RATES OF REACTION
Increasing the rates of chemical reactions is important in industry because it helps to
reduce costs.
The rate of reaction is the speed at which a chemical change takes place. It is followed
by measuring the rate at which reactants are used up or the rate at which products are
formed. This allows a comparison to be made of the changing rate of a chemical
reaction under different conditions.
Rate = amount of reactant used
time
or
Rate = amount of product formed
time
Chemical reactions can occur only when reacting particles collide with each other with
sufficient energy. The minimum amount of energy which particles must have to react is
called the activation energy.
Various factors alter the rate of a reaction:
 the state of division of the reactants (i.e. the particle size of a solid): the smaller
the particle size, the faster the reaction is.
 the concentration of dissolved reactants or the pressure of gases: the higher the
concentration or the pressure, the faster the reaction is.
 the temperature of the reaction mixture: the higher the temperature, the faster the
reaction is.
 the addition of a catalyst speeds up the reaction
1. The Particle Size of Solid Reactants
The reaction used to study the effect of particle size is the reaction of calcium
carbonate, in the form of marble, with dilute hydrochloric acid.
CaCO3(s) + 2HCl(aq)
CaCl2(aq) + H2O(l) + CO2(g)
The reaction is followed by the change in mass of the reaction flask with time as carbon
dioxide is given off. It would also be possible to measure the change in the volume of
carbon dioxide given off with time by collecting the gas in, for example, a gas syringe.
Method:
A constant mass of marble chips was weighed out and placed into a 250cm 3 conical
flask. 100cm3 of 1M hydrochloric acid was added from a measuring cylinder. The flask
was loosely stoppered with cotton wool (to allow the gas to escape but to prevent the
loss of liquid splashes) then placed onto an electronic balance. The mass of the conical
flask was recorded every 15s for the first minute, then every 30s for a total of ten
minutes. It is assumed that the temperature of the reaction mixture stayed constant.
The experiment was repeated with three different sizes of marble chips, keeping all
other variables the same. The results were tabulated and a graph of mass of carbon
dioxide (y-axis) against time (x-axis) was plotted.
TOPIC 11.2.4: RATES AND ENERGY 1
Results:
Large Particles
Time /s
Mass of
flask/g
Mass of
CO2 /g
0
15
30
60
90
120
150
180
210
240
270
300
330
360
390
420
450
480
510
540
570
600
TOPIC 11.2.4: RATES AND ENERGY 2
Medium Particles
Mass of
flask /g
Mass of
CO2 /g
Small Particles
Mass of
flask /g
Mass of
CO2 /g
Interpretation of Results:
The reaction occurs only on the surface of the calcium carbonate. For a given mass of
calcium carbonate, the smaller the size of the particles, the greater is the surface area
and the faster the reaction. This is shown on the graph by the gradient becoming
steeper as the particle size decreases.
The smaller the particle size (the bigger the surface area) , the
faster the reaction
Since the same quantities of reactants are involved in all three reactions, the same
mass of carbon dioxide is given off in all of them, if the reaction is allowed to go to
completion. This is shown by all three curves levelling off at the same total mass of
carbon dioxide.
The calcium carbonate is in excess, so the hydrochloric acid is used up completely as
the reaction takes place. Since its concentration decreases with time, the reaction
becomes slower and slower. This shown on the graph by a curve of steadily decreasing
gradient.
The expected mass of carbon dioxide can be calculated. Since the calcium carbonate is
in excess, the mass of carbon dioxide depends on the amount of hydrochloric acid
used. 100cm3 of 1M HCl contains 0.1mole of HCl.
CaCO3(s) + 2HCl(aq)
2 moles
0.1 mole
0.1 mole
0.1 mole
TOPIC 11.2.4: RATES AND ENERGY 3
CaCl2(aq) + H2O(l) + CO2(g)
1 mole
0.05 mole
0.05 x 44 g
2.2g
2. The Temperature
The reaction used to study the effect of temperature is the reaction of sodium
thiosulphate solution with dilute hydrochloric acid.
Na2S2O3(aq) + 2HCl(aq)
2NaCl(aq) + S(s) + H2O(l) + SO2(g)
The reaction is followed by the appearance of colloidal sulphur as the reaction
proceeds. The formation of sulphur begins as soon as the reactants are mixed, but it
takes time for observable amounts to be produced. The time taken to reach a particular
point in the reaction can be determined by standing the reaction flask on a piece of
paper marked with a feint cross and timing how long it takes for the cross to be
obscured when looked at from above.
Method:
10cm3 of sodium thiosulphate stock solution (concentration 40g per litre) were
measured out into a measuring cylinder and poured into a 250cm3 conical flask. 40cm3
of water were measured out in a similar way and added to the flask. The flask was
heated gently over a Bunsen burner to a temperature slightly above the desired
temperature. The flask was placed on a piece of paper marked with a feint cross. 5cm3
of 2M hydrochloric acid were measured out into a second measuring cylinder. As the
acid was poured into the conical flask, a stopwatch was started and the mixture gently
swirled. The initial temperature was recorded. The stopwatch was stopped when the
cross, viewed from above, became obscured. The time and the final temperature were
recorded. The mean of the initial and final temperatures was taken as the temperature
of the reaction.
