Chapter 16 - Reaction rates
Chemical reactions occur at different speeds.
Most reactions which we have studied, so far,
have been almost instantaneous , including acidbase reactions and redox reactions.
The rate of a chemical reaction is defined as,
the rate of change of concentration of either a
product or a reactant per unit time.
3. Particle size.
 If a quantity of reactant is present, in a
reacting vessel, as a single lump of material, it
will react more slowly than if it was crushed
into tiny pieces.
Reason;
 because the ground-up material has a much
greater surface area available for contact
with other reactant molecules.

There are a number of factors which influence
the rate of a reaction:
1.
2.
3.
4.
5.
Concentration of the reactants.
Nature of the reactants.
Particle size.
Temperature.
Catalysts.



1. Concentration of reactants.
Large number of molecules in a given volume.
More collisions between reacting molecules.
Rate (speed) of reaction is increased.
2. Nature of the reactants.
In general, reactions involving ionic compounds in
solution are faster than covalent compounds.
Reason;
 Ionic compounds, in solution, are dissociated
into individual ions.
 They are ready for combining with other ions,
almost instantaneously, without any further
effort.
 Covalent compounds, however, don’t dissociate
(break-up) easily.
 In a reaction involving covalent compounds,
considerable effort must be expended on bond
breaking before any reaction can take place.
This has the effect of slowing down any
reaction.

4. Temperature
Increasing the temperature of a chemical
reaction will increase the rate of the reaction.
As a rough guide, a 10K rise in temperature
will double the rate.
The number of collisions doesn’t increase
much. However, the effectiveness of each
collision is greatly increased because of the
increased kinetic energy, i.e. when two
molecules collide there is a greater chance
of a reaction taking place.
5. Catalysts
A catalyst is a substance which alters the rate of
a chemical reaction without being used-up itself.
Catalysts have useful properties.
 They are not chemically altered and may be
re-used. They may, however, be physically
changed (from a large lump to powder form).
 They are specific, that is, a catalyst acts on
only one reaction or a specific group of related
reactions.
 Catalysts need be present only in small
quantities to be effective. Increasing the
amount of catalyst doesn’t always significantly
increase the reaction rate.
 Catalysts do not affect the position of
equilibrium but speed up the time taken to
establish equilibrium.
 Catalysts can be rendered useless if
contaminants are introduced. These are
referred to as catalytic poisons.
Types of catalysis
1. Homogeneous catalysis.
This is where the catalyst and the reactants are
in the same phase (no boundary between the
catalyst and reactants).
2. Heterogeneous catalysis.
Catalyst and reactants in different phases.
Examples; hydrogen peroxide decomposes to
form water and oxygen gas according to the
following equation,
2H2 O2  2H 2 O  O2
This can be catalysed by powdered manganese
dioxide, in which case heterogeneous catalysis
occurs, or by potassium iodide solution, in which
case homogeneous catalysis takes place.
Autocatalysis
This is where a product of a reaction is also
capable of catalysing the reaction.
example :
MnO 4  8H   5Fe  2  Mn  2  5Fe 3  4H 2 O
in this reaction Mn 2 is a product but also
catalyses the reaction.
Mechanisms of catalysis
There are two types of mechanisms of catalysis,
 The intermediate complex theory.
And
 The surface adsorption theory.
The intermediate complex theory.
This type of catalysis is occurs in two stages.
1. One of the reactants (A) combines with
the catalyst (CAT) to form a temporary
unstable intermediate called a complex.
A  CAT  A  CAT(complex)
2. The intermediate complex then reacts with
the other reactant (B) to form the
eventual products (C and D), with the
catalyst remaining unchanged.
A  CAT  B  C  D  CAT
An example of this type, is the decomposition of
hydrogen peroxide by potassium iodide (catalyst).
This reaction may be represented by the following
series of steps,
(1)H 2 O 2  I   H 2 O  IO  (int ermediate)
(2)H 2 O 2  IO   H 2 O  O 2  I 
Surface adsorption theory
This type of catalysis occurs when the reacting
substances are adsorbed onto the surface of a
catalyst. The bonds of the reactants are loosened
and they are free to react with each other. The
process takes place in a number of stages:
 Adsorption.
 Reaction.
 Desorption.
Stage 1: The molecules are adsorbed onto the
surface of the catalyst.
H-H
O-O
H-H
Pt catalyst
This does two things,
 It concentrates the reactants in a small
area.
and
 The bonds of the reacting molecules loosen
as they form temporary bonds with the
catalyst. This makes reaction easier.
Stage 2: Reaction occurs.
The reacting species react rapidly once the bonds
have been loosened.
Unreacted
Hydrocarbons.
Lead
H
H
O
O
H
H
Stage 3: Desorption.
The new products are released off the surface of
the catalyst.
H 2O
hydrocarbons
The reaction represented by the process above
is,
Pt
2H 2  O 2 
2H 2 O
Catalytic converters.
Exhaust fumes from cars include,
Pollutant
CO
NO
NO 2
Source
Incomplete
combustion of
hydrocarbons.
N 2 in air is
Toxic oxides.
oxidized in high
temperatures of
engines, to form
NO and NO 2 on
Dissolve in
rain water to
form acids.
further
oxidation.
Problems
Highly
poisonous
Can give rise
to smog.
Toxic.
Cumulative
Poison.
Brain
damage.
Catalytic converters remove these pollutants from
the exhaust fumes. The process may be
represented by the following schematic diagram,
CO
NO
NO 2
H 2O
Fuels which
escape
combustion in the
engine.
Fuels containing
lead.
CO 2
N2
H 2O
The exhaust fumes pass through a ceramic
honeycomb, the surface of which is coated with
platinum, palladium and rhodium catalysts.
The mixture of hot gases is converted to
harmless gases by the catalysts as they pass
through the converter. The converter is encased
in steel. They last for about 50,000 miles. Lead
tends to poison the catalyst (stop it from doing
its job).
Activation energy.
This is the minimum amount of energy which
colliding particles must have for a reaction to
occur.
Activated complex
E
N
E
R
G
Y
Reactant
s
Activation
energy
Products
Reaction profile diagram
Chemical reactions can be explained in terms of
the collision theory. Molecules in a reaction vessel,
in the liquid or gas phase, are constantly moving
around. These molecules will collide with each
other. If they collide with enough energy
(activation energy)they will stick together, to
form an activated complex, and eventually form
products.