The document outlines factors affecting the rate of chemical reactions, emphasizing the need for reactant particles to collide with sufficient energy to overcome activation energy and form new bonds. Key factors for increasing reaction rates include increasing surface area, concentration, temperature, and using catalysts. The presentation serves as a refresher for A-level chemistry students, explaining how to measure reaction progress and the effects of various parameters on reaction rates.

1. Factors Affecting the Rate of a Chemical
Reaction
The following events must occur before a
reaction can proceed:
1. The reactant particles must collide with each
other.
2. The collisions must be of enough energy to
overcome the ‘energy barrier’, called the activation
energy (more about this on the next slide).
3. The reactants must form new bonds to produce
products.
This presentation is useful as a refresher for those of you
starting Unit 2 of the A-level chemistry course.

2. Activation Energy 1
Defined as:
The minimum energy required to
bring about a chemical reaction.
Gasoline for your car would ignite as soon as it came into contact with
air.
You would burst into flames.
Trees would spontaneously combust.
If there were no such thing as ‘activation
energy’ life would be very difficult:
Activation energy is why these things do not happen, there is
an energy barrier so most reactions need to be ‘started off’ by
putting in some energy.

3. Activation Energy 2
Activation energy for a reaction is shown on
reaction profile diagrams (do you remember
these?).
Reactants
Products
Activated intermediate
Activation
energy
H
energy

4. Changing the Rate of a Chemical Reaction
To change the rate of a reaction one or more of the
following things must happen:
1. Increase the number of collisions between
the reactant particles
2. Increase the energy of the collisions.
3. Decrease the activation energy.
This is all very well but how
can we follow the progress of
a chemical reaction?

5. Following a Chemical Reaction 1
To find the rate of a chemical reaction we
must be able to follow its progress with time.
We have two choices:
1. Record the increase in product concentration as
the reaction progresses.
2. Record the decrease in reactant concentration
as the reaction progresses.

6. Following a Chemical Reaction 2
As an example consider the reaction between calcium carbonate and
hydrochloric acid.
You should already know the equation but here it is:
CaCO3 + 2HCl  CaCl2 + H2O + CO2
We can follow this reaction by measuring the volume of carbon
dioxide produced as the reaction proceeds.
This apparatus can be used
to measure the gas as it is
formed. It is not the only
way, look in your text book
for more details.
Dilute acid
Marble
chips
Gas being collected

7. Following a Chemical Reaction 3
If you collect data for the total amount of gas produced as the
reaction progresses then plot this data on a graph you should get a
curve similar to that shown below.
All very well, but what
does the graph tell you?
The gradient or slope of the graph
shows the rate of the reaction.
Steeper slope = faster reaction.
Volume
of
gas/cm
3
Time/sec.
Reaction fastest at
the beginning.
Slowing down.
Reaction
finished.

8. Effect of Surface Area 1
When solids take part in chemical reactions only the surface
particles are exposed so they are the only ones that can collide
with particles of other reactants.
The surface particles
are ‘exposed’ and can
react.
‘Inner’ particles are protected and
cannot collide with other particles
until they become ‘exposed’.

9. Effect of Surface Area 2
If we break up this ‘lump’ into smaller
pieces the number of particles has not
changed but the there are now more
‘surface’ particles.
There is now a greater surface
area with more exposed particles
so more collisions can occur, hence
faster reaction.
Larger surface area =
Larger surface area =
faster reaction.
faster reaction.

10. Acid Particles
Zinc
Effect of Concentration
Consider the reaction between zinc and hydrochloric
acid: Zn + 2HCl  H2 + ZnCl2
(How could you follow the progress of this reaction? Click to find out)
1M hydrochloric acid 2M hydrochloric acid
There are more particles of acid per unit volume in the 2M acid than
there are in the 1M acid. So, there will more collisions between the
acid and zinc particles in the stronger acid, giving a faster reaction.
Higher concentration = faster reaction

11. Gas Reactions
The rate of reaction between gases is increased
by increased pressure.
In effect pressure is the gas equivalent of
concentration.
These two gas jars
contain the same
number of gas particles.
Low pressure,
particles far apart.
Higher pressure,
particles closer
together.
The higher pressure jar has
more particles per unit
volume which means a
higher concentration,
hence faster reaction.
Higher pressure = faster reaction
Higher pressure = faster reaction

12. Effect of Temperature
According to kinetic theory (do you remember
this?) as the temperature increases the particles in
a substance move about more quickly.
Reaction at 300
C Reaction at 500
C
As the temperature increases the number of collisions increases as well
as the energy of the collisions. So temperature has a big effect on the rate
of reaction. For every 100
C increase the rate approximately doubles.
Higher temperature = faster reaction

13. Effect of a Catalyst 1
A catalyst is a substance that increases the speed of a
reaction, without being used up. A catalyst can be
‘recovered at the end of a reaction and used again.
A catalyst reduces the
activation energy of a reaction.
This is easier to understand with a diagram –
see next slide.

14. Effect of a Catalyst 2
Activation energy
without catalyst.
Activation energy
with catalyst.
energy
The lower activation energy in the
presence of a catalyst means the
reaction will be faster. More of the
collisions have enough energy to
react. There is a lower ‘energy barrier’.
Catalyst = faster reaction.

15. More About Rate Graphs
Volume
of
gas/cm
3
Time/sec.
Reaction fastest at
the beginning.
Slowing down.
Reaction
finished.
1. Why is the reaction
fastest at the beginning?
2. Why does the reaction
slow down?
3. Why does the reaction
eventually stop?
1. This is where the concentration of the reactants is highest, therefore
fastest reaction.
2. As the reactants are used their concentration decreases so the rate of
reaction decreases.
3. One of the reactants is used up, so there can be no further reaction.

16. Special Note
Some exothermic reactions speed up shortly after they
start, this might be unexpected, but think about it!
The temperature increases and this overcomes, at least
to begin with, the effect of reducing the concentration.
So, in some cases the reaction will speed up then slow
down and eventually stop.
Do not get caught out by this. Questions
related to this effect are very common!

17. Summary
1. Increasing the surface area gives a faster reaction
because more particles are ‘exposed’ to the other
reactant.
2. Increasing the concentration increases the rate of
reaction because there are more collisions between the
reactant particles.
3. Increasing the temperature increases the rate of
reaction because the particles move move quickly and
so collide more often and with greater energy.
4. A catalyst increases the rate of a reaction because it
reduces the activation energy so more of the collisions
have enough energy to react.

18. The End
The End
Constructed by:
Constructed by:
Rob Dickens, Doha College Science Department.
Rob Dickens, Doha College Science Department.