The experiment was repeated for five different temperatures, keeping all other variables
constant.
The results were tabulated and from the results a graph of 1000/ time (y-axis) against
temperature (x-axis) was plotted. (1000/ time is a measure of the rate)
Results:
Initial temp.
/ oC
Final temp.
/ oC
TOPIC 11.2.4: RATES AND ENERGY 4
Mean temp.
/ oC
Time /s
1000/time /s-1
Interpretation of Results:
The graph is an exponential curve: the rate of reaction increases rapidly with
temperature, a rise in temperature of 10oC approximately doubling the rate.
The higher the temperature, the faster the reaction
In the reaction between thiosulphate
ions and hydrogen ions, as the ions
collide covalent bonds are broken
and new bonds are formed. Since
energy is needed to break bonds,
the colliding particles must have a
minimum
energy
on
collision
sufficient to break the bonds. This
energy is known as the activation
energy (Ea).
Only those collisions with energy
greater than or equal to the
activation energy will result in a
reaction.
bonds
breaking
bonds
forming
Ea
Chemical
Energy
reactants
H
products
Reaction path
An increase in temperature increases the rate of reaction in two ways:
The particles collide more energetically
The particles of a substance have a range of different energies, the average energy
being proportional to the temperature in Kelvins. As the temperature increases, the
particles move faster (i.e. they have more kinetic energy), and the proportion of particles
with higher energies increases. Therefore, the number of collisions with energy greater
than or equal to the activation energy rises rapidly as the temperature increases, and so
the rate rises rapidly. This is the major effect.
The particles collide more frequently
The particles move around faster and therefore there is a greater chance that they will
be involved in a collision. This is a minor effect.
TOPIC 11.2.4: RATES AND ENERGY 5
3. The Concentration of Reactants
The reaction used to study the effect of concentration is the reaction of sodium
thiosulphate solution with dilute hydrochloric acid.
Na2S2O3(aq) + 2HCl(aq)
2NaCl(aq) + S(s) + H2O(l) + SO2(g)
The reaction is followed by the appearance of colloidal sulphur as the reaction
proceeds. The formation of sulphur begins as soon as the reactants are mixed, but it
takes time for observable amounts to be produced. The time taken to reach a particular
point in the reaction can be determined by standing the reaction flask on a piece of
paper marked with a feint cross and timing how long it takes for the cross to be
obscured when looked at from above.
Method:
50cm3 of sodium thiosulphate stock solution (concentration 40g per litre) were
measured out into a measuring cylinder and poured into a 250cm 3 conical flask. The
flask was placed on a piece of paper marked with a feint cross. 5cm 3 of 2M hydrochloric
acid were measured out into a second measuring cylinder. As the acid was poured into
the conical flask, a stopwatch was started and the mixture gently swirled. The stopwatch
was stopped when the cross, viewed from above, became obscured. The time was
recorded.
The experiment was repeated for five different concentrations, keeping all other
variables constant.
A graph of 1000 /time (y-axis) against volume of Na2S2O3 (x-axis) was plotted.
Since the total volume of the reaction mixture is constant (at 55cm 3) the concentration
of Na2S2O3 is proportional to its volume.
Results:
Constants:
total volume of Na2S2O3 solution 50cm3
2M hydrochloric acid
5cm3
temperature
18oC
conical flask & cross
Volume of Na2S2O3
(40g/l) /cm3
50
40
30
20
10
Volume of H2O
/cm3
0
10
20
30
40
TOPIC 11.2.4: RATES AND ENERGY 6
Time /s
1000/ time /s-1
32
40
52
79
172
31.3
25.0
19.2
12.7
5.8
x
30
x
20
x
1000/ time /s-1
x
10
x
0
10
20
30
Volume of Na2S2O3 /cm3
40
50
Interpretation of Results:
The graph is a straight line through the origin, therefore the rate of the reaction is
directly proportional to the concentration of sodium thiosulphate.
The more concentrated the solution, the faster the reaction
To react, the reacting particles must collide; therefore the rate will be faster the greater
the number of collisions there are in a given volume in a given time. The more
concentrated the solution is, the greater the number of particles there are in a given
volume and therefore the greater the frequency of collisions. It is important to note
that only a very small proportion of the total number of collisions is successful and leads
to a reaction.
Effect of Pressure
Pressure is important only in reactions involving gases. Pressure affects gaseous
reactions in the same way that the concentration affects reactions in solution. As the
pressure is increased, the greater the number of particles there are in a given volume
and therefore the greater the number off collisions in a given time. Therefore, as the
pressure increases, the rate increases.
TOPIC 11.2.4: RATES AND ENERGY 7
4. Addition of a Catalyst
The reaction used to study the effect of a catalyst is the decomposition of hydrogen
peroxide:
2H2O2
2H2O + O2
The reaction is catalysed by several metal oxides; the compound used here is
manganese(IV) oxide. The reaction is followed by collecting the oxygen given off and
measuring its volume at regular intervals of time. The gas may be collected either in a
gas syringe or over water in a burette.
