3e. How does using a
catalyst affect the speed or rate of a reaction? How does a catalyst work?
GCSE level Chemistry
Notes:
The effect of a catalyst on reaction
rate
(speed)
Doc Brown's
Chemistry KS4 science GCSE/IGCSE/O Level Revision Notes - Factors
affecting the Speed-Rates of
Chemical Reactions
Doc Brown's
chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE chemistry, O level
& ~US grades 8, 9 and 10 school science courses or equivalent for ~14-16 year old
science students for national examinations in chemistry
Rates of
reaction notes INDEX
What next?
Associated Pages
3.
The Factors affecting the Rate of Chemical Reactions
REACTION RATE and
the action of CATALYSTS
Using a CATALYST
3e
The effect of a Catalyst
(see also
light effect and graph 4.8)
Experimental methods for investigating
the effect of a catalyst on the rate of a chemical reaction
Parts of the sections of 1.
Introduction and 2. collision theory are repeated here, but with extra
experimental methods and theoretical details applied to experiments and theories
linked to the effect of using a catalyst on the rate of a chemical reaction
- The apparatus can be used to investigate the how
the speed of the decomposition of hydrogen peroxide varies with different
catalysts.
- The flask and gas syringe system for measuring the rate
of a chemical reaction.
- Oxygen gas is given off which can be collected in the gas syringe
and its volume is used to measure how fast the reaction is going. The grey
'blobs' could represent the solid insoluble catalyst.
- In the diagram above, the white 'blobs'
represent oxygen gas being evolved and the grey lumps the catalyst powder.
-
hydrogen peroxide == catalyst ==> water +
oxygen
-
2H2O2(aq) ===> 2H2O(l)
+ O2(g)
- MnO2 Manganese(IV) oxide
(manganese dioxide'), is a very effective catalyst, but can also try other
transition metal oxides as catalysts like CuO copper(II) oxide.
- The variables to be kept constant are - the concentration of the hydrogen peroxide
solution, the volume of the hydrogen peroxide, the same amount of catalyst
(ideally of the same particle size) and the temperature of the reaction mixture.
- You must also swirl the flask gently to ensure a
good mixing as the reaction proceeds.
- The insoluble particles of the catalyst should
be of the same size to give the same surface area of reactant, BUT, this is
difficult to achieve in practice at school/college level.
- More details of laboratory investigations
('labs') involving 'rates of reaction' i.e. experimental methods for observing
the speed of a reaction and including the effect of a catalyst are given in
the INTRODUCTION
-
In this case, measuring the initial rate of gas formation (see left
and below diagrams) gives a reasonably accurate measure of how fast the reaction is for
that concentration.
-
- The initial gradient, giving the initial rate of
reaction, is the best method i.e. the
best straight line covering several results at the start of the
reaction by drawing the gradient line using the slope of the
tangent from time = 0, where the graph is nearly linear.
- Examples of graph data for two experiments
where one of the reactants is completely used up - all reacted.
- The two graph lines represent two typical
sets of results to explain how the rate of reaction data can be processed.
- Graph A (for a faster reaction) could
represent using a catalyst, but not in Graph B (a slower reaction).
-
- The
set of graphs (above) shows you some typical results.
- The rate of reaction order is
X > E > Y > Z,
- X is an effective catalyst,
substances
Y and Z seem to slow the reaction, acting as inhibitors.
- The more effective the
catalysts, steeper the initial gradient, the faster the reaction.
- The more effective the catalyst, the more it
lowers the activation energy, the more chance of a successful 'fruitful' collision.
- For the effect of a catalyst on the rate of
reaction, under some circumstances graph W could represent the result of
taking twice the mass of solid catalyst or twice the concentration (same volume)
of a soluble reactant, BUT it does depend on which reactant is in excess, so
take care in this particular graph interpretation.
- More details of laboratory investigations
('labs') involving 'rates of reaction' i.e. experimental methods for observing
the speed of a reaction and including the effect of a catalyst are given in
the
INTRODUCTION
-
How can you be sure the 'catalyst' was a true
catalyst?
- The definition of catalyst:
- A catalyst is a substance that increases
the rate of reaction without being chemically changed at the end of the
reaction.
- If you recover the catalyst at the end of a
reaction, and then purify it e.g. wash all the reactants off with pure
water, after this
it should still work as a catalyst.
