carbon dioxide
hydrochloric
acid
calcium carbonate
(marble)
After completing this topic you should be able to :
• Describe the mechanism of heterogeneous catalysis
• Learn how a catalyst speeds up reaction rate by lowering
the activation energy, and how to represent this on a
potential energy diagram.
Catalysts at Work
Heterogeneous
When the catalyst and reactants are in different states
you have ‘Heterogeneous Catalysis’. They work by the adsorption of
reactant molecules.
E.g. Ostwald Process (Pt) for making nitric acid and the Haber
Process (Fe) for making ammonia and the Contact Process (Pt)
for making Sulphuric Acid.
Homogeneous
When the catalyst and reactants are in the same state
you have ‘Homogeneous Catalysis’. E.g. making ethanoic acid from
methanol and CO using a soluble iridium complex.
Enzymes are biological catalysts, and are protein molecules
that work by homogeneous catalysis. E.g. invertase and lactase.
Enzymes are used in many industrial processes
How a heterogenous catalyst works
Heterogenous Catalysis are thought to work in three stages...
Adsorption
Reaction
Desorption
Higher Chemistry Eric Alan and John Harris
How a heterogenous catalyst works
For an explanation of what happens click on the numbers in turn, starting with

How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of
the bonding electrons in the molecules thus weakening them and making a subsequent reaction
easier.
How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of the
bonding electrons in the molecules thus weakening them and making a subsequent reaction easier.
Reaction (STEPS 2 and 3)
Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur.
This increases the chances of favourable collisions taking place.
How a heterogenous catalyst works
Adsorption (STEP 1)
Incoming species lands on an active site and forms bonds with the catalyst. It may use some of
the bonding electrons in the molecules thus weakening them and making a subsequent reaction
easier.
Reaction (STEPS 2 and 3)
Adsorbed gases may be held on the surface in just the right orientation for a reaction to occur.
This increases the chances of favourable collisions taking place.
Desorption (STEP 4)
There is a re-arrangement of electrons and the products are then released from the active sites
Examples of heterogenous catalysts
Format
Metals
Ni, Pt
Fe
Rh, Pd
hydrogenation reactions
Haber Process
catalytic converters
Oxides
Al2O3
V2O5
dehydration reactions
Contact Process
FINELY DIVIDED
increases the surface area
provides more collision sites
IN A SUPPORT MEDIUM
maximises surface area and reduces costs
How a heterogenous catalyst works
In some cases the choice of catalyst can influence the products
Ethanol undergoes different reactions depending on the metal used as the catalyst.
The distance between active sites and their similarity with the length of bonds
determines the method of adsorption and affects which bonds are weakened.
Copper
Dehydrogenation (oxidation)
C2H5OH ——>
CH3CHO + H2
Alumina Dehydration
C2H5OH ——>
C2H4 + H2O
How a heterogenous catalyst works
Poisoning
Impurities in a reaction mixture can also adsorb onto the surface of a
catalyst thus removing potential sites for gas molecules and decreasing
efficiency.
expensive
because... the catalyst has to replaced
the process has to be shut down
examples
Sulphur
Lead
Haber process
catalytic converters in cars
Investigating Catalysis
The problem
Transition metals and their compounds are said
to be effective catalysts
You are asked to link the transition metal
compounds to their ability to act as a catalyst
for the decomposition of hydrogen peroxide
Materials available
hydrogen peroxide solution (5-10 vol)
cobalt(II) chloride (aq)
copper(II) sulphate (aq)
manganese(II) chloride (aq)
nickel(II) nitrate (aq)
bench dilute sodium hydroxide
(approximately 1 mol l-1).
Decomposition of hydrogen peroxide
Hydrogen Peroxide → water + oxygen
2H2O2
→ 2H2O + O2
What will you see during the reaction?
What is the test for oxygen?
Testing the catalysts
Collect 5 test tubes in a rack
To each add 5 cm3 of hydrogen peroxide
Add 3 cm3 of a different catalyst to 4 test
tubes, leaving the last one as a control
Note the results
Which type of catalysis?
