CHM Ch6 - Rates of chemical reactions

Rate of reaction

The change in concentration of reagents or products per unit time (M/s)

Ways of measuring rate of reaction

• Measure how much reactant has been used

• Measure how much product has been formed

• Do this through measuring change in mass, gas volume, colour and pH

Collision theory

For a reaction to occur: Molecules must collide in the correct orientation with enough energy to overcome the activation energy barrier

Activation energy

The amount of energy required to break the bonds of the reactants

Transition state

Occurs when the activation energy is absorbed, causing an unstable, temporary arrangement of atoms where bonds are both breaking and forming

Factors affecting the rate of reaction (5)

• Concentration of reactants

• Gas pressure

• Surface area (solids)

• Presence of a catalyst

• Temperature

Effect of concentration of reactants on the rate of reaction

Increasing the concentration of reactants increases the frequency of successful collisions, increasing the rate of reaction

Effect of gas pressure on the rate of reaction

Increasing gas pressure by adding more gas to a fixed-volume or decreasing the volume of a container increases the concentration of gas molecules, leading to a higher frequency of successful collisions, increasing the rate of reaction

Effect of surface area on the rate of reaction

Smaller particles have a larger surface area than large particles, allowing more reactants to be at the surface, increasing the frequency of successful collisions for a higher rate of reaction

Effect of temperature on the rate of reaction

Increasing temperature increases the rate of reaction in two ways:

• It causes particles to move faster, increasing the frequency of successful collisions

• It increases the average kinetic energy of particles, meaning a greater proportion of particles have sufficient energy to overcome the activation energy barrier

Kinetic energy (definition and equation)

The energy of motion

KE = (1/2)mv²

Catalysts

Catalysts are substances that increase the rate of a reaction by creating an alternative energy pathway with a lower activation energy. This means a greater proportion of particles have sufficient energy to overcome the activation energy barrier, causing a higher frequency of successful collisions, increasing the rate of reaction.

• Catalysts are not consumed in the reaction

• A catalyst does not change the delta H of a reaction

Homogenous catalysts

Are in th same physical state as the reactants and products

Heterogenous catalysts

Are in a different physical state from the reactants and products

Catalysts in industry are usually heterogenous because they are:

• More easily separated from the products

• Easier to reuse

• Usable at high temperatures

Solid catalysts

Some solid catalysts work by using a high surface area to adsorb (bond with) gas molecule. This distorts or breaks bonds in the gas molecules, lowering the activation energy to increase the rate of reaction