Determining 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 rate of reaction between calcium carbonate and hydrochloric acid is investigated using a continuous monitoring method.

Method

• Place a conical flask on a balance and add approximately 20 g of large marble chips.

• Add 50 cm³ of 0.4 mol dm^–3 hydrochloric acid.

• Place a loose cotton wool plug in the neck of the flask.

• Zero the mass reading on the balance.

• Start a timer.

• Record the loss in mass (m_t) every 30 seconds for 4 minutes.

• Wait for the reaction to finish and record the total mass loss (m_total).

• Plot a graph of (m_total – m_t) against time.

Figure 1 shows a graph of the results obtained during the first 240 s

Suggest why a loose cotton wool plug is placed in the neck of the flask, instead of leaving the flask open or inserting a bung.

Instead of leaving the flask open

Instead of inserting a bung

The mass of carbon dioxide produced in time t is equal to m_t.

The total mass of CO₂ produced during the reaction is equal to m_total.

Explain why (m_total – m_t) is proportional to the concentration of hydrochloric acid remaining in the flask at time t.

Table 1 shows the rate of reaction, calculated from the gradient of the curve, at five different times.

(m_total – m_t) is proportional to the concentration of unreacted HCl at time t.

On the grid in Figure 2 plot the rate of reaction (y-axis) against (m_total – m_t) (x-axis).

State how the graph in Figure 2 confirms that the rate equation for this reaction is

Rate = k [HC l ]

In this experiment the variable measured is mass loss.

The rate of this reaction at a constant temperature can be investigated in other ways.

Suggest two other variables that can be measured instead of mass loss.

A student investigated the reaction between magnesium and hydrochloric acid.

The results are plotted on this graph.

Which value is closest to the rate of reaction, in cm³ s^–1, at 70 s?

This question is about rates of reaction.

Figure 1 shows apparatus used to measure the rate of reaction when an acid reacts with an excess of solid sodium hydrogencarbonate, NaHCO₃

When different monoprotic organic acids are used, the rates at which gas escapes can be used to compare the strengths of the acids.

A timer is started when the NaHCO₃ is added to the acid and the mass of CO₂ gas lost is recorded at regular intervals. (It is assumed that any change in mass is due to the loss of CO₂)

Suggest a reason why using a conical flask instead of a beaker would give more accurate results in this experiment.

Figure 2 shows the results of this experiment when 25.0 cm³ of a 2.23 mol dm^–3 solution of ethanoic acid reacts with an excess of NaHCO₃

Use Figure 2 to calculate the rate of reaction at 2 minutes.

Deduce the units of your calculated rate.

Chloroethanoic acid is a stronger acid than ethanoic acid.

Sketch, on Figure 2, the curve you would expect when 25.0 cm³ of a 2.23 mol dm^–3 solution of chloroethanoic acid reacts with an excess of NaHCO₃

Suggest why chloroethanoic acid is a stronger acid than ethanoic acid.

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

Hydrogen peroxide solution decomposes to form water and oxygen.

The reaction is catalysed by manganese(IV) oxide.

A student determines the order of this reaction with respect to hydrogen peroxide. The student uses a continuous monitoring method in the experiment.

The student places hydrogen peroxide solution in a conical flask with the catalyst and uses a gas syringe to collect the oxygen formed. The student records the volume of oxygen every 10 seconds for 100 seconds.

Explain why the reaction is fastest at the start.

The graph in Figure 1 shows how the concentration of hydrogen peroxide changes with time in this experiment.

Tangents to the curve in Figure 1 can be used to determine rates of reaction.

Draw a tangent to the curve when the concentration of hydrogen peroxide solution is

Use your tangent to calculate the gradient of the curve at this point.

The concentration of hydrogen peroxide solution at time t during the experiment can be calculated using this expression.

H₂O_2t = concentration of hydrogen peroxide solution at time t / mol dm_–3

H₂O_2initial = concentration of hydrogen peroxide solution at the start / mol dm_–3

V_max = total volume of oxygen gas collected during the whole experiment / cm³

V_t = volume of oxygen gas collected at time t / cm³

Use Figure 1 and the expression to calculate H₂O_2t when 20 cm₃ of oxygen has been collected.

Table 5 shows data from a similar experiment.

Plot the data from Table 5 on the grid in Figure 2.

Draw a line of best fit.

Use Figure 2 to determine the order of reaction with respect to H₂O₂

State how the graph shows this order.

How the graph shows this order

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

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

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.

A series of experiments was carried out to find the order of reaction with respect to reactant X. In these experiments, only the concentration of X was changed.

Which graph would show that the reaction is second-order with respect to X ?

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.