve becomes progressively longer as the reaction proceeds.

Which row correctly gives the order of reaction with respect to R for Reactions 1, 2 and 3?

An investigation into the kinetics of a reaction between three aqueous reactants, A, B and C, $\text{A(aq)} + \text{B(aq)} + \text{C(aq)} \rightarrow \text{D(aq)}$ measured the initial rate of reaction for four different combinations of starting concentrations, at the same constant temperature. The results are shown below.

(a) Using an appropriate pair of experiments in each case, deduce the order of reaction with respect to A, with respect to B, and with respect to C. Show your reasoning. [3]

(b) Write the rate equation for this reaction, and state the overall order of reaction. [2]

(c) Calculate the rate constant, $k$, for this reaction at this temperature, including its units. [3]

(d) Calculate the initial rate of reaction, in mol dm⁻³ s⁻¹, for an experiment in which $[\text{A}] = 0.25\text{ mol dm}^{-3}$, $[\text{B}] = 0.50\text{ mol dm}^{-3}$ and $[\text{C}] = 0.30\text{ mol dm}^{-3}$. [2]

A reaction between two aqueous compounds, E and F, produces G and H: $\text{E(aq)} + 2\text{F(aq)} \rightarrow \text{G(aq)} + \text{H(aq)}$ The initial rate of this reaction was measured for three combinations of starting concentrations, at the same constant temperature.

(a) Deduce the order of reaction with respect to E and with respect to F, explaining your reasoning. [3]

(b) Write the rate equation for this reaction, and state the overall order of reaction. [1]

Two mechanisms have been proposed for this reaction, each consisting of two steps:

Mechanism 1: Step 1: $\text{E} \rightarrow \text{I}$ (slow) Step 2: $\text{I} + 2\text{F} \rightarrow \text{G} + \text{H}$ (fast)

Mechanism 2: Step 1: $\text{E} + \text{F} \rightarrow \text{I}$ (slow) Step 2: $\text{I} + \text{F} \rightarrow \text{G} + \text{H}$ (fast)

(c) Deduce, with reference to the rate equation from part (b), which of these two mechanisms is consistent with the experimental data. Explain your reasoning, including why the other mechanism is not consistent. [3]

(d) Identify the role of species I in the mechanism you chose in part (c), and explain why F, although a reactant in the overall equation, does not appear in the rate equation. [2]

A chemist compares the reaction pathway for the same exothermic reaction, once without any catalyst present and once with a catalyst present, at the same constant temperature.

Which statement correctly compares the two pathways?

Hydrogen gas and iodine vapour react together in a sealed, rigid container at constant temperature: $\text{H}_2\text{(g)} + \text{I}_2\text{(g)} \rightarrow 2\text{HI(g)}$

(a) Explain, in terms of collision theory, why increasing the pressure inside the container (by reducing its volume, with the same amounts of each gas present and no change in temperature) increases the rate of this reaction. [3]

(b) A second, identical mixture of $\text{H}_2\text{(g)}$ and $\text{I}_2\text{(g)}$, at the original pressure, is instead heated to a higher temperature. State two distinct ways in which this temperature rise increases the rate of reaction that the pressure increase in part (a) does not, explaining each briefly. [3]