Learning Objectives

The Requirements for All Reactions (Collision Theory)

Activation Energy

The initial required energy for the reaction to occur (minimum amount of energy). (affected by temperature and the presence of a catalyst)

Orientation

Reactants must be in the correct orientation for a reaction to occur.

(affected by temperature and pressure)

Collisions

For a reaction to occur particles must collide. (w/ sufficient energy and the correct orientation)

(affected by concentration, pressure and surface area)

Kinetic Energy = Ek= 1/2 mV²

Reaction Rate

Reaction Rate is the frequency of collisions (or the rate of particle collisions)

The proportion of successful collisions is the fraction of collisions that have sufficient energy and the appropriate orientation.

Reaction Rate = Collision Frequency X Proportion of Successful Collisions

1/time is used for the inversely proportional relationship (e.g. pressure vs reaction rate) to show an increasing effect.

Factors Effecting Reaction Rate

Nature of reactants

Re-arranging of bonds decreases the rate at which a reaction occurs. Ions can easily form products without having to rearrange bonds, thus are capable of reacting quickly. If bonds need to be broken, they need to have enough energy and suitable orientation, slowing the reaction.

Concentration/Pressure

Higher concentration or pressure increases the likelihood of particles colliding, hence increasing the frequency of successful collisions. Number of products formed increases per unit time, resulting in the reaction's fast rate of progression.

Surface Area (the state of subdivision)

Higher surface area increases the area for collisions to occur, increasing the frequency of collisions, and resulting successful collisions. Number of products formed increases per unit time, resulting in the reaction's fast rate of progression.

Temperature

Increasing temperature increases the average kinetic energy of particles, this increases the proportion of particles with sufficient kinetic energy to overcome the activation energy barrier, increasing the frequency and percentage of successful collisions.

Catalysts

Catalysts are chemical agents which increase reaction rate but are not consumed during the reaction. The presence of a catalyst loosens the bonds between reactant molecules, providing an alternative pathway for the reaction to take place, with a lower activation energy. The lower activation energy means that a higher proportion of molecules will have kinetic exceeding the activation energy barrier. This results in more frequent collisions, increasing reaction rate.

The increase of reaction rate due to the presence of a catalyst is known as "catalysis".

Homogenous Catalysts are in the same state as the reactants and products

Heterogeneous Catalysts are in a different state to the reactants and products (this is more favourable for industry, as it makes the catalyst easier to separate, and can be used for adsorption)