How Fast? The Rate of Chemical Change (HL IB Chemistry)

When investigating the reaction between sulfuric acid and calcium carbonate, it was observed that a small increase of temperature of around 10 ^oC caused a doubling in the rate of the reaction.

Sketch and label Maxwell-Boltzmann curve for the two temperatures T and T+10, and use this diagram to help to explain this effect of temperature on rate.

Identify and explain another factor that affects the number of particles present in a solution with sufficient energy to react.

Some groups investigating the effect of temperature on rate stirred their reactions, some did not.

Explain the effect of stirring upon the rate of the reaction.

0.5 g of magnesium reacts with 50 cm³ of 0.01 moldm^-3 nitric acid. Magnesium is in excess.

A graph monitoring the volume of hydrogen gas produced is shown below:

Compare the expected rate and progress of the reaction if 25 cm³ of 0.2 mol dm^-3 nitric acid was used instead of 50 cm³ of 0.1 mol dm^-3 nitric acid.

Suggest one change to the reaction that could be made to produce more hydrogen gas in total and explain your choice.

Suggest why it is often better to study a slower reaction instead of a faster one.

The following energy profile diagram shows the pathways for both a catalysed and uncatalysed reversible reaction:

Identify the letter(s) representing the activation energy for the catalysed reverse reaction.

State and explain the effect that this catalyst will have on the equilibrium yield.

Vehicles with combustion engines usually have catalytic convertors added to catalyse the oxidation of carbon monoxide into carbon dioxide and to catalyse the reduction of nitrogen oxides to nitrogen. These catalysts are usually rhodium or platinum.

Leaded fuels were phased out as they were found to poison these catalysts, binding irreversibly to the metal surface.

Explain the problems for drivers of the catalysts being poisoned.

Suggest a situation in which using a catalyst would not be appropriate.

The conversion of hydrogen and iodine into hydrogen iodide proceeds via a three step reaction mechanism:

1. I₂ (g)   2I (g)                   fast
2. H₂ (g) + I (g)  H₂I (g)      fast
3. H₂I (g) + I (g) → 2HI (g)    slow

Write the rate equation for this reaction and show how the mechanism is consistent with the stoichiometric equation.

An investigation into the rate of reaction between hydrogen and iodine was carried out at 298 K and the data obtained is shown below.

Determine the rate equation for the reaction and justify your answer.

Calculate the rate constant using Expt 2 data, including its units.

Using section 12 of the Data booklet, determine whether the forward reaction is favoured by an increase in temperature.

The reaction between iodide ions and persulfate ions is a 'clock' reaction and often used to study reaction kinetics.

2I^- (aq) + S₂O₈^2- (aq) → I₂ (aq) + 2SO₄^2- (aq)

Deduce the redox changes taking place in the reaction.

A persulfate-iodide clock reaction was studied and the following rate data obtained.

Deduce the order with respect to persulfate ions and iodide ions.

Determine the rate equation for the reaction and calculate rate constant, including the units.

Four mechanisms are proposed for the persulfate-iodide reaction. Deduce which mechanism(s) is/are consistent with the rate equation in part c) and justify your answer.

Mechanism 1:

1. I^- (aq) + I^- (aq) → I₂^2- (aq)                                         slow
2.  I₂^2- (aq) + S₂O₈^2- (aq) → I₂ (aq) + 2SO₄^2- (aq)    fast

Mechanism 2:

1. I^- (aq) + S₂O₈^2- (aq) → S₂O₈I^3- (aq)                     slow
2.  S₂O₈I^3- (aq) + I^- (aq) → I₂ (aq) + 2SO₄^2- (aq)    fast

Mechanism 3:

1. I^- (aq) + S₂O₈^2- (aq)  → S₂O₈I^3- (aq)                    fast
2.  S₂O₈I^3- (aq) + I^- (aq) → I₂ (aq) + 2SO₄^2- (aq)    slow

Mechanism 4:

1.  2I^- (aq) + S₂O₈^2- (aq) → I₂ (aq) + 2SO₄^2- (aq)    slow

The reaction between nitrogen monoxide and hydrogen produces nitrogen and water:

2NO (g) + 2H₂ (g) →  N₂ (g) + 2H₂O (g)

Rate data for this reaction is shown below.

What is the molecularity of the reaction?

Draw a sketch graphs of:

Suggest a possible mechanism for the reaction.

Suggest a Lewis structure for N₂O₂ and draw the shape of the molecule.

The rate of reaction between manganate(VII) ions and oxalate ions, C₂O₄^2-, can be investigated by measuring how the concentration of manganate(VII) varies with time.

