Rates of Reaction
The rate of a reaction is defined as the change in concentration per unit time in any
one reactant or product.
The rate of reaction depends on 5 factors.
 Nature of reactants
 Particle size
 Concentration
 Temperature
 Catalysts.
Nature of reactants
The reaction between acidified sodium dichromate and Ammonium Iron (11) sulphate
is instantaneous. (Ionic)
The reaction between acidified sodium dichromate and ethanal occurs much more
slowly. (Covalent)
Particle Size.
Large Marble chips
Small marble chips
Powdered Marble
CaCO3 +2HCl 
CaCl2 +CO2 + H2O
The Rate of the reaction increases as the particle size decreases. The smaller the
particle size the greater the surface area. Therefore the greater the number of
collisions, the greater the number of successful collisions.
Concentration
The greater the concentration the greater the rate of reaction
This reaction is studied using the reaction between sodium thiosulfate and
Hydrochloric acid.
Na2S2O3 + 2HCl -> S + 2NaCl + SO2 + H2O
If the concentration of the reactants is increased, the number of collisions will also be
increased. If the number of collisions is increased then the number of effective
collisions will be increased.
Temperature
The greater the temperature the greater the rate of reaction
This reaction is studied using the reaction between sodium thiosulfate and
Hydrochloric acid.
Na2S2O3 + 2HCl -> S + 2NaCl + SO2 + H2O
The rate of a reaction increases as the temperature increases because more of the
colliding molecules have the minimum activation energy needed to react.
Catalysts
A catalyst increases the rate of a chemical reaction by lowering the activation energy
of the reactants. A catalyst speeds up the rate of a chemical reaction but does not take
part in the reaction itself.
General properties of catalysts
 Catalysts are unchanged at the end of a reaction.
 Catalysts tend to be specific. I.e. Lipase breaks down fats
Pepsin breaks down protein
Amylase breaks down starch.
 Catalysts need only be present in small amounts.
 In the case of equilibrium reactions, catalysts increase the rate at which
equilibrium is reached but have no effect on the position of equilibrium.
 Catalytic poisons can destroy catalysts.
Types of Catalysis.
Homogeneous catalysis. This is catalysis in which both the reactants and the catalyst
are in the same phase. (Iodine snake experiment)
Heterogeneous catalysis. This is catalysis in which the reactants and catalyst are in
different phases. (Hydrogen peroxide (liquid) and Manganese dioxide (solid))
Auto catalysis.
One of the products in the reaction catalyses the reaction. (Permanganate ions and
Fe2+ ions.)
Mechanisms of catalysis
Intermediate compound theory.
A + B  AB
SLOW
A + C  AC
FAST
AC + B AB + C
FAST
The decomposition of hydrogen peroxide catalysed by the presence of I ions (iodine
snake reaction) illustrates the formation of an intermediate.
Also
Sodium Hydrogen tartrate +Hydrogen peroxide + Co2+ ions (pink)
Pink
Blue/Green
Pink
Cobalt
Intermediate
Cobalt
Surface Adsorption Theory
Platinum
2H2 + O2 
2H2O
A good example is the reaction of Hydrogen and Oxygen to form water using finely
divided Platinum as the catalyst.
The Hydrogen and oxygen molecules settle on the surface of the catalyst. The
adsorbed atoms form weak bonds with the metal atoms. Transition metals can act as
catalysts because they have vacant d orbitals.
The hydrogen and oxygen molecules then react to form water.
The products leave the surface of the catalyst. (desorption)
Catalytic converters.
Exhaust fumes contain carbon monoxide (CO), nitrogen monoxide (NO), nitrogen
dioxide (NO2) and unburnt hydrocarbons. a catalytic converter converts these gases to
environmentally friendly gases.
The catalytic converter consists of a thin coating of platinum, palladium, and rhodium
on a ceramic or metal honeycomb inside a stainless steel case.
Temp= 300 C
2CO + 2NO
Pt/Pd/Rh

2CO2 + N2
The unburnt hydrocarbons react with oxides of nitrogen to form carbon dioxide
nitrogen and water..
Collision Theory
For a reaction to occur, the reacting particles must collide with each other.
For the formation of product a certain minimum energy is required in the collision.
Such a collision is called an effective collision.
The Activation energy is the minimum energy which colliding particles must have
for a reaction to occur.
Energy profile diagram.
A catalyst works by reducing the activation energy.