Kinetics
Chemists can find out many different pieces of information about particular
reactions. The most useful are the enthalpy changes associated with reactions and
information about how fast a reaction will occur. Controlling reaction rates and
understanding the processes by which reactants change into products are essential
skills for a chemist.
Reaction rate
You will be familiar with many everyday uses of the term ‘rate’. For example the rate
of movement of a car is referred to as speed: the distance travelled by the car in
relation to the time taken for the journey. The rate of a chemical reaction is
measured as the amount of product made or reactant used up in a certain time.
Collision theory
When a chemical reaction takes place the reactant particles must collide. The theory
which explains the reactions that take place as the result of collisions is called
collision theory. Imagine a gaseous substance in which the particles are in constant
random motion. The particles are continuously colliding with each other and with
the walls of their container. Not all of the collisions between the gaseous particles
result in a reaction.
Collision between identical molecules – no reaction
Molecules collide too slowly – no reaction
Collision at wrong angle – no reaction
The diagram shows that in order for a reaction occur:
• the correct particles must collide
• the collision must be of the correct energy
In order to maximize the energy of the collision the particles must be moving
quickly and collide head on. Reaction conditions can be altered to maximize the
probability of a collision occurring or to increase the energy with which particles
collide. The factors that affect the rates of chemical reactions are
• the temperature of the reactants
• the concentration of the reactants and products
• the surface area of reactants
• the presence of any catalysts
Activation energy
The minimum energy with which particles must collide in order for a reaction to
occur is referred to as the ‘activation energy’. If the particles collide with energy less
than the activation energy then the particles will simply bounce off each other and
no reaction will occur. If the particles collide with energy equal to or greater than the
activation energy the collision is described as successful and a reaction will take
place.
Activation energy and enthalpy changes
The diagram shows the relationship between the activation energy, Ea, of a reaction
and the enthalpy change of the reaction.
The reaction profiles for (a) an
exothermic reaction and (b) an
endothermic reaction. Ea is the
activation energy barrier that
reactants must overcome before
they can change into products.
H indicates the overall
enthalpy change for the
reaction.
The effect of temperature on reaction rate
Temperature is used as a measure of the amount of energy of the particles in a
substance. For example if a sample of gas is heated up to a higher temperature this tells
us that the average amount of energy each particle possesses has increased. You know
that whether a reaction takes place or not is dependent on:
• the energy of the colliding particles
• the number of collisions
Maxwell–Boltzmann distribution of molecular energies
In a sample of gas at a given temperature, the molecules do not all have the same
energy. They are all moving at different speeds. At any instant some of the particles have
a very low energy, a small proportion have a very high energy while the majority of the
particles have an intermediate energy.
In 1859 the Scottish physicist James Clerk Maxwell calculated this distribution of
energies in a sample of gas. His ideas were applied by the Austrian physicist Ludwig
Edward Boltzmann in 1871. The resulting graph showing the distribution of molecular
energies is known as the Maxwell–Boltzmann distribution.
Only molecules in the
shaded area collide with
enough energy to react.
Note the following features of the curve:
• Only very small fractions of the molecules have extremely high or extremely low
energies.
• The curve is not symmetrical; the average energy is to the right of the peak of the
curve.
• The curve passes through the origin.
• The curve does not touch the x (horizontal) axis on the right hand side. It is an
asymptote, a line or curve which approaches another but never touches it.
The shaded area under the right of the curve shows the proportion of molecules that
possess the activation energy (the minimum energy with which particles must collide in
order to react). Only molecules in this portion of the curve are able to react. Each
reaction has its own activation energy. At a given temperature a reaction with higher
activation energy will be slower than one with lower activation energy.
Effect of temperature
The shape of the Maxwell-Boltzmann distribution changes as temperature is altered. As
the temperature increases the energy distribution moves to the right and the height of
the peak decreases. The total area under the curve is constant as this represents the total
number of particles.
For a small increase in temperature the shape of the graph remains broadly the same but
note that the area of the shaded portion has increased, so more molecules have energy
greater than or equal to the activation energy
For a large increase in temperature the shape of the graph alters more
dramatically.
Note that altering the temperature does not have any effect on the value of the
activation energy; this is constant for a given reaction. Increasing the temperature
does not influence this value it increases the energy with which particles collide so
that more of these collisions possess the activation energy.
Catalysis
Many of the reactions essential for our everyday life would not occur without the
presence of catalysts. They can be recovered at the end of the reaction and used
many times over. Industrially, catalysts are used in the manufacture of ammonia,
sulphuric acid, margarines, plastics, and fertilizers. Without the presence of enzymes
(biological catalysts) our bodies would not function, we would not have biological
washing powder, nor have bread to eat or alcohol to drink.
Catalysts
A catalyst can be described as a substance that alters the rate of a chemical reaction
and remains unchanged at the end of the reaction. The catalyst acts by providing an
alternative route of lower activation energy. This reduction in activation energy
enables many more of the collisions between reactants to achieve this minimum
requirement for a reaction to take place. Therefore the rate of the reaction increases.
Catalysts and the Maxwell–Boltzmann distribution
The addition of a catalyst to a reaction has no effect on the energies of the reactant
molecules or on the total number of molecules in the reaction system. It therefore
has no effect on the shape or size of the Maxwell–Boltzmann distribution. The
activation energy line is marked further to the left hand side of the curve to show
the reduction in its value. This has a significant effect on the number of particles that
are found in the shaded area of the curve and can therefore take part in the
reaction.
Heterogeneous and homogeneous
catalysts
There are two important classes of catalysts: heterogeneous and homogeneous.
A heterogeneous catalyst is in a different phase from the reactants.
For example in the hardening of vegetable oils for the production of
margarine a nickel catalyst is used to reduce the activation energy for the
reaction of two gases: hydrogen and an alkene. The reactant molecules attach
themselves to the nickel surface breaking the C=C in the alkene as they attach. The
reaction then takes place on the surface and the alkane molecules formed detach
from the surface. You will learn about the catalytic converter, an important use of
heterogeneous catalysts, when you study the combustion of fossil fuels.
Catalytic converters in the exhaust systems of modern cars
show all the main features of a heterogeneous catalyst. The
catalyst consists of about 2g of finely divided
platinum/rhodium, on a rigid ceramic support. The primary
effect is to catalyse the conversion of the pollutants carbon
monoxide and nitrogen monoxide to carbon dioxide and
nitrogen. 2CO(g) + 2NO(g) r 2CO2(g) + N2(g)
Note that leaded petrol will rapidly poison a catalytic
converter.
A homogeneous catalyst is in the same phase as the reactants. Chlorine radicals
act as a homogeneous catalyst in the upper atmosphere, and have a devastating
effect on the sequence of reactions which take place constantly making and
destroying ozone. One chlorine radical can catalyse as many as one hundred
thousand reactions. The series of reactions that take place is complex and you will
study them in more detail elsewhere. Some of them are given here:
The chlorine radical is represented by Cl•. It destroys ozone in the first reaction and
is regenerated in the second.
The bombardier beetle stores hydrogen peroxide, water, and
noxious substances in an abdominal sac. When threatened,
it injects a catalyst into this mixture. The almost
instantaneous exothermic decomposition of hydrogen
peroxide generates steam, which ejects the contents of the
sac as a hot and highly offensive spray.