Rate equations and orders of reaction

This question is about rates of reaction.

The rate of hydrolysis of a bromoalkane (RBr) is investigated using this method.

• A boiling tube has a cross drawn onto it.

• An aqueous solution of silver nitrate is added to an ethanolic solution of RBr in the boiling tube in a warm water bath.

• The time is recorded for the cross on the boiling tube to become invisible.

The method is repeated at different temperatures.

State why the cross on the boiling tube becomes invisible.

The volume and concentration of each substance used at each temperature is kept constant.

Identify one other variable that should be kept constant to ensure that the investigation is fair.

The reaction between the bromoalkane RBr and an aqueous solution of KOH is investigated.

The rate constant k for this reaction is determined at a series of temperatures T.

The Arrhenius equation can be written in this form.

Table 2 shows an incomplete set of results for this investigation.

Complete Table 2.

The results in Table 2 can be used to plot a graph to determine the activation energy E_a for this reaction.

Complete the x-axis in Figure 10 by adding an appropriate scale.

Plot the graph on the grid in Figure 10.

Use your graph to determine a value, in kJ mol^–1, for the activation energy E_a

The gas constant, R = 8.31 J K^–1 mol^–1

The rate equation for the reaction between a different bromoalkane (R₂CHBr) and an aqueous solution of KOH is

Which of these is a correct mechanism for the reaction between R₂CHBr and KOH?

Tick (✓) one box.

Explain, in terms of bond polarity, why R₂CHBr and KOH react together.

The effect of reactant concentrations on the rate of a reaction is investigated.

The equation for the reaction is shown.

The experiment is done five times in the presence of thiosulfate ions and starch. The reactant concentrations are varied by changing the volume of each reactant in the mixture. The thiosulfate ions react with the iodine as it is produced, until all the thiosulfate has reacted.

After all the thiosulfate has reacted, the iodine produced in the mixture then reacts with the starch and the solution turns blue-black. The time taken for the solution to turn blue-black is recorded. If the amount of thiosulfate is the same in each experiment, the time recorded represents the time to produce the same amount of iodine in each experiment.

Table 3 shows the results of the investigation.

The amount of thiosulfate is the same in each experiment. To allow the order of reaction with respect to each reactant to be calculated, several other variables need to be controlled.

Identify three other control variables.

Control variable 1

Control variable 2

Control variable 3

The total uncertainty in the times recorded for each experiment is ±0.5 s but the stopwatch used has a measurement uncertainty of ±0.01 s

Suggest why, in each experiment, the uncertainty in the times recorded is greater than the measurement uncertainty of the stopwatch.

Calculate the percentage uncertainty in the time recorded in experiment 3.

The five experiments are repeated with the same stopwatch and the same

Suggest and explain how the method can be changed to decrease the percentage uncertainty in the times recorded.

Propanone reacts with bromine in alkaline conditions.

The rate equation for this reaction is

Sketch a graph on the axes provided to show how, at constant temperature, the concentration of bromine changes during this reaction.

Table 1 shows the initial rate of this reaction for experiments using different mixtures containing propanone, bromine and hydroxide ions.

Complete Table 1.

Use the data from experiment 1 to calculate the rate constant k for this reaction.

Give the units for the rate constant.

Figure 1 shows an incomplete mechanism for this reaction.

Complete the mechanism in Figure 1 by adding four curly arrows and any relevant lone pair(s) of electrons.

Use evidence from the rate equation to explain why Step 1 is the rate determining step.

The thermal decomposition of but-3-en-1-ol is investigated at different temperatures (T).

CH₂=CHCH₂CH₂OH → CH₂=CHCH₃ + HCHO

The results from the investigation are used to calculate the rate constant, k, at each temperature.

Table 2 shows some of the results.

Complete Table 2 with the missing values at 573 K

The overall order of the reaction can be deduced from a piece of information in one of the column headings in Table 2.

Identify this piece of information and deduce the overall order.

