Reaction Kinetics

The rate of the reaction is measured in $\text{mol dm}^{-3} \text{ s}^{-1}$. State the units of $k$.

In experiment 1, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are equal.

The initial rate of the reaction in experiment 1 is $2.57 \times 10^{-6} \text{ mol dm}^{-3} \text{ s}^{-1}$.

Calculate the initial concentration of $\text{NO}$.

Show your working.

In experiment 2, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are both ten times greater than the initial concentrations used in experiment 1.

Calculate the initial rate of the reaction in experiment 2.

The graph of $[\text{Cl}_2]$ against time shows that the reaction has a constant half-life, $t_{1/2}$.

Explain this observation.

Under the conditions used in experiment 3, the value of the rate constant is 105.6.

Show that $t_{1/2}$ of $[\text{Cl}_2]$ is $6.56 \times 10^{-3} \text{ s}$ under these conditions.

Calculate the time taken, in s, for $[\text{Cl}_2]$ to fall to $1.25 \times 10^{-5} \text{ mol dm}^{-3}$ in experiment 3.

Sulfur dioxide, $\text{SO}_2$, reacts very slowly with oxygen in the atmosphere, forming sulfur trioxide, $\text{SO}_3$. This reaction is much faster in the presence of $\text{NO}$.

Explain the role of $\text{NO}$ in this process.

Include chemical equations in your answer.

Explain what is meant by order of reaction.

Complete Table 4.1.

Table 4.1

Sketch a line on Fig. 4.1 to show the relationship between $[\text{I}^-]$ and time.

Two experiments are carried out to measure the rate of reaction 2.

In the first experiment, the initial rate is measured starting with known concentrations of $\text{NO}_2$ and $\text{O}_3$. In the second experiment, the concentrations of $\text{NO}_2$ and $\text{O}_3$ are both increased by a factor of four.

Predict how the initial rate for reaction 2 would change.

The reaction mechanism for reaction 2 has two steps.

Define rate-determining step.

Suggest equations for the two steps of the reaction mechanism for reaction 2.

step 1 ................................................................................................................................

step 2 ................................................................................................................................

Use Fig. 4.2 to calculate a value for the rate constant, $k$, for this decomposition.

Use your answer to (c)(i) to calculate the half-life, $t_{\frac{1}{2}}$, in seconds, for this decomposition.

The rate of the reaction is measured in $\text{mol dm}^{-3}\text{ s}^{-1}$. State the units of $k$.

units of $k$ = ..............................

In experiment 1, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are equal.

The initial rate of the reaction in experiment 1 is $2.57 \times 10^{-6}\text{ mol dm}^{-3}\text{ s}^{-1}$.

Calculate the initial concentration of $\text{NO}$.

Show your working.

initial concentration of $\text{NO}$ = .............................. $\text{mol dm}^{-3}$

In experiment 2, the initial concentrations of $\text{NO}$ and $\text{Cl}_2$ are both ten times greater than the initial concentrations used in experiment 1.

Calculate the initial rate of the reaction in experiment 2.

initial rate of reaction in experiment 2 = .............................. $\text{mol dm}^{-3}\text{ s}^{-1}$

The graph of $[\text{Cl}_2]$ against time shows that the reaction has a constant half-life, $t_{1/2}$.

Explain this observation.

Under the conditions used in experiment 3, the value of the rate constant is 105.6.

Show that $t_{1/2}$ of $[\text{Cl}_2]$ is $6.56 \times 10^{-3}\text{ s}$ under these conditions.

Calculate the time taken, in s, for $[\text{Cl}_2]$ to fall to $1.25 \times 10^{-5}\text{ mol dm}^{-3}$ in experiment 3.

time = .............................. s

Sulfur dioxide, $\text{SO}_2$, reacts very slowly with oxygen in the atmosphere, forming sulfur trioxide, $\text{SO}_3$. This reaction is much faster in the presence of $\text{NO}$.

Explain the role of $\text{NO}$ in this process.

Include chemical equations in your answer.

Solid manganese(IV) oxide, $\text{MnO}_2$, catalyses the decomposition of hydrogen peroxide.

State the type of catalysis for this reaction. Explain your answer.

Use the information in Table 3.1 to deduce the rate equation for this reaction.

