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# Reaction rates and collision theory

It should be clear now that the rate of a reaction varies depending on a number of factors. But how can we explain why reactions take place at different speeds under different conditions? Why, for example, does an increase in the surface area of the reactants also increase the rate of the reaction? One way to explain this is to use collision theory.

For a reaction to occur, the particles that are reacting must collide with one another. Only a fraction of all the collisions that take place actually cause a chemical change. These are called 'successful' collisions. When there is an increase in the concentration of reactants, the chance that reactant particles will collide with each other also increases because there are more particles in that space. In other words, the collision frequency of the reactants increases. The number of successful collisions will therefore also increase, and so will the rate of the reaction. In the same way, if the surface area of the reactants increases, there is also a greater chance that successful collisions will occur.

Definition 1: Collision theory

Collision theory is a theory that explains how chemical reactions occur and why reaction rates differ for different reactions. The theory states that for a reaction to occur the reactant particles must collide, but that only a certain fraction of the total collisions, the effective collisions, actually cause the reactant molecules to change into products. This is because only a small number of the molecules have enough energy and the right orientation at the moment of impact to break the existing bonds and form new bonds.

When the temperature of the reaction increases, the average kinetic energy of the reactant particles increases and they will move around much more actively. They are therefore more likely to collide with one another (Figure 1). Increasing the temperature also increases the number of particles whose energy will be greater than the activation energy for the reaction (refer Section 3.5).

## Exercise 1: Rates of reaction

Hydrochloric acid and calcium carbonate react according to the following equation:

$CaCO 3+2 HCl → CaCl 2+H2O+ CO 2$

The volume of carbon dioxide that is produced during the reaction is measured at different times. The results are shown in the table below.

 Time (mins) Total Volume of $CO2$ produced (cm3) 1 14 2 26 3 36 4 44 5 50 6 58 7 65 8 70 9 74 10 77

Note: On a graph of production against time, it is the gradient of the graph that shows the rate of the reaction.

Questions:

1. Use the data in the table to draw a graph showing the volume of gas that is produced in the reaction, over a period of 10 minutes.

2. At which of the following times is the reaction fastest? Time = 1 minute ; time = 6 minutes or time = 8 minutes .

3. Suggest a reason why the reaction slows down over time.

4. Use the graph to estimate the volume of gas that will have been produced after 11 minutes.

5. After what time do you think the reaction will stop?

6. If the experiment was repeated using a more concentrated hydrochloric acid solution...

1. would the rate of the reaction increase or decrease from the one shown in the graph?

2. draw a rough line on the graph to show how you would expect the reaction to proceed with a more concentrated HCl solution.

1.

2. Time = 1 minute. This is where the gradient of the graph is the steepest.

3. Carbon dioxide is a gas and so can escape from the reaction vessel. This means that there is less carbon dioxide available to reform the reactants and so the reactants are used up.

4. About $80{\text{cm}}^{3}$

5. Any answer between 15 and 25 minutes is reasonable. To see this extend the line and find approximate the time that the gradient flattens out.

6. a. The rate would increase.

b. The red line indicates roughly how the reaction would proceed. Note that the reaction does not produce more carbon dioxide, it just reacts faster.