R2.2 How fast? The rate of chemical change

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Practice exam-style IB Chemistry questions for How fast? The rate of chemical change, aligned with the syllabus and grouped by topic.

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Question 1

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SL • Paper 1A

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A clock reaction is repeated using different reactant concentrations. In each trial, the time is measured for the same fixed amount of product to form. What quantity is most directly proportional to the initial rate?

A.

$t$

B.

$t^2$

C.

$\ln t$

D.

$1/t$

Question 2

SL • Paper 1A

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SL • Paper 1A

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The temperature of a reacting gas mixture is increased from $300\ \text{K}$ to $330\ \text{K}$ . What is the main reason for the increase in reaction rate according to collision theory?

A.

A larger fraction of particles has energy at least equal to $E_a$ .

B.

The activation energy of the uncatalysed reaction decreases.

C.

The enthalpy change of the reaction becomes more negative.

D.

All collisions occur with the correct collision geometry.

Question 3

SL • Paper 1A

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SL • Paper 1A

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For the reaction $2\text{NO(g)} + \text{O}_2\text{(g)} \to 2\text{NO}_2\text{(g)}$ , the concentration of $\text{NO}$ decreases from $0.240\ \text{mol dm}^{-3}$ to $0.160\ \text{mol dm}^{-3}$ in $50.0\ \text{s}$ . What is the average rate of reaction over this time interval?

A.

$8.00\times 10^{-4}\ \text{mol dm}^{-3}\ \text{s}^{-1}$

B.

$4.00\times 10^{-4}\ \text{mol dm}^{-3}\ \text{s}^{-1}$

C.

$3.20\times 10^{-3}\ \text{mol dm}^{-3}\ \text{s}^{-1}$

D.

$1.60\times 10^{-3}\ \text{mol dm}^{-3}\ \text{s}^{-1}$

Question 4

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SL • Paper 1A

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Hydrogen gas is produced when magnesium reacts with excess dilute hydrochloric acid. The gas is to be collected and its volume measured continuously. What apparatus is most suitable?

A.

Gas syringe connected to the reaction flask

B.

Thermometer placed above the reaction flask

C.

pH probe placed in the reaction mixture

D.

Colorimeter and cuvette containing the reaction mixture

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Question 5

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SL • Paper 1A

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In repeated trials measuring the volume of gas produced in a reaction, all final gas volumes are lower than the expected stoichiometric volume by a similar amount. What is the most likely type of error?

A.

Random error from variation in reaction temperature

B.

Systematic error from gas leaking from the apparatus

C.

Random error from using a smaller measuring cylinder

D.

Systematic error from repeating the experiment several times

Question 6

HL • Paper 1A

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HL • Paper 1A

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The elementary step $\text{NO}_2\text{(g)} + \text{CO(g)} \to \text{NO(g)} + \text{CO}_2\text{(g)}$ occurs in a proposed mechanism. What is the molecularity of this elementary step?

A.

Unimolecular

B.

Zeromolecular

C.

Termolecular

D.

Bimolecular

Question 7

SL • Paper 2

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SL • Paper 2

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A student investigates the rate of reaction between zinc granules and hydrochloric acid by measuring the hydrogen gas produced.

State one suitable piece of apparatus for measuring the volume of hydrogen continuously.

[1]

State two variables, other than acid concentration, that should be controlled when investigating the effect of acid concentration on the rate.

[2]

Question 8

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SL • Paper 1A

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An exothermic reaction is carried out with and without a catalyst. The catalyst changes only the activation energy. The correct energy profile is shown by which diagram?

A.

B.

C.

D.

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Question 9

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HL • Paper 1A

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Initial-rate data for a reaction are shown.

Experiment	$[A]/\text{mol dm}^{-3}$	$[B]/\text{mol dm}^{-3}$	Initial rate/ $\text{mol dm}^{-3}\ \text{s}^{-1}$
1	$0.100$	$0.100$	$2.00\times 10^{-4}$
2	$0.200$	$0.100$	$8.00\times 10^{-4}$
3	$0.100$	$0.300$	$6.00\times 10^{-4}$

What is the rate equation?