Theoretical interpretation of the
results of the effect of catalyst on the rate of a chemical reaction
For each factor I've presented
several particle diagrams to help you follow the text explaining how the
particle collision theory accounts for your observations of reaction rate
varying with a catalyst (some 'work' better than others!)

A picture of a particles of
gaseous molecules or molecules/ions in solution undergoing a chemical changes on
the surface of a catalyst - look for the fruitful collision on the catalyst
surface.
-
WHAT IS A CATALYST?
-
HOW DOES A CATALYST AFFECT THE SPEED OF A CHEMICAL
REACTION?
-
HOW DOES A CATALYST WORK?
-
Why does a catalyst speed up a reaction?
-
I was once asked "what is the opposite of
a catalyst? There is no real opposite to a catalyst, other
than the uncatalysed reaction! However, catalysed reactions are very important
in industry and most biochemistry involves enzyme catalysts.
-
The word catalyst means an
added substance, in contact with the reactants, that changes the rate of a
reaction without itself being chemically changed in the end. The catalyst
may temporarily changed though! -
There are the two phrases you may come across:
-
a 'positive catalyst' meaning speeding up
the reaction (plenty of examples in most chemistry courses)
-
OR a 'negative catalyst' slowing down a
reaction (rarely mentioned at GCSE, sometimes at AS-A2 level, e.g. adding a
chemical that 'mops up' free radicals or other reactive species).
-
One other point, never say
anything remotely like 'the catalyst doesn't take part in the reaction',
it just speeds it up. It does take a crucial part in the reaction and may be
temporarily chemically changed, but the original catalyst does reform to
perform its task again i.e. speed up the reaction again!
-
Catalysts increase the rate of a reaction by helping break chemical bonds in reactant
molecules and provide a 'different pathway' for the reaction.
-
This effectively means the Activation Energy
Ea is reduced,
irrespective of whether its an exothermic or endothermic reaction (see diagrams below
for the smaller 'humps' of the lowered activations energy produced by using a
catalyst).
In both cases, the activation energy
for a catalysed reaction (top of the green
hump shown on the diagrams) is the minimum kinetic energy the
reactant molecules must have to break bonds and undergo a chemical
change.
The black higher activation energy curve
represents the non-catalysed reaction pathway.
-
Therefore at the same temperature,
using a catalyst,
more reactant molecules have enough kinetic energy to react compared to the uncatalysed situation
- this increases the frequency of fruitful collisions leading to product
formation - more successful collisions per unit time.
-
The purple arrow
indicates the greater activation energy for the uncatalysed
reaction.
-
The green hump and
green arrow indicates the lower activation energy of the
catalysed reaction - less particle energy required due to the action of the
catalyst.
-
The catalyst does NOT increase the energy
of the reactant molecules!
-
Neither does a catalyst increase the
frequency of reactant particle collisions.
-
BUT, the catalyst does increase the
frequency of 'fruitful' collisions because less kinetic energy is needed by
the particles to break bonds and undergo chemical change.
-
Many solution or gaseous catalysed
reactions involve a solid catalyst.
-
The reactant molecules are adsorbed
onto the surface, and this 'sticking on to the surface' enables the bonds of the
reactant molecules to be more easily broken.
-
This is actually what 'lowering of
the activation energy' means at the molecular level, how easy is it to break
bonds, so we are getting a bit technical here!
-
Although a true catalyst does take part in the
reaction and may change chemically temporarily, but it does not get used up and can be reused/regenerated with more reactants.
It does not change chemically or get used up in the end.
- Black manganese(IV) oxide (manganese
dioxide) catalyses the decomposition of hydrogen peroxide.
-
hydrogen peroxide ===> water + oxygen
-
2H2O2(aq)
===> 2H2O(l)
+ O2(g)
- Note that the catalyst is NOT included in the
equation.
- This is because overall the catalyst doesn't
change chemically.
-
The manganese dioxide is chemically the same at the end of
the reaction but it may change a little physically if its a solid.
-
In the hydrogen peroxide solution
decomposition by the solid black catalyst manganese dioxide, the solid can
be filtered off when reaction stops 'fizzing' i.e. all of the hydrogen
peroxide has reacted-decomposed.