2MnO₄^- (aq) + 16H⁺ (aq) + 5C₂O₄^2-  →  2Mn²⁺(aq) + 8H₂O(l) + 10CO₂(g) 

The rate is first order with respect to oxalate ions and the general rate equation for the reaction is:

rate = k [MnO₄^-]^p[C₂O₄^2-]^q[H⁺]^r

[1]

[2]

The student used an acid concentration of 1.0 mol dm^-3. She then varied it, keeping the other concentrations constant. She measured the rate of reaction and found the following results:

Identify the relationship between the relative rate of reaction and H⁺, and hence determine the order of reaction with respect to H⁺ ions.

The student varied the concentration of [MnO₄^-] and plotted the rate against time at three different concentrations:

[2]

A reaction proceeds by a three step mechanism. The energy profile for the reaction is shown below:

Explain the difference between points A, C, E and B, D on the profile.

Deduce which step is the rate determining step of the reaction, giving a reason.

A series of experiments were carried out to investigate how the rate of the reaction of bromate and bromide in acidic conditions varies with temperature.

The time taken, t, was measured for a fixed amount of bromine to form at different temperatures. The results are shown below. 

Complete the table above.

The Arrhenius equation relates the rate constant, k, to the activation energy, E_a, and
temperature, T.

ln k = ln A +

In this experiment, the rate constant, k, is directly proportional to . Therefore, 

ln = ln A +

Use your answers from part (a) to plot a graph of ln against x 10^-3 on the graph below.

Use section 2 of the data booklet along with your graph and information from part (b) to calculate a value for the activation energy, in kJ mol^–1, for this reaction.

To gain full marks you must show all of your working.

Three experiments were carried out at a temperature, T₁,  to investigate the rate of the reaction between compounds F and G. The results are shown in the table below:

Use the data in the table to deduce the rate equation for the reaction between compounds F and G.

Use the information in the table in part (a) to calculate a value for the rate constant, k, for this reaction between 0.0534 mol dm^-3 F and 0.0624 mol dm^-3 G. 

Give your answer to the appropriate number of significant figures.

State the units for k.

(If you did not get an answer for (a), you may assume that rate = k [F]² [G]². This is not the correct answer)

The Arrhenius equation shows how the rate constant, k, for a reaction varies with temperature, T.

k =

For the reaction between 0.0534 mol dm^-3 F and 0.0624 mol dm^-3 G at 25 °C, the activation energy, E_a, is 16.7 kJ mol^–1. 

Use section 2 of the data booklet and your answer to part (b) to calculate a value for the Arrhenius constant, A, for this reaction. 

Give your answer to the appropriate number of significant figures.

(If you did not get an answer for (b), you may assume that k has a value of 4.97. This is not the correct answer)

The temperature of the reaction is increased to twice the original temperature, T₁.
The value of k increases to 0.28 mol^-1 dm³ s^-1 at this new temperature.
 

Using sections 1 and 2 of the data booklet and your answer to part (b), determine the original temperature, T₁.

(If you did not get an answer for (b), you may assume that k = 16700 mol^-1 dm³ s^-1 This is not the correct answer)

The rate constant for a reaction doubles when the temperature is increased from 25.0 °C to 35 °C.

Calculate the activation energy, E_a, in kJ mol⁻¹ for the reaction using section 1 and 2 of the data booklet.

The rate constant is 6.2 x 10³ s^-1 when the temperature is reduced by a factor of a fifth from the original starting temperature, 25 °C.

Calculate the rate constant, in min^-1, using sections 1 and 2 of the data booklet.

A different reaction route is used which reduces the activation energy of the reaction. 

Explain how the rate constant calculated in part(b) would differ.

State the meaning of the term rate of reaction.

A group of students were completing a practical, investigating the factors which affect the rate of the chemical reaction shown below. 

A (s) + B (aq) → C (g)

The students collected the gas produced and plotted the graph shown in Figure 1. 

Figure 1

Explain why the gradient of the curve in part (b) decreases as the time of the reaction progresses.

Another way to increase the rate of reaction is to increase the temperature. 

Explain why a small increase in temperature has a large effect on the initial rate of a chemical reaction.  

The decomposition of hydrogen peroxide into water and oxygen is a very slow chemical reaction. 

Write the equation for the decomposition of hydrogen peroxide.

The rate of decomposition of hydrogen peroxide can be found by collecting and measuring the volume of gas formed at specific time intervals.

The decomposition of hydrogen peroxide is a slow reaction, so a catalyst is often added to speed up the rate of the reaction. Catalysts are used in many chemical reactions to increase the rate.

The following shows a two-step reaction mechanism of a chemical reaction, where a catalyst, X is used. 