A.

$r=k[A]^2[B]^2$

B.

$r=k[A]^2[B]$

C.

$r=k[A][B]^2$

D.

$r=k[A][B]$

Question 10

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HL • Paper 1A

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A proposed mechanism is shown.

Step 1: $A + B \to I$ slow

Step 2: $I + C \to D$ fast

What are the intermediate and the rate equation predicted by the rate-determining step?

A.

Intermediate $D$ ; $r=k[D]$

B.

Intermediate $I$ ; $r=k[A][B]$

C.

Intermediate $C$ ; $r=k[I][C]$

D.

Intermediate $B$ ; $r=k[A][B][C]$

Question 11

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HL • Paper 1A

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For a reaction with rate equation $r=k[A][B]^2$ , rate is measured in $\text{mol dm}^{-3}\ \text{s}^{-1}$ and concentration in $\text{mol dm}^{-3}$ . What are the units of $k$ ?

A.

$\text{s}^{-1}$

B.

$\text{dm}^3\ \text{mol}^{-1}\ \text{s}^{-1}$

C.

$\text{mol dm}^{-3}\ \text{s}^{-1}$

D.

$\text{dm}^6\ \text{mol}^{-2}\ \text{s}^{-1}$

Question 12

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HL • Paper 1A

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For an Arrhenius plot of $\ln k$ against $1/T$ , what are the gradient and vertical intercept?

A.

Gradient $E_a/R$ ; intercept $A$

B.

Gradient $\ln A$ ; intercept $-E_a/R$

C.

Gradient $-E_a/R$ ; intercept $\ln A$

D.

Gradient $-R/E_a$ ; intercept $\ln A$

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Question 13

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SL • Paper 2

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The concentration of bromine was monitored during its reaction with methanoic acid at constant temperature.

Time / s	$[\text{Br}_2]$ / $\text{mol dm}^{-3}$
0	$0.0100$
20	$0.00640$
40	$0.00440$

Determine the average rate of disappearance of bromine between $20$ s and $40$ s.

[2]

State how the instantaneous rate at $20$ s would be obtained from a concentration-time graph.

[1]

Question 14

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SL • Paper 2

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The rate of reaction between iodide ions and peroxodisulfate ions increases when the temperature is raised from $298\ \text{K}$ to $318\ \text{K}$ .

Define activation energy.

[1]

Explain why increasing the temperature increases the rate of reaction, using collision theory.

[2]

Question 15

SL • Paper 2

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SL • Paper 2

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Magnesium reacts with excess dilute hydrochloric acid. In one experiment a single strip of magnesium is used. In a second experiment the same mass of magnesium is cut into many small pieces.

Predict the effect on the initial rate in the second experiment.

[1]

Explain your prediction in terms of collisions.

[2]

State why changing the pressure would have little direct effect on this reaction mixture.

[1]

Question 16

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HL • Paper 2

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The energy profile for a multistep reaction is shown.

State the number of elementary steps and the number of intermediates shown.

[2]

Identify the rate-determining step and justify your answer.

[2]

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Question 17

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SL • Paper 1B

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Magnesium ribbon was reacted with excess dilute hydrochloric acid. The volume of hydrogen gas was collected in a gas syringe and recorded over time.

Determine the total volume of hydrogen gas produced.

[1]

Calculate the average rate of production of hydrogen gas between $20\ \text{s}$ and $60\ \text{s}$ .

[2]

Explain why the rate decreases during the reaction.

[2]

Question 18

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SL • Paper 1B

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A clock reaction was carried out between sodium thiosulfate solution and hydrochloric acid. The time taken for a fixed amount of sulfur precipitate to obscure a mark was recorded at different sodium thiosulfate concentrations. Temperature and total volume were kept constant.

$\text{Na}_2\text{S}_2\text{O}_3$ concentration / $\text{mol dm}^{-3}$	Time / s
0.0250	160
0.0500	80
0.100	40

State why $1/t$ can be used as a measure of rate in this experiment.