Explain your reasoning.

Use your rate equation from (b)(i) and the data from Experiment 1 to calculate the rate constant, $k$, for this reaction. Include the units of $k$.

Use Fig. 3.1 to calculate two half-lives, $t_{1/2}$, to show that the reaction is first order.

Use your answer to (c)(i) to calculate the rate constant, $k$, for the decomposition of azomethane.

Describe the effect of increasing temperature on the rate constant and on the rate of a reaction.

Solid manganese(IV) oxide, $\text{MnO}_2$, catalyses the decomposition of hydrogen peroxide.

State the type of catalysis for this reaction. Explain your answer.

Use the information in Table 3.1 to deduce the rate equation for this reaction.

Explain your reasoning.

Use your rate equation from (b)(i) and the data from Experiment 1 to calculate the rate constant, $k$, for this reaction. Include the units of $k$.

Use Fig. 3.1 to calculate two half-lives, $t_{\frac{1}{2}}$, to show that the reaction is first order.

Use your answer to (c)(i) to calculate the rate constant, $k$, for the decomposition of azomethane.

Describe the effect of increasing temperature on the rate constant and on the rate of a reaction.

The rate equation for reaction 3 is shown.

Complete Table 4.1.

Table 4.1

Predict how the initial rate for reaction 3 changes when the concentration of $\text{NO}$ is halved.

Predict how the initial rate for reaction 3 changes when the concentrations of $\text{NO}$ and $\text{H}_2$ are both increased three times.

Suggest why reaction 3 is unlikely to proceed by a mechanism involving only a single step.

Suggest equations for the three steps of the reaction mechanism for reaction 3.

Each step involves a reaction between two molecules.

step 1 .................................................. $\rightarrow$ ..................................................

step 2 ...................... $+$ ...................... $\rightarrow \text{N}_2\text{O} +$ .....................................

step 3 $\text{N}_2\text{O} +$ ..................................... $\rightarrow$ ....................... $+$ .......................

Suggest the role of $\text{N}_2\text{O}$ in this mechanism.

Explain your reasoning.

Iodine, $\text{I}_2$, reacts with thiosulfate ions, $\text{S}_2\text{O}_3^{2-}$, as shown in reaction 4.

Reaction 4 is carried out in the presence of a large excess of $\text{I}_2$. Under these conditions, the reaction is first order with respect to $[\text{S}_2\text{O}_3^{2-}]$ and zero order with respect to $[\text{I}_2]$.

The half-life, $t_{\frac{1}{2}}$, for reaction 4 is $720\text{ s}$ under certain conditions.

Calculate the value of the rate constant, $k$, for reaction 4. Include the units of $k$.

State the type of catalysis that occurs in this reaction.

Explain your reasoning.

Write two equations to show how $\text{Co}^{3+}(\text{aq})$ catalyses this reaction.

equation 1 .........................................................................................................................

equation 2 .........................................................................................................................

Suggest why this reaction is slow in the absence of $\text{Co}^{3+}(\text{aq})$.

Use Fig. 5.1 to determine the initial rate of reaction 1. Show your working.

The rate equation for reaction 1 is $\text{rate} = k [\text{S}_2\text{O}_8^{2-}] [\text{I}^-]$.

Suggest why a large excess of iodide ions allows the rate constant to be determined from the half-life in this investigation.

The reaction of persulfate ions, $\text{S}_2\text{O}_8^{2-}$, with iodide ions is catalysed by $\text{Fe}^{2+}$ ions.

Write two equations to show how $\text{Fe}^{2+}$ catalyses reaction 1.

equation 1 .................................................................................................................................

equation 2 .................................................................................................................................

Describe the effect of an increase in temperature on the rate constant and the rate of reaction 1.

In aqueous solution, thiosulfate ions, $\text{S}_2\text{O}_3^{2-}$, react with hydrogen ions, as shown in reaction 2.

The rate of reaction is first order with respect to $[\text{S}_2\text{O}_3^{2-}]$ and zero order with respect to $[\text{H}^+]$ under certain conditions.

The rate constant, $k$, for this reaction is $1.58 \times 10^{-2}\text{ s}^{-1}$.

Calculate the half-life, $t_{\frac{1}{2}}$, for reaction 2.

