7.5 Chapter summary

7.5 Chapter summary (ESCNC)

Presentation: 27V5

• The average rate of a reaction describes how quickly reactants are used, or how quickly products form. The units used are usually moles per second.

• A number of factors can affect the average rate of a reaction. These include the nature of the reactants, the concentration of solutions, surface area of solids or pressure of gases, the temperature of the reaction and the presence or absence of a catalyst.

• Collision theory provides one way of explaining why each of the factors mentioned in the previous bullet can affect the average rate of a reaction. For example, higher temperatures mean increased reaction rates because the reactant particles have more energy and are more likely to collide successfully with each other.

• Different methods can be used to measure the average rate of a reaction. The method used will depend on the nature of the product or products. Reactions that produce gases can be measured by collecting the gas in a syringe or by measuring a change in mass of the reaction vessel. In reactions that produce a precipitate the mass of the dried precipitate can be measured.

• A change in colour can show that the product has formed or that the reaction has reached completion. (This is known as colorimetry).

• For any reaction to occur, a minimum amount of energy is needed so that bonds in the reactants can break, and new bonds can form in the products. The minimum energy that is required is called the activation energy of a reaction.

• In reactions where the particles do not have enough energy to overcome this activation energy, one of two methods can be used to facilitate a reaction to take place: increase the temperature of the reaction or add a catalyst.

• Increasing the temperature of a reaction means that the average energy of the reactant particles increases and they are more likely to have enough energy to overcome the activation energy and the chances of two molecules having the correct orientation are increased.

• A catalyst is used to lower the activation energy so that the reaction is more likely to take place. A catalyst does this by providing an alternative, lower energy pathway for the reaction.

• A catalyst therefore speeds up a reaction but remains unchanged after the reaction is complete.