[1]

Use the data to compare the rate when the sodium thiosulfate concentration is doubled from $0.0500\ \text{mol dm}^{-3}$ to $0.100\ \text{mol dm}^{-3}$ .

[2]

Explain the effect of increasing the sodium thiosulfate concentration using collision theory.

[2]

Question 19

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SL • Paper 1B

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Equal masses of calcium carbonate were reacted separately with excess hydrochloric acid. In experiment 1, large chips were used. In experiment 2, powdered calcium carbonate was used. The volume of carbon dioxide was recorded over time.

Compare the initial rates and final volumes of carbon dioxide for the two experiments.

[2]

Explain the difference in rate between the two experiments.

[2]

Question 20

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HL • Paper 1A

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A two-step reaction is exothermic. The second step is rate-determining and there is one intermediate. The correct energy profile is shown by which diagram?

A.

B.

C.

D.

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Question 21

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SL • Paper 2

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A Maxwell-Boltzmann distribution can be used to show the effect of temperature on the proportion of particles able to react.

Sketch Maxwell-Boltzmann energy distribution curves for the same sample at two temperatures, $T_1$ and $T_2$ , where $T_2 > T_1$ . Label the activation energy, $E_a$ , and show the relative number of particles with energy greater than $E_a$ at each temperature.

[3]

Question 22

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SL • Paper 2

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Hydrogen peroxide decomposes exothermically. Manganese(IV) oxide acts as a catalyst for the decomposition.

Sketch an energy profile for the reaction with and without the catalyst.

[3]

State the effect of the catalyst on the enthalpy change of the reaction.

[1]

Question 23

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HL • Paper 2

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The following mechanism is proposed for the reaction $2\text{NO}(g) + \text{Cl}_2(g) \to 2\text{NOCl}(g)$ .

Step 1: $\text{NO}(g) + \text{Cl}_2(g) \to \text{NOCl}_2(g)$ slow

Step 2: $\text{NOCl}_2(g) + \text{NO}(g) \to 2\text{NOCl}(g)$ fast

Identify the intermediate in the mechanism.

[1]

State the molecularity of the slow step.

[1]

The experimentally determined rate equation is $r = k[\text{NO}][\text{Cl}_2]$ . Evaluate whether the mechanism is consistent with the kinetic and stoichiometric data.

[2]

Question 24

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HL • Paper 2

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Initial-rate data were obtained for the reaction $\text{A} + \text{B} \to \text{products}$ at constant temperature.

Experiment	$[\text{A}]$ / $\text{mol dm}^{-3}$	$[\text{B}]$ / $\text{mol dm}^{-3}$	Initial rate / $\text{mol dm}^{-3}\ \text{s}^{-1}$
1	$0.100$	$0.100$	$2.00 \times 10^{-5}$
2	$0.200$	$0.100$	$8.00 \times 10^{-5}$
3	$0.100$	$0.200$	$2.00 \times 10^{-5}$

Deduce the order of reaction with respect to A and with respect to B.

[2]

Write the rate equation and state the overall order.

[2]

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Question 25

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HL • Paper 2

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For a reaction at $298\ \text{K}$ , the rate equation is $r = k[\text{P}][\text{Q}]^2$ . When $[\text{P}] = 0.120\ \text{mol dm}^{-3}$ and $[\text{Q}] = 0.200\ \text{mol dm}^{-3}$ , the initial rate is $5.76 \times 10^{-5}\ \text{mol dm}^{-3}\ \text{s}^{-1}$ .

Calculate the value of the rate constant, $k$ , including its units.

[3]

Question 26

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HL • Paper 2

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The decomposition of X was followed at constant temperature. Three plots were tested using the same concentration-time data. Only the plot of $\ln[\text{X}]$ against time was linear, with gradient $-1.80 \times 10^{-2}\ \text{s}^{-1}$ .