Piece of information

The Arrhenius equation can be written in the form shown.

On the grid in Figure 3 plot a graph of ln k against

Use your graph from Question 03.3 to calculate a value for E_a, in kJ mol^–1, for the thermal decomposition of but-3-en-1-ol.

The gas constant, R = 8.31 J K^–1 mol^–1

2-Methylpent-4-en-2-ol decomposes in a similar way to but-3-en-1-ol, to produce an alkene and a carbonyl compound.

Deduce the structures of the alkene and the carbonyl compound.

alkene
carbonyl compound

A and B react together in the presence of an acid catalyst.

A(aq) + 2 B(aq) → C(aq) + D(aq)

The rate equation for this reaction is

Table 1 shows how the values of the relative initial rate vary with different concentrations of each reagent at the same temperature.

Complete Table 1 by calculating the missing values.

A suggested mechanism for the reaction is shown.

Deduce the rate-determining step for this reaction.

Give a reason for your answer.

Rate-determining step

This question is about ethanoic anhydride.

In the gas phase, ethanoic anhydride (CH₃CO)₂O decomposes to form ethenone.

The equation is

Ethenone is the simplest member of the ketene homologous series. Ketenes all contain one C=C double bond and one C=O double bond.

Deduce the general formula for the ketene homologous series.

Figure 11 shows an incomplete suggested mechanism for the decomposition of ethanoic anhydride.

Complete the mechanism in Figure 11 by adding three curly arrows and any relevant lone pairs of electrons.

For a chemical reaction the relationship between the rate constant, k, and the temperature, T, is shown by the Arrhenius equation.

For the decomposition of gaseous ethanoic anhydride

the activation energy, E_a = 34.5 kJ mol^–1

At temperature T₁ the rate constant, k = 2.48 10⁸ s^–1

The gas constant, R = 8.31 J K^–1 mol^–1

Sketch the Maxwell–Boltzmann distribution of molecular energies for gaseous ethanoic anhydride at temperature T₁ and at a higher temperature T₂

Include a label for each axis, and mark on the appropriate axis a typical position for the activation energy.

Explain why the rate of reaction is faster at T₂

Explanation

An acidified solution of butanone reacts with iodine as shown.

Draw the displayed formula for CH₃CH₂COCH₂I

Give the name of CH₃CH₂COCH₂I

Displayed formula

The rate equation for the reaction is

Table 1 shows the initial concentrations used in an experiment.

The initial rate of reaction in this experiment is 1.45×10⁻⁴ mol dm⁻³ s⁻¹

Calculate the value of the rate constant, k, for the reaction and give its units.

Calculate the initial rate of reaction when all of the initial concentrations are halved.

An experiment was done to measure the time, t, taken for a solution of iodine to react completely when added to an excess of an acidified solution of butanone.

Suggest an observation used to judge when all the iodine had reacted.

The experiment was repeated at different temperatures. Figure 1 shows how varied with temperature for these experiments.

Describe and explain the shape of the graph in Figure 1.

Deduce the time taken for the reaction at 35^oC

For a different reaction, Table 2 shows the value of the rate constant at different temperatures.

This equation can be used to calculate the activation energy, E_a

Calculate the value, in kJ mol⁻¹, of the activation energy, E_a

The gas constant, R = 8.31 J K⁻¹mol⁻¹

Name and outline the mechanism for the reaction of butanone with KCN followed by dilute acid.

Name of mechanism

Outline of mechanism

This question is about rates of reaction. Iodine and propanone react together in an acid-catalysed reaction

A student completed a series of experiments to determine the order of reaction with respect to iodine.

Method

• Transfer 25 cm³ of 1.0 mol dm^–3 propanone solution into a conical flask.

• Add 10 cm³ of 1.0 mol dm^–3 HCl(aq)

• Add 25 cm³ of 5.0 10^–3 mol dm^–3 I₂(aq) and start a timer.