Method:
7cm3 of 20 volume hydrogen peroxide were measured out using a measuring cylinder
and poured into a 100cm3 conical flask. 43cm3 of water were measured out in a similar
way and added to the conical flask. 0.5g of powdered manganese(IV) oxide was
weighed out and added to the flask. The flask was connected by a delivery tube to a
100cm3 gas syringe. A stopwatch was started. The volume of oxygen, which had
collected in the gas syringe, was recorded at 30 second intervals for ten minutes.
Keeping all other variables constant, the experiment was repeated using 0.75g and 1.0g
of the same manganese(IV) oxide.
Results:
Constants:
Time /s
0
30
60
90
120
150
180
210
240
270
300
330
360
390
420
450
480
510
540
570
600
20 volume H2O2
water
0.5g of MnO2
0
1.8
4.1
6.2
8.3
9.9
11.5
12.8
14.4
16.2
17.5
18.8
19.9
21.7
22.8
23.5
25.0
26.0
27.0
27.7
28.8
TOPIC 11.2.4: RATES AND ENERGY 8
7cm3
43cm3
temperature
Volume of O2 /cm3
0.75g of MnO2
0
2.9
6.0
7.9
10.3
12.1
13.8
15.4
16.8
18.2
19.7
20.8
22.1
23.1
24.2
25.2
26.3
27.4
28.3
29.0
30.1
18oC
1.0g of MnO2
0
3.2
5.7
9.4
12.0
14.6
17.1
19.0
21.1
23.3
25.1
26.8
28.5
30.0
31.6
32.9
34.2
35.5
36.6
37.7
39.0
All three reactions are being observed only in the early stages and are far from
complete. Since all three reactions use exactly the same quantity of hydrogen peroxide,
they should, when complete, give the same total volume of oxygen.
A graph of volume of oxygen (y-axis) against time (x-axis) was plotted.
Interpretation of Results:
The reaction is faster the more catalyst there is present; this is shown by the increasing
steepness of the curves.
The reaction takes place on the surface of the catalyst. Increasing the mass of catalyst
increases the surface area and therefore speeds up the rate.
A catalyst does not affect the outcome
of a reaction; the same product is
formed but in a shorter time. A catalyst
works by weakening bonds, which
lowers the activation energy for the
reaction. If the activation energy is
lowered, more particles have enough
energy to react and therefore the
reaction goes faster.
A catalyst is not used up in the reaction
but is recovered unchanged at the end.
Ea
Chemical
Energy
Ea cat
reactants
H
products
Reaction path
A catalyst is a substance which increases the rate of a chemical reaction but is
not used up in the reaction.
A catalyst can be used over and over again to speed up the conversion of reactants to
products. Different reactions need different catalysts.
TOPIC 11.2.4: RATES AND ENERGY 9
ENERGY & CHEMICAL CHANGE
During a chemical reaction, existing chemical bonds are broken and new bonds are
formed.
 energy must be supplied to break existing bonds
 energy is released when new bonds are formed
The overall change in energy, which takes place during the reaction, is called the heat
of reaction and is given the symbol H (pronounced ‘delta H’).
If the energy needed to break existing bonds is less than the energy released when new
bonds are formed, the reaction will give out energy (usually as heat) to its surroundings.
The temperature of the reaction mixture will rise. This is called an exothermic reaction.
For an exothermic reaction, the sign of the heat change (H) is negative. Ea is called
the activation energy and is the minimum energy needed for a reaction to take place.
When a catalyst is added to a reaction, it provides an alternative reaction path with a
lower activation energy. If the activation energy is lowered, the reaction goes faster.
All reactions involving burning, e.g. the burning of magnesium, are exothermic.
bonds
breaking
Ea
bonds
forming
Chemical
Energy
reactants
H
products
Reaction path
We can make use of exothermic reactions is devices such as hand warmers and self
heating cans where heat may be needed in a remote environment.
TOPIC 11.2.4: RATES AND ENERGY 10
If the energy needed to break existing bonds is greater than the energy released when
new bonds are formed, the reaction will absorb energy (usually as heat) from its
surroundings. The temperature of the reaction mixture will fall. This is called an
endothermic reaction. For an endothermic reaction, the sign of the heat change (H) is
positive.
Ea
Chemical
Energy
products
H
reactants
Reaction path
The reaction of sodium hydrogencarbonate with hydrochloric acid is endothermic.
NaHCO3 + HCl
NaCl + H2O + CO2
We can make use of endothermic reactions is devices such as chemical cold packs to
put on an injury to reduce swelling and numb pain.
If a reversible reaction is exothermic in one direction it will be equally endothermic in the
other. Eg
CuSO4.5H20
CuSO4.
+
5H20
You need to heat blue hydrated copper sulphate to produce white anhydrous copper
sulphate so this process requires energy and is therefore endothermic.
However, when water is added to white anhydrous copper sulphate a lot of heat is given
out as it turns blue and so is exothermic.
TOPIC 11.2.4: RATES AND ENERGY 11