Time / s	$[\text{X}]$ / mol dm $^{-3}$
0	0.100
20	0.0698
40	0.0487
60	0.0340
80	0.0237
100	0.0165
120	0.0115

Determine the order of reaction with respect to X.

[1]

Determine the value and units of the rate constant, $k$ .

[2]

Question 27

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SL • Paper 1B

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The diagram shows Maxwell-Boltzmann energy distributions for the same reacting gas mixture at two different temperatures. The activation energy, $E_a$ , for the reaction is also shown.

State what the area under each distribution curve represents.

[1]

Describe two differences between the distribution at higher temperature and the distribution at lower temperature.

[2]

Explain why the rate is greater at the higher temperature.

[2]

Question 28

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SL • Paper 1B

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An energy profile is shown for an exothermic reaction, with and without a catalyst.

Identify which pathway has the lower activation energy.

[1]

State the effect of the catalyst on the overall enthalpy change of the reaction.

[1]

Explain how the catalyst increases the rate of reaction.

[2]

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Question 29

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HL • Paper 1B

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A proposed mechanism for a reaction is shown. The experimentally determined rate equation is also given.

Item	Details
Step 1 (slow)	$\text{NO}_2 + \text{Cl}_2 \to \text{NO}_2\text{Cl} + \text{Cl}$
Step 2 (fast)	$\text{Cl} + \text{NO}_2 \to \text{NO}_2\text{Cl}$
Experimental rate law	$r = k[\text{NO}_2][\text{Cl}_2]$

Identify the intermediate in the mechanism.

[1]

Deduce the overall equation by adding the elementary steps.

[1]

Evaluate whether the mechanism is consistent with the experimental rate equation.

[2]

Question 30

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HL • Paper 1B

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The concentration of reactant A was monitored during a reaction at constant temperature. Three possible linearized plots were produced from the same concentration-time data.

Time / s	$[A]$ / mol dm $^{-3}$	$\ln([A])$	$1/[A]$ / dm $^{3}$ mol $^{-1}$
0	1.000	0.000	1.000
25	0.3679	-1.000	2.719
50	0.1353	-2.000	7.390
75	0.04979	-3.000	20.09
100	0.01832	-4.000	54.60
125	0.006738	-5.000	148.4

Identify the order of reaction with respect to A.

[1]

Determine the rate constant, including its units.

[2]

State how the value of $k$ would be affected by increasing the temperature, assuming no other change is made.

[1]

Question 31

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HL • Paper 2

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For a first-order decomposition, an Arrhenius plot of $\ln k$ against $1/T$ is linear. The gradient of the best-fit line is $-9.20 \times 10^3\ \text{K}$ and the intercept is $27.5$ . Use $R = 8.31\ \text{J mol}^{-1}\ \text{K}^{-1}$ .

Calculate the activation energy, $E_a$ , in $\text{kJ mol}^{-1}$ .

[2]

Determine the Arrhenius factor, $A$ , including units.

[2]

Question 32

SL • Paper 1B

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SL • Paper 1B

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A student investigated the effect of hydrochloric acid concentration on the initial rate of reaction with magnesium ribbon. The initial rate was determined from the gradient of a tangent to the volume-time curve for each trial. Several repeats were carried out at each acid concentration.

Use the graph to determine whether the initial rate is directly proportional to the hydrochloric acid concentration over the range investigated.

[2]

Comment on the reliability of the data, referring to random error and the anomalous point.

[2]

State one variable, other than acid concentration, that should be controlled in this investigation.

[1]

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Question 33

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HL • Paper 1B

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Initial-rate data were obtained for the reaction between reactants A and B at constant temperature.

Experiment	[A] / mol dm $^{-3}$	[B] / mol dm $^{-3}$	Initial rate / mol dm $^{-3}$ s $^{-1}$
1	0.100	0.100	$2.00 \times 10^{-4}$
2	0.200	0.100	$4.00 \times 10^{-4}$
3	0.100	0.200	$8.00 \times 10^{-4}$
4	0.200	0.200	$1.60 \times 10^{-3}$

Deduce the order of reaction with respect to A and with respect to B.