The compound nitrosyl bromide, $\text{NOBr}$, can be formed as shown in reaction 3.

The rate is first order with respect to $[\text{NO}]$ and first order with respect to $[\text{Br}_2]$.

The reaction mechanism has two steps.

Suggest equations for the two steps of this mechanism. State which is the rate-determining step.

step 1 ........................................................................................................................

step 2 ........................................................................................................................

rate-determining step = ...............................

Use the data in Table 5.1 to deduce the rate equation for this reaction.

Explain your reasoning.

Use your rate equation from (a)(i) and the data from experiment 1 to calculate the rate constant, $k$, for this reaction. Include the units of $k$.

$\text{NO}_2\text{Cl}$ is another compound containing nitrogen, oxygen and chlorine.

In sunlight, $\text{NO}_2\text{Cl}$ can undergo homolytic fission to release chlorine radicals which can catalyse the conversion of ozone, $\text{O}_3$, into oxygen.

Complete the mechanism for this process.

Ozone reacts with nitrogen dioxide, as shown.

The rate of reaction is first order with respect to $\text{O}_3$ and first order with respect to $\text{NO}_2$.

Suggest equations for a two-step mechanism for this reaction.

step 1 ........................................................................................................................................

step 2 ........................................................................................................................................

Use Fig. 5.1 to determine the initial rate of reaction 1. Show your working.

The rate equation for reaction 1 is $\text{rate} = k [\text{S}_2\text{O}_8^{2-}] [\text{I}^-]$.

Suggest why a large excess of iodide ions allows the rate constant to be determined from the half-life in this investigation.

The reaction of persulfate ions, $\text{S}_2\text{O}_8^{2-}$, with iodide ions is catalysed by $\text{Fe}^{2+}$ ions.

Write two equations to show how $\text{Fe}^{2+}$ catalyses reaction 1.

equation 1 .................................................................................................................................

equation 2 .................................................................................................................................

Describe the effect of an increase in temperature on the rate constant and the rate of reaction 1.

In aqueous solution, thiosulfate ions, $\text{S}_2\text{O}_3^{2-}$, react with hydrogen ions, as shown in reaction 2.

The rate of reaction is first order with respect to $[\text{S}_2\text{O}_3^{2-}]$ and zero order with respect to $[\text{H}^+]$ under certain conditions.

The rate constant, $k$, for this reaction is $1.58 \times 10^{-2}\text{ s}^{-1}$.

Calculate the half-life, $t_{\frac{1}{2}}$, for reaction 2.

The compound nitrosyl bromide, $\text{NOBr}$, can be formed as shown in reaction 3.

The rate is first order with respect to $[\text{NO}]$ and first order with respect to $[\text{Br}_2]$.

The reaction mechanism has two steps.

Suggest equations for the two steps of this mechanism. State which is the rate-determining step.

step 1 ........................................................................................................................................

step 2 ........................................................................................................................................

rate-determining step = ...............................

Explain what is meant by order of reaction.

Complete Table 3.2.

Table 3.2

Use your answer to (a)(ii) to sketch lines in Fig. 3.1 to show the relationship between the initial rates and the concentrations of $[\text{IO}_3^{-}]$ and $[\text{I}^{-}]$.

Use data from Table 3.1 to calculate the rate constant, $k$, for this reaction.

Include the units of $k$.

Use data from Table 3.1 to calculate the concentration of hydrogen ions, $[\text{H}^{+}]$, in experiment 2.

This reaction is repeated in two separate experiments.

The experiments are carried out at the same temperature and with the same concentrations of $\text{I}^{-}$ and $\text{IO}_3^{-}$.

One experiment takes place at pH 1.0 and the other experiment takes place at pH 2.0.

Calculate the value of $\frac{\text{rate at pH 1.0}}{\text{rate at pH 2.0}}$.

In aqueous solution, iron(III) ions react with iodide ions, as shown.

The initial rate of reaction is first order with respect to $\text{Fe}^{3+}$ and second order with respect to $\text{I}^{-}$.

The mechanism for this reaction has three steps.

Each step involves only two ions reacting together.

Suggest equations for the three steps of this mechanism. Identify the rate-determining step.

step 1:

step 2:

step 3:

rate-determining step =