Transition metals: Homogenous catalysts

Edexcel A level Chemistry (2017)

Topic 15: Transition metal chemistry

Learning Objectives related to homogenous catalysis

15/29. To know that transition metals and their compounds can act as heterogeneous and homogeneous catalysts.

15/33. To know that a homogeneous catalyst is in the same phase as the reactants and appreciate that the catalysed reaction will proceed via an intermediate species.

15/34. To understand the role of Fe2+ ions in catalysing the reaction between I− and S2O82− ions.

Transition metals as catalysts

Catalysts are substances that change the rate of chemical reaction and remain unchanged at the end of that reaction. 

They act so as to reduce the activation energy of a chemical change.

The reduction of the activation energy of a reaction means that more molecules within that reaction possess the activation energy.

A greater percentage of collisions between these molecules can be effective if the activation energy is lower.  It is this increase in the number of effective collisions that leads to an increase in the reaction rate.   

With transition metals, catalysts can be either homogenous or heterogeneous.

Homogenous transition metal catalysts

These homogenous catalysts are in the same state as the reactants and products in the chemical change.

A classic example of a homogenous catalysts that has been thoroughly studied is iron(III) ions in the reaction between iodide ions and peroxodisulphate ions. 

The reaction between iodide ions (2I^—_(aq)) and peroxodisulphate ions (S₂O₈^2—_(aq)) happens very slowly because the activation energy is very high.  We can see why when we look at the equation for the reaction:

2I^—_(aq)   +    S₂O₈^2—_(aq) ⟶  I_{2 (aq)}  +   2SO₄^2—_(aq)

The reaction is between two negatively charged ions and these ions, in the ordinary course of events, would not be attracted to each other but would repel each other. 

Therefore, these two particles would need to have high activation energies for them to make an effective collision.

How is the activation energy reduced?

The reaction speeds up considerably if iron(III) ions are added to the reaction mixture.  Iron(III) ions act as a catalyst.  Why?

Let’s look at the electrode potentials of the half-cell reactions for the three species involved in the catalysed reaction.

Here they are:

I_{2 (aq)}       +   2e^—   ⟶   2I^— _(aq)      E^o  + 0.54v

Fe³⁺_(aq)   +   e^—     ⟶   Fe²⁺_(aq)     E^o  + 0.77v

S₂O₈^2—_(aq) + 2e^—  ⟶   2SO₄^2—_(aq)   E^o  + 2.01v

The principle we work is that top right species will reduce bottom left species.

From these values, we can see that iodide ions will reduce iron(III) ions to iron(II) ions.  Then the iron(II) ions formed will reduce the peroxodisulphate ions to sulphate ions.

Both these reactions involve the combination of ions of opposite charge and therefore the reactions require lower activation energies. 

The iron(III) ions are reformed in the second reaction and can therefore engage in further reduction/oxidation action with other iodide and peroxodisulphate ions. 

The reaction overall is faster than without the iron(III) ions. 

Iron(III) ions engage in the reaction as a catalyst but remain unchanged at the end of the redox process. They are the intermediate species. 

All ions are in the same aqueous state so the iron(III) ion catalyst is a homogenous catalyst. 

We have seen then

15/33. that a homogeneous catalyst is in the same phase as the reactants and appreciated that the catalysed reaction will proceed via an intermediate species.