Reaction Kinetics (CIE AS Chemistry)

Fig. 1.1 below shows, for a given temperature T, a Boltzmann distribution of the kinetic energy of the molecules of a mixture of two gases that will react together.

The activation energy for the reaction, E_a, is marked.

Fig. 1.1

On Fig. 1.1 above,

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Explain the meaning of the term activation energy.

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Two reactions involving aqueous NaOH are given below.

The reagents in reaction 1 must be heated together for some time for the reaction to occur.
Whereas, reaction 2 is almost instantaneous at room temperature.

Suggest brief explanations why the rates of these two reactions are very different.

A platinum-rhodium catalyst used in the reaction forming nitrogen(II) oxide from ammonia.

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

A (s) + B (aq) → C (g)

The students collected the gas produced and plotted the graph shown in Fig. 1.1. 

Fig. 1.1

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.

This question is about how different factors affect the rate of a chemical reaction.

A reaction between reactant molecules can only occur when collisions between the reactant molecules are effective.

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A student made the following statement about factors that affect the rate of a chemical reaction.

Discuss why this statement is only true to an extent and what other factors need to be considered to fully explain how the rate of a reaction is affected by temperature.

Hydrogen and oxygen can react to form water in an exothermic reaction.

The rate of this reaction depends on the pressure as well as the temperature.

Describe and explain the effect of increasing the pressure on this reaction.

If the rate of reaction increases suddenly, an explosion can occur.

A highly exothermic reaction, such as the reaction in (c), is at more risk of exploding than a less exothermic reaction.

Suggest why.

A series of experiments was carried out to investigate the factors which affect the rate of reaction between calcium carbonate and dilute hydrochloric acid.

CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)

• 50.0 cm³ of hydrochloric acid was added to 10 g of small pieces of calcium carbonate (an excess) in a conical flask placed on an electronic balance.
• The loss in mass of the flask and its contents was recorded every 30 seconds for 10 minutes.
• The experiment was repeated using a different concentration of hydrochloric acid or a different temperature.
• The results are shown in Table 1.1

Table 1.1

The results of Experiment 1 are shown in Fig. 1.1.

Fig. 1.1

Draw curves on the graph in Fig 1.1 to show the results you would expect for Experiments 2 and 3.

Label the curves 2 and 3.

Determine the initial rate of reaction for Experiment 1.

You must show your working on the graph in Fig 1.2.
Include units in your answer.

Calculate the average rate of reaction over the first 240 seconds of Experiment 1.

The student repeated each experiment but used 10.0 g of larger sized pieces of calcium carbonate.

Suggest what effect this would have on the rate of reaction. Give your reasoning.

Ammonia can be produced by the reaction of nitrogen and hydrogen.

N₂ (g) + 3H₂ (g)  2NH₃ (g)          ΔH = -92 kJ mol^-1

The reaction can be catalysed and the activation energy for this catalysed reaction is +109 kJ mol^-1

Complete the reaction pathway diagram in Fig. 1.1 for the uncatalysed and the catalysed reaction between nitrogen and hydrogen.

You should label the following:

• products
• the enthalpy change of reaction, ΔH
• the activation energy of the forward, uncatalysed reaction, E_a
• the activation energy of the forward, catalysed reaction, E_c

Fig. 1.1

Calculate the value of the activation energy of the catalysed decomposition of ammonia into nitrogen and hydrogen.

Show your working.

activation energy .............................................. kJ mol^-1

Catalysts, such as iron used in the production of ammonia, increase the rate of reaction.

Explain why. Use a labelled Boltzmann distribution to explain your answer.

Platinum is used as a catalyst in catalytic converters which are fitted to vehicle exhaust systems to remove nitrogen oxide from the exhaust gases.

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Reaction rates can be affected by a range of factors including changes in pressure and temperature.

Fig 1.1

On Fig 1.1:

Hydrogen iodide can be used in the manufacturing of pharmaceuticals and can be broken down back into its elements in standard form, iodine and hydrogen. 

The activation energy when uncatalysed is +183 kJ mol^-1 and when catalysed with gold it is +105 kJ mol^-1. 

Catalysts are often used in industrial processes and can be used in a variety of forms. 

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The Contact process is an important industrial process, contributing to the production of sulfuric acid. In the Contact process, solid vanadium (V) oxide, a heterogeneous catalyst, is used to make sulfur trioxide from sulfur dioxide and oxygen. This process is reversible.

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.

A Maxwell-Boltzmann distribution curve is shown below in Fig 3.1 

Fig 3.1

For the changes detailed in part (i) and (ii) state and explain the effect 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 greater than or equal to E_a 

A chemist performed a reaction at three different temperatures, 100 K, 300 K and 700 K as shown by the Maxwell-Boltzmann distribution graph in Fig 3.2.

Fig 3.2

Hydrogen will react with chlorine to form the hydrogen halide, hydrogen chloride, a colourless gas.

H₂ (g) + Cl₂ (g) → 2HCl (g)

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