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Revision:Collision Theory

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TSR Wiki > Study Help > Subjects and Revision > Revision Notes > Chemistry > GCSE Chemistry Revision Notes > Collision Theory

Collision Theory

In a chemical reaction the reacting particles need to collide. The collision must also have enough energy so that the chemical bonds can be broken. Collision without enough energy will not lead to a reaction. An effective reaction is a reaction which does have enough energy and does lead to a reaction.

Chemical reactions occur at different speeds. Some reactions are faster while others are much slower. For example an explosive chemical reaction between two reactants tells us that this is a very fast reaction whereas rusting is a much slower reaction.

Particles need enough kinetic energy to break the bonds and cause a chemical reaction to occur. The minimum amount of kinetic energy that a reaction requires to occur is known as the activation energy. So when the particles collide there must be enough kinetic energy to exceed the activation energy in order for a chemical reaction to occur. Slow reactions such as rusting generally have high activation energies, while explosive chemical reactions generally have low activation energies.

Controlling the Rate of Reaction

The rate of the reaction can be controlled by changing the frequency of the collisions or changing the energy of the reactant particle.

Factor Effect on Rate Explanation
Concentration Increasing the concentration of a reactant increases the rate Increasing the concentration means that the particles are more crowded, so there will be more frequent effective collisions increasing the rate of reaction.
Temperature Increasing the temperature of a reaction increases the rate Increasing the temperature increases the kinetic energy that the reactants have, which means that they will collide more often and have more energy available to exceed the activation energy. At lower temperatures the particles have less kinetic energy, are moving much slower and so there are less effective collisions.
Surface area Increasing the surface area of the solid or cutting the solid into smaller pieces increases the rate of reaction Increasing the surface area increases the number of particles that are exposed to the reactants which increases the number of collisions and so increases the number of effective reactions. So a powder form of a reactant will have a higher rate of reaction than the same reactant in a block form.
Catalyst Adding a catalyst to the reaction changes the rate of reaction A catalyst changes the activation energy in a reaction. It basically speeds up a reaction without being used up itself. It can raise the activation energy in some reactions but it usually lowers it. This means that more particles will have enough kinetic energy to overcome the activation energy level.
Pressure Increasing the pressure of reactions involving gasses increases the rate of reaction Once the pressure has been increased the gas particles are in a smaller space meaning that more frequent and effective reactions will occur.

Changing the rate of reaction does not change the volume or mass of product formed, instead it only changes how quickly the product is formed.

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