The rate of a reaction is dependent on the number of collisions between particles that occur each
second with sufficient energy to react (i.e. with energy that exceeds the activation energy). The
distribution of energies of particles in a sample is shown by the Maxwell-Boltzmann curve. A typical
curve for a gas a room temperature is given below together with a value for the activation energy of
a typical reaction. As can be seen below, only a small proportion of the molecules have an energy
that exceeds the activation energy. That is, only a small proportion of collisions will occur at room
temperature with sufficient energy to react. If molecules collide with less than the minimum required
energy (i.e. less than the activation energy) they will simply bounce off each other and will not form
products. Successful collisions will result in reaction and energy will be released as new bonds are
made.
**Note: Rate of reaction = (change in concentration)/(change in time).
The Kinetic Molecular Theory proposes that there is a distribution of molecular kinetic energies for a
substance at a given temperature and that the mean kinetic energy of the molecules is proportional
to the absolute temperature.
Clearly, as temperature increases, the percentage of ‘reactive’ molecules increases. It follows that
as temperature increases:
•
Collision frequency increases
•
The proportion of molecules having the minimum kinetic energy necessary for reaction to
occur increases
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Unit 3 & 4 Chemistry – Notes
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**Note: In the given graph, two distribution curves are shown where T2 > T1 and it is assumed that
the area under the whole curve is the same for both temperatures i.e. the same number/population
of molecules. The area under the graph is proportional to the number of particles present.
**Note: For many reactions, the rate of reaction doubles for every rise in 10oC.
Now, if we consider the activation energy Ea – the minimum kinetic energy required to break the
bonds in the reactants – the proportion of the population able to react at T1 is given by the blue area.
However, at the higher temperature T2, the proportion with enough kinetic energy to react is given
by the combined blue area plus the red area. Therefore, because of the shift in the distribution at
the higher temperature T2, a greater proportion of particles have the minimum kinetic energy to
react and hence a greater chance of an effective collision happening i.e. reactant molecule bonds
breaking to form products.
**Note: Despite the proportion of energetically qualified molecules increasing with increase in
temperature, not all the possible collisions are effective collisions. The alignment of the energetically
qualified molecules must also be correct for a given collision to be an effective (or fruitful) collision.
The relative orientations of the reactants must allow formation of any new bonds necessary to
produce products.
The methods of measuring the rate of reaction include:
•
Measuring the volumes of gas evolved using a gas syringe over a convenient period of time
•
A conductivity meter may be used if there is a change in conductivity during the reaction (i.e. if
the number of ions present is changing)
•
The pH of the reaction can be monitored over time using a pH probe if there involves the
consumption of H+ ions
•
Measuring the change in pressure over time for gas reactions
•
Colorimetry may be used if one of the species is coloured as it can monitor the change in
intensity of colour
•
Titration of aliquots of reaction mixture against a primary standard or standard solution
 The School For Excellence 2019
Unit 3 & 4 Chemistry – Notes
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1.
Concentration
It is generally true to say that reactions are accelerated if the concentrations of the reactants
are increased. This is due to an increased collision frequency at higher concentrations which
increases the number of effective collision and therefore the rate at which products are formed.
2.
Pressure
An increase in pressure by decreasing the volume results in the particles having less space in
which they can move. This increases the collision frequency and thus the number of effective
collisions. The reaction rate is therefore increased.
**Note: Increasing the total pressure increases the rate of reaction. Increasing the partial
pressure of a reactive gas also increases reaction rate as it increases collision frequency.
3.
Temperature
Increasing the temperature of a reaction increases the average speed of the particles and
therefore their collision frequency. HOWEVER, this is not the only contribution affecting rate
otherwise every gaseous reaction would be almost instantaneous. Far more important is that a
temperature rise increases the proportion (or fraction) or particles that possess enough energy
to exceed the activation energy and therefore form products. The MINIMUM quantity of energy
required to break the bonds in the reactants is the activation energy. This minimum collision
energy is the energy required to activate the molecules into a state from which reactant bonds
can change into product bonds.
4.
Surface Area
Maximisation of contact area between reactant molecules results in reaction rate increase
because collision frequency is increased.
5.
Catalysts
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Unit 3 & 4 Chemistry – Notes
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A catalyst is a non-consumed reaction accelerator. Catalysts are often said to lower the
activation energy of a reaction system by providing an alternative pathway. In general,
reactions with catalysts proceed from reactants to products via unstable, high energy
intermediates known as activated complexes. The actual mechanisms involved in catalytic
action are quite complex and not fully understood. The most common catalysts are present as
solids (usually solids of high surface energy like certain metals, particular ionic compounds
etc.). The regions of incompletely bonded atoms on a surface can act as active sites for other
substances to become attached (adsorbed). If substances become attached to a surface,
albeit temporarily, via what is called chemisorption, electrons must be involved. This results in
a different activated complex for the reaction, where the original bonds in the reactants are
weakened and are more easily broken than in the absence of the catalyst. This different
activated complex is produced of lower energy requirements and the reaction occurs via an
alternate pathway.
**Note: Homogeneous catalysts are in the same state as the reactants and products.
**Note: Heterogeneous catalysts are in a different state to the reactants and products.
**Note: Catalysts DO NOT affect the equilibrium yield as they increase the rate of forward and
back reactions EQUALLY.
6.
Light
Some reactions are accelerated when irradiated with light of suitable wavelength. Light is
radiant energy, and absorption of photons of appropriate energy value by reactant species is
the mechanism by which light-induced reactions occur.
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Unit 3 & 4 Chemistry – Notes
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