[2]

Write the rate equation for the reaction.

[1]

Calculate the rate constant, including its units, using the first experiment.

[2]

Question 34

HL • Paper 1B

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HL • Paper 1B

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An energy profile is shown for a reaction that occurs in three elementary steps.

State the number of transition states and intermediates shown.

[2]

Identify the rate-determining step and justify your answer.

[1]

Calculate the overall enthalpy change for the reaction.

[1]

Suggest why changing the concentration of a reactant involved only in step 1 may have little effect on the overall rate under these conditions.

[1]

Question 35

SL • Paper 2

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SL • Paper 2

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A student investigates the reaction between marble chips, $CaCO_3(s)$ , and excess hydrochloric acid.

$CaCO_3(s) + 2HCl(aq) \to CaCl_2(aq) + H_2O(l) + CO_2(g)$

The volume of carbon dioxide is recorded using a gas syringe.

The tangent at $30\ \text{s}$ passes through $(10\ \text{s}, 14.0\ \text{cm}^3)$ and $(50\ \text{s}, 42.0\ \text{cm}^3)$ .

I.

Calculate the instantaneous rate of formation of carbon dioxide at $30\ \text{s}$ in $\text{cm}^3\ \text{s}^{-1}$ .

[2]

II.

State why the rate calculated from a tangent is an instantaneous rate rather than an average rate.

[1]

The experiment is repeated using the same mass of powdered calcium carbonate instead of marble chips. Explain the expected change in the initial rate and the final volume of carbon dioxide.

[2]

Suggest one improvement to the apparatus if the graph showed a lower final gas volume than expected. Explain how this improvement would affect the data.

[2]

Question 36

SL • Paper 2

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SL • Paper 2

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The reaction between iodide ions and peroxodisulfate ions is followed using a clock method. A fixed amount of thiosulfate ion and starch indicator is added. The time taken for a blue-black colour to appear is measured for different initial concentrations.

$S_2O_8^{2-}(aq) + 2I^-(aq) \to 2SO_4^{2-}(aq) + I_2(aq)$

The clock endpoint corresponds to the formation of the same small amount of iodine in every trial.

I.

Explain why $\frac{1}{t}$ can be used as a measure proportional to the initial rate in this experiment.

[2]

II.

Explain, using collision theory, why increasing the concentration of iodide ions is expected to decrease the time to the endpoint.

[2]

Evaluate two variables, other than the concentration being investigated, that must be controlled to make the comparison of rates valid.

[2]

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Question 37

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SL • Paper 2

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A student investigates the reaction between sodium thiosulfate solution and hydrochloric acid by timing how long it takes for a cross under the flask to disappear.

$S_2O_3^{2-}(aq) + 2H^+(aq) \to SO_2(aq) + S(s) + H_2O(l)$

The time for the cross to disappear in one trial is $48.0\ \text{s}$ .

I.

Calculate the relative rate, using $\frac{1}{t}$ , in $\text{s}^{-1}$ .

[1]

II.

Explain why the disappearance of the cross is a fixed endpoint for comparing rates.

[1]

Design a fair comparison to investigate the effect of thiosulfate concentration on the rate.

[3]

Use the graph to distinguish between random error and systematic error in this investigation.

[2]

Question 38

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HL • Paper 1B

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Rate constants were measured for a first-order reaction at several temperatures. The data were used to construct an Arrhenius plot of $\ln k$ against $1/T$ .

Determine the activation energy, $E_a$ , from the gradient of the Arrhenius plot.

[2]

Determine the Arrhenius factor, $A$ , including units.

[2]

Explain why temperature must be expressed in kelvin in the Arrhenius equation.

[1]

Question 39

HL • Paper 1B

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HL • Paper 1B

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Two mechanisms have been proposed for the reaction $2\text{NO}+\text{Br}_2\to 2\text{NOBr}$ . The experimentally determined rate equation is $r=k[\text{NO}]^2[\text{Br}_2]$ .