• At intervals of 1 minute, remove a 1.0 cm³ sample of the mixture and add each sample to a separate beaker containing an excess of NaHCO₃(aq)

• Titrate the contents of each beaker with a standard solution of sodium thiosulfate and record the volume of sodium thiosulfate used.

Suggest why the 1.0 cm³ portions of the reaction mixture are added to an excess of

Suggest why the order of this reaction with respect to propanone can be ignored in this experiment.

The volume of sodium thiosulfate solution used in each titration is proportional to the concentration of iodine in each beaker.

Table 5 shows the results of the experiment.

Use the results in Table 5 to draw a graph of volume of sodium thiosulfate solution against time.

Draw a line of best fit.

Explain how the graph shows that the reaction is zero-order with respect to iodine in the reaction between propanone and iodine.

The Arrhenius equation can be written as

Figure 8 shows a graph of ln k against for the reaction

Use Figure 8 to calculate a value for the activation energy (E_a), in kJ mol⁻¹, for this reaction.

Use the graph in Figure 1 to determine the initial rate of reaction for this experiment. State its units. Show your working on the graph.

Another student reacted different initial concentrations of phosphinate ions with an excess of hydroxide ions. The student measured the time (t) taken to collect 15 cm³ of hydrogen gas. Each experiment was carried out at the same temperature. Table 1 shows the results.

State the relationship between the initial concentration of phosphinate and time (t).

Deduce the order of the reaction with respect to phosphinate.

Complete the diagram in Figure 2 to show how the hydrogen gas could be collected and measured in the experiments in Questions 01.1 and 01.2.

Figure 2

The rate equation for a different reaction is

Deduce the overall effect on the rate of reaction when the concentrations of both L and M are halved.

Calculate the concentration of M if the rate constant is 21.3 mol^–2 dm⁶ s^–1

Define the term overall order of reaction.

Cisplatin, [Pt(NH₃)₂Cl₂], is used as an anti-cancer drug.

Cisplatin works by causing the death of rapidly dividing cells.

Name the process that is prevented by cisplatin during cell division.

After cisplatin enters a cell, one of the chloride ligands is replaced by a water molecule to form a complex ion, B.

Give the equation for this reaction.

When the complex ion B reacts with DNA, the water molecule is replaced as a bond forms between platinum and a nitrogen atom in a guanine nucleotide. The remaining chloride ligand is also replaced as a bond forms between platinum and a nitrogen atom in another guanine nucleotide.

Figure 1 represents two adjacent guanine nucleotides in DNA.

Complete Figure 1 to show how the platinum complex forms a cross-link between the guanine nucleotides.

An experiment is done to investigate the rate of reaction in Question 04.2.

During the experiment the concentration of cisplatin is measured at one-minute intervals.

Explain how graphical methods can be used to process the measured results, to confirm that the reaction is first order.

In another experiment, the effect of temperature on the rate of the reaction in Question 04.2 is investigated.

Table 1 shows the results.

Complete Table 1.

The Arrhenius equation can be written in the form

Use the data in Table 1 to plot a graph of ln k against on the grid in Figure 2. Calculate the activation energy, E_a, in kJ mol¹

The gas constant, R = 8.31 J K¹ mol¹

The rate expression for the reaction between X and Y is

Which statement is correct?

What are the units of the rate constant for a third order reaction?

Substances P and Q react in solution at a constant temperature. The initial rate of reaction was studied in three experiments by measuring the change in concentration of P over the first five seconds of the reaction. The data obtained are shown in Table 1.

Complete Table 2 to show the initial rate of reaction of P in each experiment.

Determine the order of reaction with respect to P and the order of reaction with respect to Q.

