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# Factors affecting reaction rates

Several factors affect the rate of a reaction. It is important to know these factors so that reaction rates can be controlled. This is particularly important when it comes to industrial reactions, so that productivity can be maximised. The following are some of the factors that affect the rate of a reaction.

1. Nature of reactants

Substances have different chemical properties and therefore react differently and at different rates.

2. Concentration (or pressure in the case of gases)

As the concentration of the reactants increases, so does the reaction rate.

3. Temperature

If the temperature of the reaction increases, so does the rate of the reaction.

4. Catalyst

Adding a catalyst increases the reaction rate.

5. Surface area of solid reactants

Increasing the surface area of the reactants (e.g. if a solid reactant is broken or ground up into smaller pieces) will increase the reaction rate.

## General experiment 1: The nature of reactants.

Aim

To determine the effect of the nature of reactants on the rate of a reaction.

Apparatus

Oxalic acid $((COOH)2)$, iron(II) sulphate $(FeSO4)$, potassium permanganate $(KMnO4)$, concentrated sulphuric acid $(H2SO4)$, spatula, test tubes, medicine dropper, glass beaker and glass rod.

Method

1. In the first test tube, prepare an iron (II) sulphate solution by dissolving about two spatula points of iron (II) sulphate in 10 cm3 of water.

2. In the second test tube, prepare a solution of oxalic acid in the same way.

3. Prepare a solution of sulphuric acid by adding 1 cm3 of the concentrated acid to about 4 cm3 of water. Remember always to add the acid to the water, and never the other way around.

4. Add 2 cm3 of the sulphuric acid solution to the iron(II) and oxalic acid solutions respectively.

5. Using the medicine dropper, add a few drops of potassium permanganate to the two test tubes. Once you have done this, observe how quickly each solution discolours the potassium permanganate solution.

Results

• You should have seen that the oxalic acid solution discolours the potassium permanganate much more slowly than the iron(II) sulphate.

• It is the oxalate ions $(C2O42-)$ and the $Fe2+$ ions that cause the discolouration. It is clear that the $Fe2+$ions act much more quickly than the $(C2O42-)$ ions. The reason for this is that there are no covalent bonds to be broken in the ions before the reaction can take place. In the case of the oxalate ions, covalent bonds between carbon and oxygen atoms must be broken first.

Conclusions

The nature of the reactants can affect the rate of a reaction.

## Interesting Fact:

Oxalic acids are abundant in many plants. The leaves of the tea plant (Camellia sinensis) contain very high concentrations of oxalic acid relative to other plants. Oxalic acid also occurs in small amounts in foods such as parsley, chocolate, nuts and berries. Oxalic acid irritates the lining of the gut when it is eaten, and can be fatal in very large doses.

## General experiment 2: Surface area and reaction rates.

Marble $(CaCO3)$ reacts with hydrochloric acid $(HCl)$ to form calcium chloride, water and carbon dioxide gas according to the following equation:

$CaCO 3+2 HCl → CaCl 2+H2O+ CO 2$

Aim

To determine the effect of the surface area of reactants on the rate of the reaction.

Apparatus

2 g marble chips, 2 g powdered marble, hydrochloric acid, beaker, two test tubes.

Method

1. Prepare a solution of hydrochloric acid in the beaker by adding 2 cm3 of the concentrated solution to 20 cm3 of water.

2. Place the marble chips and powdered marble into separate test tubes.

3. Add 10 cm3 of the dilute hydrochloric acid to each of the test tubes and observe the rate at which carbon dioxide gas is produced.

Results

• Which reaction proceeds the fastest?

• Can you explain this?

Conclusions

The reaction with powdered marble is the fastest. The smaller the pieces of marble are, the greater the surface area for the reaction to take place. The greater the surface area of the reactants, the faster the reaction rate will be.

## General experiment 3: Reactant concentration and reaction rate.

Aim

To determine the effect of reactant concentration on reaction rate.

Apparatus

Concentrated hydrochloric acid $(HCl)$, magnesium ribbon, two beakers, two test tubes, measuring cylinder.

Method

1. Prepare a solution of dilute hydrochloric acid in one of the beakers by diluting 1 part concentrated acid with 10 parts water. For example, if you measure 1 cm3 of concentrated acid in a measuring cylinder and pour it into a beaker, you will need to add 10 cm3 of water to the beaker as well. In the same way, if you pour 2 cm3 of concentrated acid into a beaker, you will need to add 20 cm3 of water. Both of these are 1:10 solutions. Pour 10 cm3 of the 1:10 solution into a test tube and mark it 'A'. Remember to add the acid to the water, and not the other way around.

2. Prepare a second solution of dilute hydrochloric acid by diluting 1 part concentrated acid with 20 parts water. Pour 10 cm3 of this 1:20 solution into a second test tube and mark it 'B'.

3. Take two pieces of magnesium ribbon of the same length. At the same time, put one piece of magnesium ribbon into test tube A and the other into test tube B, and observe closely what happens.

The equation for the reaction is:

$2HCl+Mg→MgCl2+H2$

Results

• Which of the two solutions is more concentrated, the 1:10 or 1:20 hydrochloric acid solution?

• In which of the test tubes is the reaction the fastest? Suggest a reason for this.

• How can you measure the rate of this reaction?

• What is the gas that is given off?

• Why was it important that the same length of magnesium ribbon was used for each reaction?

Conclusions

The 1:10 solution is more concentrated and this reaction therefore proceeds faster. The greater the concentration of the reactants, the faster the rate of the reaction. The rate of the reaction can be measured by the rate at which hydrogen gas is produced.

## Group discussion 1: The effect of temperature on reaction rate

1. In groups of 4–6, design an experiment that will help you to see the effect of temperature on the reaction time of 2 cm of magnesium ribbon and 20 ml of vinegar. During your group discussion, you should think about the following:

• What equipment will you need?

• How will you conduct the experiment to make sure that you are able to compare the results for different temperatures?

• How will you record your results?

• What safety precautions will you need to take when you carry out this experiment?

2. Present your experiment ideas to the rest of the class, and give them a chance to comment on what you have done.

3. Once you have received feedback, carry out the experiment and record your results.

4. What can you conclude from your experiment?