STEP 1:                                  W + X  →  Y + Z

STEP 2:                                  Y + W → Z + A + X

OVERALL REACTION:         2W → 2Z + A

Give a reason, other than the rate of reaction increasing, why it can be deduced from the three equations above that X is a catalyst.

The graph shown below represents the decomposition of hydrogen peroxide.

Figure 1

The graph starts to level out as the reaction slows down.

State why the rate of the reaction slows down over time.  

During the following reaction, A and B react together to produce C. 

A  +  2B  ⇌  C

Figure 1 shows the production of C over time.   

Figure 1

In the reaction in part (a), large pieces of A were used.

Use collision theory to explain what would happen to the rate of the reaction if powdered A was used instead of large pieces.

In a different reaction, gaseous W and X were added together to produce Y and Z as shown in the equation below:

2W (g)  +  X (g)  →  Y (g)  + 2Z (g)

A catalyst was added to speed up the rate of reaction. 

Some changes were made individually to the experiment completed in part (c). 

Consider your Maxwell-Boltzmann distribution curve from part (c). For each of the changes in parts (i), (ii) and (iii) below, state and explain the effect that the change would have on:

• The area under the curve 
• The value of the most probable energy of the molecules (E_mp) 
• The proportion of molecules with energy greater than or equal to E_a
  

The number of molecules in the original reaction mixture is increased, but no other changes are made.

[2] 

A catalyst is added to the original reaction mixture, but no other changes are made.

[2]

Iodine reacts with propanone in an acid catalyzed reaction, according to reaction equation below.  

CH₃COCH₃ (aq) + I₂ (aq) → CH₃COCH₂I (aq) + H⁺ (aq) + I^- (aq)

Suggest how the change in concentration of iodine could be used to determine the rate of the above reaction.  

A group of students completed the iodination of propanone reaction using the same acid catalyst, but with different concentrations. The results achieved are shown in the table below:

Table 1

Use the table to state and explain the relationship between the concentration of acid used in the reaction and the rate.

Sodium thiosulfate and hydrochloric acid will react together readily, as shown by the equation below:

Na₂S₂O₃  +  2HCl  →  2NaCl  +  S  +  SO₂  +  H₂O

This reaction is often referred to as the ‘disappearing cross’ experiment. The cross disappears when viewed from above because the solution turns cloudy as a sulfur precipitate is formed, covering the cross.  

Figure 1

The speed of the reaction can be increased, by raising the temperature of the sodium thiosulfate solution in the reaction. The thiosulfate solution is heated to different temperatures before the acid is added, and the time it takes for the cross to disappear is recorded. The times can then be compared.

Suggest one reason why the value for the rate of reaction when a higher temperature was used may be less accurate than at a lower temperature.

Collision theory can be used to explain why different factors affect the rate of a chemical reaction.

Describe collision theory.

A student carried out a metal displacement reaction between zinc powder and copper(II) sulfate solution. The equation for the reaction is

Zn (s) + CuSO₄ (aq) → ZnSO₄ (aq) + Cu (s)

3.78 g of zinc powder was added to 50.0 cm³ of 0.250 moldm^-3 copper(II) sulfate solution.

Determine the limiting reagent showing your working.

The reaction between the zinc and copper sulfate was carried out in a polystyrene cup and the temperature change was measured using a temperature probe. The maximum temperature rise the student recorded was 8.5 ^oC. 

Using sections 1 and 2 of the data booklet, calculate the enthalpy change, ∆H, for the reaction, in kJ.

Assume that all the heat evolved was absorbed by the solution, and that the density and specific heat capacity of the copper(II) sulfate solution are the same as pure water.

State two further assumptions made in the calculation of ∆H.

Using Figure 1, sketch a graph of the concentration of zinc sulfate, ZnSO₄ (aq), versus time and show how the graph may be used to find the initial rate of reaction.

                                                                  Figure 1

A student investigated the rate of decomposition of hydrogen peroxide, H₂O₂, at a temperature of 45 ^o The decomposition reaction occurs in the presence of a catalyst, MnO₂.

The results she obtained are shown in Table 1 below.

Table 1

Plot a graph on the axes below in Figure 1 and from it determine the rate of reaction after 60 s.

Figure 1

On the same graph sketch the shape obtained if the student had carried out the same reaction at 60 ^oC. Explain the shape of the graph at 60 ^oC.

The decomposition of hydrogen peroxide can be investigated by measuring the volume of oxygen given off using the apparatus shown in Figure 2.

Figure 2

Two students decide to measure the rate of decomposition for H₂O₂ using the change in mass as oxygen escapes from the reaction container.