A reaction between substances R and S was second order with respect to R and second order with respect to S. At a given temperature, the initial rate of reaction was 1.20 × 10^–3 mol dm^–3 s^–1 when the initial concentration of R was 1.00 × 10^–2 mol dm^–3 and the initial concentration of S was 2.45 × 10^–2 mol dm^–3

Calculate a value for the rate constant, k, for the reaction at this temperature. Give the units for k

The rate constant, k, for a reaction varies with temperature as shown by the equation

For this reaction, at 25 °C, k = 3.46 × 10⁸ s¹ The activation energy E_a = 96.2 kJ mol¹ The gas constant R = 8.31 J K¹ mol¹

Calculate a value for the Arrhenius constant, A, for this reaction. Give the units for A.

The results of an investigation of the reaction between P and Q are shown in this

The rate equation is: rate = k [P] [Q]²

What is the initial concentration of Q in experiment 2?

Iodide ions are oxidised to iodine by hydrogen peroxide in acidic conditions.

The rate equation for this reaction can be written as

In an experiment to determine the order with respect to H⁺(aq), a reaction mixture is made containing H⁺(aq) with a concentration of 0.500 mol dm^–3

A large excess of both H₂O₂ and I^– is used in this reaction mixture so that the rate equation can be simplified to

Explain why the use of a large excess of H₂O₂ and I^– means that the rate of reaction at a fixed temperature depends only on the concentration of H⁺(aq).

Samples of the reaction mixture are removed at timed intervals and titrated with alkali

State and explain what must be done to each sample before it is titrated with alkali.

A graph of the results is shown in Figure 1.

Explain how the graph shows that the order with respect to H⁺(aq) is zero.

Use the graph in Figure 1 to calculate the value of k₁

A second reaction mixture is made at the same temperature. The initial concentrations of H⁺(aq) and I^–(aq) in this mixture are both 0.500 mol dm^–3 There is a large excess of H₂O₂

In this reaction mixture, the rate depends only on the concentration of I^–(aq).

The results are shown in Table 1.

Plot these results on the grid in Figure 2. The first three points have been plotted.

Draw a line of best fit on the grid in Figure 2.

Calculate the rate of reaction when [H⁺] = 0.35 mol dm^–3 Show your working using a suitable construction on the graph in Figure 2.

A general equation for a reaction is shown.

In aqueous solution, A, B, C and D are all colourless but E is dark blue.

A reagent (X) is available that reacts rapidly with E. This means that, if a small amount of X is included in the initial reaction mixture, it will react with any E produced until all of the X has been used up.

Explain, giving brief experimental details, how you could use a series of experiments to determine the order of this reaction with respect to A. In each experiment you should obtain a measure of the initial rate of reaction.

Solutions of two compounds, W and X, react together in the presence of a soluble catalyst, Y, as shown in the equation

When the concentrations of W, X and Y are all doubled, the rate of reaction increases by a factor of four.

Which is a possible rate equation for this reaction?

The rate equation for the reaction between compounds A and B is

Figure 2 shows how, in an experiment, the concentration of A changes with time, t, in this reaction.

0 2. Draw a tangent to the curve at t = 0

Use this tangent to deduce the initial rate of the reaction.

Initial rate
mol dm⁻³s⁻¹

The experiment was repeated at the same temperature and with the same initial concentration of B but with a different initial concentration of A. The new initial rate was 1.7 times greater than in the original experiment.

Calculate the new initial concentration of A.

rate = k [ C ][ D ] In one experiment at 25 °C, the initial rate of reaction is 3.1 10⁻³ mol dm⁻³ s⁻¹ when the initial concentration of C is 0.48 mol dm⁻³ and the initial concentration

Calculate a value for the rate constant at this temperature and give its units.

An equation that relates the rate constant, k, to the activation energy, E_a, and the temperature, T, is

Use this equation and your answer from Question 3.1 to calculate a value, in kJ mol⁻¹, for the activation energy of this reaction at 25 °C. For this reaction ln A = 16.9 The gas constant R = 8.31 J K⁻¹ mol⁻¹ (If you were unable to complete Question 3.1 you should use the value of 3.2 10⁻³ for the rate constant. This is not the correct value.)