Did the experiment involve heterogeneous or
homogeneous catalysis?
Which catalyst was the most effective?
Changing the solubility of the
catalysts
The transition metal ion in the compound has
the catalytic effect.
The ion can be found in solution (aqueous) or in a
solid state.
How could you use the chemicals provided to
create transition metal ions in the solid state
i.e. a heterogeneous catalyst?
Investigate the effect of changing
the state of the catalysts
Which was the best catalyst overall?
The students should find that none of the substances on
their own is an effective catalyst but that the solid
precipitates made by mixing any of the transition metal
solutions with sodium hydroxide solution are good
catalysts, ie this is heterogeneous catalysis. You may
have to join in the discussion to ensure that the students
realise this.
An Example of a homogenous catalyst
Dissolve 4 spatulas full of potassium sodium
tartrate in a small beaker with about 2.5 cm depth
of water. Repeat with another beaker (the
control)
To one beaker, add enough 10% cobalt chloride
solution (catalyst) for the solution to be pink.
Add 10 ml hydrogen peroxide to the beakers.
Heat the control to around 60 - 80 oC and note
observations
Now heat the one with the catalyst and note
observations
An Example of a homogenous catalyst
Higher Chemistry Eric Alan and John Harris
This is an impressive demonstration of how a catalyst is involved in the progress of a
reaction. Students can add another 10 cm3 of the hydrogen peroxide solution and if
there is any potassium sodium tartrate remaining they will see a similar reaction.
The reaction is an oxidation of the tartrate ion (proper name is 2, 3dihydroxybutandioate ion) to carbon dioxide gas and the methanoate ion. Hydrogen
peroxide oxidises the tartrate ion very slowly if there is no catalyst, even at elevated
temperatures.
Cobalt(ll) ions are pink. The hydrogen peroxide initially oxidises the cobalt(II), Co2+, to
cobalt(lll), Co3+, which is green. The cobalt(III) bonds to the tartrate ion, allowing the
oxidation to take place. The CO3+ is then reduced back to CO2+ and the pink colour
returns.
The cobalt catalyst provides an alternative route for the reaction to occur. This
alternative route has a lower activation energy and the reaction proceeds much more
quickly.
Catalysts and Potential energy diagrams
Catalysts work by providing…
“AN ALTERNATIVE REACTION PATHWAY WHICH HAS A LOWER ACTIVATION
ENERGY”
Visualising activated complex
Potential energy graphs and catalysts
Uncatalysed
75 -
60 50 -
Reactants
Catalysed
P.E.
Products
25 -
Reaction path
Potential energy graphs and catalysts
Activation energy Ea for the forward uncatalysed reaction
75 -
60 -
Activation energy Eafor the
forward catalysed reaction
Reactants
50 -
P.E.
Products
25 -
Reaction path
Catalysts lower the activation energy needed for a successful collision.
Potential energy graphs and catalysts
Activation energy Ea for the reverse uncatalysed reaction reaction
75 -
Activation energy Ea for the
reverse catalysed reaction
60 50 -
Reactants
P.E.
Products
25 -
Reaction path
Catalysts lower the activation energy needed for a successful collision.
Potential energy graphs and catalysts
75 -
Catalysts lower the activation
energy needed for a successful
collision.
60 50 -
Reactants
P.E.
Products
25 -
Reaction path
∆H
Activation energy
Effect of catalyst – forward reaction
No change
Lowered
Effect of catalyst – reverse reaction
No change
Lowered
Catalysts
A catalyst speeds up the reaction by lowering the
activation energy.
A catalyst does not effect the enthalpy change for a
reaction
A catalyst speeds up the reaction in both directions
and therefore does not alter the position of
equilibrium or the yield of product, but does decrease
the time taken to reach equilibrium.
Energy distribution and catalysts
Ea
No of
Collisions
with a
given
K.E.
Un-catalysed reaction
Kinetic energy
Ea
Total number of collisions (area under
the graph) with sufficient K.E.
energy to create new products.
Catalysed reaction
Ea is reduced