One student says that they should use a three decimal place rather than two decimal place balance because it will make their results more accurate. The second student disagrees and says it will make their results more precise, but not more accurate.

Which student is correct?

For the reaction below, consider the following experimental data.
 

X (aq) + Y (aq) → Z (aq)

Deduce the order of reaction with respect to X.

Deduce the order of the reaction with respect to Y.

Write the rate expression for the reaction between X and Y.

Determine the rate constant, k, correct to three significant figures and state its units, using data from Experiment 2. 

Explain why the reaction represented below is a redox reaction. 

2ClO₂ (aq) + 2NaOH (aq) → NaClO₃ (aq) + NaClO₂ (aq) + H₂O (l) 

For the reaction below, consider the following experimental data. 

2ClO₂ (aq) + 2OH^- (aq) → ClO₃^- (aq) + ClO₂^- (aq) + H₂O (l)

Deduce the rate expression.

Determine the rate constant, k, and state its units, using data from Experiment 3. 

Calculate the rate when [ClO₂ (aq)] = 3.10 x 10^-2 mol dm^-3 and [OH^- (aq)] = 1.50 x 10^-2 mol dm^-3.

Sketch a graph to show how the rate constant, k, varies with temperature.

The following mechanism is proposed for the reaction where ethanal dimerises in dilute alkaline solution to form 3-hydroxybutanal: 

Step 1: CH₃CHO + :OH^- → :CH₂CHO + H₂O 
Step 2: CH₃CHO + :CH₂CHO → CH₃CH(O:^-)CH₂CHO
Step 3: CH₃CH(O:^-)CH₂CHO + H₂O → CH₃CH(OH)CH₂CHO + :OH^- 

Classify OH^-, CH₂CHO and CH₃CH(O:^-)CH₂CHO as reactant, product, catalyst or intermediate, based on the proposed mechanism.

Using the following information about the proposed mechanism, deduce the rate expression. 

Step 1: CH₃CHO + :OH^- → :CH₂CHO + H₂O                                                slow step
Step 2: CH₃CHO + :CH₂CHO → CH₃CH(O:^-)CH₂CHO                                fast step
Step 3: CH₃CH(O:^-)CH₂CHO + H₂O → CH₃CH(OH)CH₂CHO + :OH^-         fast step

Calculate the initial rate of reaction for experiment 2, if measured under the same conditions.

State the effect, if any, increasing the concentration of a reactant would have on the value of the rate constant, k.

Nitrogen dioxide and carbon monoxide react according to the following equation. 

            NO₂ (g) + CO (g) → NO (g) + CO₂ (g) 

Using the following graph, what is the order with respect to NO₂?

 

The rate expression for the reaction of nitrogen dioxide and carbon monoxide at T < 227 ^OC is: 

            Rate = k [NO₂]² 

Sketch a labelled graph of concentration against time for carbon monoxide.

A student proposed the following single step mechanism for the reaction of nitrogen dioxide and carbon monoxide. 

NO₂ + CO → NO + CO₂         slow 

Rate = k [NO₂]² 

Justify whether the student’s proposed mechanism is correct.

Another student proposed the following mechanism for the reaction of nitrogen dioxide and carbon monoxide. 

Step 1:             NO₂ + NO₂ → NO + NO₃
   Step 2:             NO₃ + 2CO → NO + 2CO₂ 

Rate = k [NO₂]² 

Explain which of the student’s proposed mechanism steps is the rate determining step.

Nitrogen(II) oxide is oxidised according to the following equation. 

2NO (g) + O₂ (g) → 2NO₂ (g) 

The following mechanism is proposed for the two-step oxidation of nitrogen(II) oxide. 

Step 1:             NO (g) + NO (g) → N₂O₂ (g)
   Step 2:             N₂O₂ (g) + O₂ (g) → 2NO₂ (g) 

The potential energy profile for this two-step reaction is shown.

Explain which step is the rate determining step.

Energy profile diagrams give evidence for or against a proposed mechanism or proposed rate expression.

Explain why the following reaction between iodide ions and peroxodisulfate ions has a high activation energy. 

S₂O₈^2- (aq) + 2I^- (aq) → 2SO₄^2- (aq) + I₂ (aq)

Sketch the potential energy diagram for the reaction of iodide ions with peroxodisulfate ions catalysed by iron(II) ions according to the following mechanism. 

2Fe²⁺ (aq) + S₂O₈^2- (aq) → 2Fe³⁺ (aq) + 2SO₄^2- (aq)             slow
2Fe³⁺ (aq) + 2I^- (aq) → 2Fe²⁺ (aq) + I₂ (aq)                             fast