Mechanism	Step 1	Step 2
A	$\text{NO} + \text{Br}_2 \to \text{NOBr}_2$ slow	$\text{NOBr}_2 + \text{NO} \to 2\text{NOBr}$ fast
B	$\text{NO} + \text{Br}_2 \rightleftharpoons \text{NOBr}_2$ fast	$\text{NOBr}_2 + \text{NO} \to 2\text{NOBr}$ slow

State the molecularity of the slow step in mechanism B.

[1]

Show that mechanism B gives the correct overall equation.

[1]

Evaluate which mechanism is more consistent with the experimental rate equation.

[2]

State why a mechanism that matches the rate equation is still described as a possible mechanism.

[1]

Question 40

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SL • Paper 2

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The rate of a gas-phase reaction increases when the temperature is raised from $T_1$ to $T_2$ , where $T_2>T_1$ .

Use the axes provided.

I.

Sketch Maxwell-Boltzmann energy distribution curves for $T_1$ and $T_2$ .

[2]

II.

Explain why the increase in temperature causes a large increase in rate.

[2]

A catalyst is added at constant temperature. Explain how the Maxwell-Boltzmann diagram can be used to show the effect of the catalyst on the rate.

[3]

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Question 41

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SL • Paper 2

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An uncatalysed exothermic reaction is represented by an energy profile. The same reaction can also occur by a catalysed pathway.

Refer to the energy profile shown.

I.

Identify the activation energy and the transition state on the diagram.

[2]

II.

State how the diagram shows that the reaction is exothermic.

[1]

Sketch the catalysed pathway on the same diagram and explain the effect of the catalyst on the rate.

[3]

Discuss the statement: "A catalyst increases the yield of product because it makes the products more stable."

[2]

Question 42

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SL • Paper 2

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Hydrogen peroxide decomposes slowly at room temperature but much faster in the presence of catalase, an enzyme found in potato tissue.

$2H_2O_2(aq) \to 2H_2O(l) + O_2(g)$

The volume of oxygen collected is used to determine the reaction rate.

I.

Explain why using a gas syringe is suitable for following the rate continuously.

[2]

II.

At one instant oxygen is forming at $1.20\times10^{-4}\ \text{mol dm}^{-3}\ \text{s}^{-1}$ . Determine the rate of disappearance of hydrogen peroxide at this instant.

[2]

Explain the roles of sufficient energy and proper orientation in the decomposition of hydrogen peroxide.

[2]

Compare the effect of catalase with the effect of finely divided manganese(IV) oxide, $MnO_2$ , on this reaction.

[2]

Question 43

HL • Paper 2

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HL • Paper 2

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The initial rate of a reaction involving aqueous reactants $A$ , $B$ and $C$ is measured at constant temperature. The balanced equation is not sufficient to determine the rate equation.

Experiment	[A] / mol dm $^{-3}$	[B] / mol dm $^{-3}$	[C] / mol dm $^{-3}$	Initial rate / mol dm $^{-3}$ s $^{-1}$
1	0.100	0.100	0.100	$2.00 \times 10^{-4}$
2	0.200	0.100	0.100	$8.00 \times 10^{-4}$
3	0.100	0.200	0.100	$4.00 \times 10^{-4}$
4	0.100	0.100	0.300	$2.00 \times 10^{-4}$

Use the initial-rate data.

I.

Deduce the order of reaction with respect to $A$ , $B$ and $C$ .

[3]

II.

Write the rate equation and state the overall order.

[2]

In one experiment, $[A]=0.100\ \text{mol dm}^{-3}$ , $[B]=0.100\ \text{mol dm}^{-3}$ and the initial rate is $2.00\times10^{-4}\ \text{mol dm}^{-3}\ \text{s}^{-1}$ . Calculate $k$ and give its units.

[2]

Explain why the orders in the rate equation cannot be deduced from the coefficients in the overall balanced equation.

[1]

Question 44

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HL • Paper 2

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A possible mechanism for a reaction is shown.

