1
REACTION RATES
1.1 What is a chemical reaction?
A chemical reaction is a change that involves the formation of at least one new substance.
Reactions only occur between certain substances. Not every chemical substance reacts with
any other. The drive for one substance to react with another is for greater stability, for
example, pure iron is not stable and will react with oxygen to form iron oxide – rusting.
Reactions can be detected:

if a gas is evolved, as long as the temperature is low enough that the gas is not simply
one of the components boiling, eg hydrochloric acid and zinc
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if a solid (precipitate) is formed when two solutions are mixed eg silver nitrate and
sodium chloride
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if there is a change in colour eg potassium permanganate and hydrogen peroxide
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if there is a significant change in temperature eg burning of magnesium
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if there is disappearance of a solid, not attributed to dissolving eg magnesium hydroxide
and hydrochloric acid
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by instrumental methods which can show the presence of materials different to the
reactants, where there are no other obvious changes
1.2 What causes a reaction to occur?
The current understanding of how chemical reactions occur is known as the collision
theory. It states that a reaction can only occur if the reactant molecules collide. If you have
reactant A in one beaker and reactant B in another beaker, then no reaction is possible. If
they are combined, then the reactant molecules are free to mix and to crash into each other.
Even a collision between reactant molecules does not guarantee that a reaction will
occur. Two other factors are critical in determining whether the collision will result in
reaction:

the direction of collision – unless the two reactants are both single atoms, the success of
the collision will depend on which part of each molecule comes into contact, as shown in
Figure 1.1
(a)
(b)
FIGURE 1.1 (a) Successful collision (b) unsuccessful collision
1. Reaction Rates

the energy which each molecule has in the collision – the more energy involved, the
better the chance of success; imagine two cars running into to each other in the
supermarket carpark – very little damage is caused because they are moving slowly;
compare this to a collision on the highway
The amount of energy involved in the collision is important because the reaction must pass
over an energy barrier or hill to get to the product side, as shown in Figure 1.2. This barrier is
known as the activation energy (Ea). Some reactions have very high activation energies and
are less likely to occur than those with lower hills. If the reactant energies are not enough to
get over the Ea hill, the reaction will not occur.
Energy
Ea
Reactants
Products
Progress of reaction
FIGURE 1.2 The energy barrier
Consider a reaction where four reactant molecules are involved, such as the production of
ammonia from nitrogen and hydrogen.
N2 + 3H2  2NH3
Does a successful collision for this reaction involve all four molecules colliding at the same
time with right amount of energy and in the right directions? The chances are very remote
that such an event would occur. What happens is a series of collisions, and the formation of
intermediate species which are part way to the final product. This is shown below in the
following hypothetical reaction sequence, where AB is the intermediate “product”.
A + B  AB
AB + A  A2B
1.3 What affects the speed of reaction?
The speed at which a chemical reaction goes is termed its rate. The reaction rate is
expressed as the number of successful collisions per second, or more easily measured, the
rate of change of concentration (increase for a product, decrease for a reactant).
It is important to realise that not everything in a chemical reaction happens at the same
rate even at the same point in time. Returning to the ammonia example from above,
whenever a molecule of nitrogen reacts, 3 molecules of hydrogen also react and 2 molecules
of ammonia are produced. Hence the rate of disappearance of hydrogen is three times as
great as the rate of disappearance of nitrogen, and the rate of production of ammonia is twice
as great as the rate of disappearance of nitrogen.
Chemistry 2
1.2
1. Reaction Rates
The rate of chemical reaction depends on the following factors. In each case, the effect on the
rate is given, but in most cases, you are required to explain why in terms of the collision
theory and the energy hill.

the reaction itself – the size of the energy barrier differs for different reactions

the concentration of the reactants – the higher the concentration, the faster the reaction

the temperature at which reaction is occurring – the higher the temperature, the faster
the reaction

the physical state of the reactants – liquids react faster than solids; for solids, powders
react faster than large granules

mixing – this increases the reaction rate

presence of a catalyst – a catalyst is defined as a substance that changes the reaction
rate without being consumed in the reaction; it can increase (normal) or decrease the
rate; it works by changing the activation energy (see section 1.4 also)
Chemistry 2
1.3
1. Reaction Rates

pressure of the system (if gas is involved) – higher pressures speed up the reaction

electromagnetic radiation – for some reactants, the presence of UV light speeds the
reaction
1.4 Heterogeneous Catalysis
An important industrial use of catalysts involves heterogenous catalysts, where the reactants
are of a different physical state to the catalyst. Usually this means gaseous or liquid reactants
and a solid catalyst. Many industrial processes use this, including plastics manufacture and
petroleum refining.
The reaction occurs on the surface of the catalyst where one reactant has been adsorbed
– physically stuck to the surface (Figure 1.3a). The other reactant, which is still free to “swim”
around can collide with the adsorbed reactant in only a limited number of ways (Figure 1.3b).
This ensures that the collision direction is more likely to be the correct one for the reaction to
occur. Once the reaction occurs, the products, which are not attracted to the catalyst surface
float away (Figure 1.3c).
(a)
(b)
(c)
FIGURE 1.3 Reaction using a heterogeneous catalyst
Chemistry 2
1.4
1. Reaction Rates
What You Need To Be Able To Do

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
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
list observable signs of a chemical reaction
explain why reactions occur
define the term reaction rate
list factors affecting the reaction rate
explain these effects
describe the use of heterogeneous catalysts
Chemistry 2
1.5