-
After washing with water,
the catalyst can be
collected and added to fresh colourless hydrogen peroxide solution and the
oxygen production 'fizzing' is instantaneous! In other words the catalyst
hasn't changed chemically and is as effective as it was fresh from the
bottle!
-
Note: At the end of the experiment the
solution is sometimes stained brown from minute manganese dioxide
particles. The reaction is exothermic and the heat has probably
caused some disintegration of the catalyst into much finer particles
which appear to be (but not) dissolved. In other words the catalyst
has changed physically BUT NOT chemically.
-
You can try salts of transition
metals (notable for catalytic effects) e.g. copper sulfate, iron
sulfate or cobalt sulfate to see if they have any effect on the rate
of decomposition of hydrogen peroxide. Its best to use sulfate salts
rather than chlorides because the hydrogen peroxide may oxidise the
chloride ion and give rise to further unintentioned reactions.
-
You can investigate the rate of reaction
of hydrochloric acid with zinc and find that adding a few granules of copper
or a few drops of copper sulfate solution speeds of the reaction, which you
can measure the rate of reaction by following the rate of evolution of
hydrogen gas.
- Using an effective catalyst can reduce
costs in the chemical and food industries by increasing the rate of reaction
(more efficient) and lowering the energy requirements if the process can be done
at lower temperatures.
- Increasing the rate of reaction saves time and
operating at a lower temperature saves energy and therefore saves money.
- However, catalysts can be very specialised and
expensive to produce.
- They also get contaminated ('poisoned') and
become less efficient and might have to be extracted and cleaned up, but if a
true catalyst, this' refurbishment' should enable the catalyst to be reused
(theoretically catalysts take part in the reaction, but are not consumed in the
reaction).
- Sulfur compounds poison the iron catalyst used
in the Haber synthesis of ammonia.
- One way of minimising the
poisoning-contamination of catalysts is to purify the reactant molecules before
they enter the chemical reactor chamber.
- See 'chemical
economics' for other commercial aspects of chemical production.
-
Different reactions need different catalysts
and they are all extremely important in industry:
-
Catalysts make chemical
industrial processes much more efficient and economic e.g.
-
Nickel catalyses the hydrogenation of
unsaturated fats to margarine
-
iron catalyses the combination of
unreactive nitrogen and hydrogen to form ammonia in the Haber Synthesis.
-
Zeolite minerals catalyse the cracking of big
hydrocarbon molecules into smaller ones
-
Most polymer making reactions require a
catalyst surface or additive in contact with or mixed with the monomer molecules.
-
Catalysts in car exhausts
(catalytic converters) change harmful carbon monoxide and nitrogen monoxide into
harmless carbon dioxide and nitrogen, they also convert unburned potentially
carcinogenic hydrocarbons into carbon dioxide and water.
-
Enzymes are biochemical catalysts
are dealt with on other pages
Enzyme catalysed synthesis of a larger molecule from two smaller
ones.
Enzyme catalysed decomposition a larger molecules into two
smaller molecules.
For more advanced details on catalysis see
Advanced Level Chemistry Theory pages on "CHEMICAL
KINETICS"
APPENDIX - activation energy,
catalysts and reaction profiles
Reaction profile diagram |
Comments related to the reaction activation energy and use of
a catalyst |
 |
An exothermic
reaction with a small activation energy
The reaction may go very well without a
catalyst at a practical temperature, perhaps even at room
temperature |
 |
An exothermic
reaction with a moderately high activation energy.
This reaction might benefit from
using a catalyst if a suitable one is available |
 |
An endothermic
reaction with a big activation energy
This reaction would benefit from using a
catalyst e.g. to avoid using an excessively high temperature,
catalysts used to crack crude oil into useful fractions |
What next?
Associated Pages
Rates of
reaction notes INDEX
GCSE
Level (~US grade 8-10) School Chemistry Notes
(students age ~14-16)
Find your GCSE
science course for more help links to revision notes
ALL my Advanced Level pre-university
Chemistry Notes
Advanced A Level KINETICS
index
[SEARCH
BOX]
email doc
brown
GCSE level 'Rates of Reaction' multiple
choice quiz
Website content © Dr
Phil Brown 2000+. All copyrights reserved on Doc Brown's Chemistry revision notes, images,
quizzes, worksheets etc. Copying of website material is NOT
permitted. Exam revision summaries & references to science course specifications
are unofficial.
What next?
Associated Pages
|
|