$\begin{aligned} \text{Step 1: }& A \to I &&\text{fast}\\ \text{Step 2: }& I+B \to P &&\text{slow}\\ \text{Step 3: }& P+B \to C+D &&\text{fast} \end{aligned}$

Consider the proposed sequence of elementary steps.

I.

Show that the steps give the overall equation and identify the intermediates.

[3]

II.

State the molecularity of step 2 and explain why termolecular elementary steps are uncommon.

[2]

Use the energy profile.

I.

Identify the rate-determining step and the transition state for this step.

[2]

II.

Explain why the expression $r=k[I][B]$ is not normally written as the experimental rate equation for the overall reaction.

[1]

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Question 45

HL • Paper 2

Hard

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HL • Paper 2

Hard

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The decomposition of a compound $X$ is monitored at constant temperature. The data are processed using three possible linear plots.

The gradient of the straight-line plot is $-4.62\times10^{-2}\ \text{s}^{-1}$ .

I.

Deduce the order of reaction with respect to $X$ .

[1]

II.

Determine the rate constant, including units.

[2]

III.

Calculate the half-life of $X$ .

[2]

Calculate the instantaneous rate when $[X]=0.0800\ \text{mol dm}^{-3}$ .

[1]

Explain why identifying a first-order graph does not prove a unique reaction mechanism.

[1]

Question 46

HL • Paper 2

Hard

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HL • Paper 2

Hard

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The hydrolysis of a tertiary halogenoalkane, $RBr$ , in aqueous alkali is investigated.

$RBr(aq)+OH^-(aq) \to ROH(aq)+Br^-(aq)$

Initial-rate data show that doubling $[RBr]$ doubles the rate, but doubling $[OH^-]$ has no effect on the rate.

Use the kinetic information.

I.

Deduce the rate equation.

[2]

II.

State the units of $k$ .

[1]

Two possible mechanisms are proposed.

Mechanism 1:

$\begin{aligned} \text{Step 1: }&RBr \to R^+ + Br^- &&\text{slow}\\ \text{Step 2: }&R^+ + OH^- \to ROH &&\text{fast} \end{aligned}$

Mechanism 2:

$RBr+OH^- \to ROH+Br^- \quad \text{slow}$

[3]

Evaluate whether the kinetic data prove mechanism 1.

[2]

Question 47

HL • Paper 2

Hard

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HL • Paper 2

Hard

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The rate constant for a first-order reaction is measured at different temperatures. An Arrhenius plot of $\ln k$ against $1/T$ is obtained.

The gradient of the line is $-7.20\times10^3\ \text{K}$ and the intercept is $18.4$ .

I.

Calculate the activation energy, $E_a$ , in $\text{kJ mol}^{-1}$ .

[2]

II.

Determine the Arrhenius factor, $A$ , including units.

[2]

Calculate $k$ at $310\ \text{K}$ using the line equation.

[2]

Explain the physical significance of the temperature scale and the Arrhenius factor in this analysis.

[2]

Question 48

HL • Paper 2

Hard

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HL • Paper 2

Hard

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A clock reaction is studied at two temperatures using identical initial concentrations and the same fixed visible endpoint. The reciprocal of the time to the endpoint is used as a value proportional to $k$ .

At $298\ \text{K}$ , $k=1.25\times10^{-2}\ \text{s}^{-1}$ . At $318\ \text{K}$ , $k=1.00\times10^{-1}\ \text{s}^{-1}$ .

Analyse the temperature-dependence data.

I.

Explain why $\frac{1}{t}$ can be used as a value proportional to $k$ only under the stated experimental conditions.

[2]

II.

Calculate the activation energy using the two values of $k$ .

[3]

Use the value $E_a=81.9\ \text{kJ mol}^{-1}$ to estimate the Arrhenius factor, $A$ , from the data at $298\ \text{K}$ .

[2]

Evaluate one major experimental limitation when using a clock reaction to determine Arrhenius parameters.

[1]

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Test your knowledge with a custom paper containing questions from selected topics.

R2.1 How much? The amount of chemical change

R2.3 How far? The extent of chemical change

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