Transition metals as catalysts

In a nutshell

Transition metal elements and compounds make good catalysts as they can change oxidation number. Heterogenous catalysts are in a different state to the reactants. Homogenous catalysts are in the same state as the reactants. Impurities reduce the efficiency of catalysts and increase costs.

Oxidation numbers

Transition metal elements and compounds can their change oxidation number by gaining or losing d-orbital electrons. This ability to transfer electrons makes them good catalysts. 

Example

The Contact Process is used to make sulfuric acid and uses vanadium ($$V$$​) oxide as a catalyst. Vanadium ($$V$$​) oxide is able oxidise $$SO_2$$ to $$SO_3$$ because it can be reduced to vanadium ($$IV$$​). 

​$$vanadium \ (V) \rightarrow vanadium \ (IV) \\ V_2O_5 \space + \space SO_2 \rightarrow V_2O_4 \space + \space SO_3$$​​

Vanadium ($$IV$$) oxide is oxidised back into vanadium ($$V$$​) oxide using oxygen and is used again. 

​$$vanadium \ (IV) \rightarrow vanadium \ (V) \\ V_2O_4 \space + \space \frac{1}{2}O_2 \rightarrow V_2O_5$$​​

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Heterogenous catalysts 

A heterogenous catalyst is a catalyst which is in a different phase to the reactants. At the surface of heterogenous catalysts are active sites; reactants are absorbed onto the active site and the reaction takes place. Increasing the surface area of the catalyst increases the rate of reaction, as more molecules can react at the same time.

The area of a catalyst is often maximised by using support mediums. Support mediums are coated with a small amount of catalyst which reduces the cost of a reaction significantly.

Example

Vanadium ($$V$$​) oxide is a heterogenous catalyst used in the Contact Process.

Example

Iron is a heterogenous catalyst used in the Haber Process for making ammonia. The gas reactants are passed over the solid iron catalyst.

​$$N_2(g) \space + \space 3H_2(g) \space \space \underrightarrow {^{Fe(s)} } \space \space 2NH_3(g)$$​​

Catalytic converters 

Catalytic converters in cars use a platinum or rhodium catalyst to convert nitrogen monoxide and carbon monoxide to nitrogen and carbon dioxide.

​$$2NO(g) \space + \space 2CO(g) \rightarrow N_2(g) \space + \space 2CO_2(g)$$

​Homogenous catalysts​

A homogenous catalyst is a catalyst in the same phase as the reactants. Usually homogenous catalysts catalyse the reaction between two aqueous solutions.

Homogenous catalysts bind with reactants to form an intermediate species. When the intermediate species reacts, the products form and the catalyst is reformed and reused.

The diagram below shows the enthalpy profile for a reaction involving a homogenous catalyst. The two peaks correspond to the two steps involved in the reaction. The activation energy needed to form the intermediate, and therefore the products, is lower when a catalyst is used. 

1. Reactants 2. Intermediates formed 3. Uncatalysed reaction 4. Products

Iodide and peroxodisulfate ions

Iodide ions and peroxodisulfate ions undergo a very slow redox reaction because both ions are negatively charged and repel each other.

When $$Fe^{2+}$$ ions are added, the rate of reaction increases significantly. This is because there is no repulsion as each step of the reaction involves a positive and negative ion. 

First, the $$S_2O_8 \ ^{2-}$$​ ions oxidise the $$Fe^{2+}$$ ions to $$Fe^{3+}$$ ions.

​$$S_2O_8 \ ^{2-}(aq) \space + \space 2Fe^{2+}(aq) \rightarrow 2Fe^{3+}(aq) \space + \space 2SO_4 \ ^{2-}(aq)$$​​

Then the $$Fe^{3+}$$ ions are intermediates and can easily oxidise $$I^-$$ ions to iodine. The catalyst is regenerated. This is an example of homogenous catalysis.

​$$2Fe \ ^{3+}(aq) \space + \space 2I^{-}(aq) \rightarrow I_2(aq) \space + \space 2Fe \ ^{2+}(aq)$$​​

Starch solution can be added to a solution to test if iodine is present. The solution will turn blue-black if iodine is present.

Autocatalysis 

In an autocatalysis reaction the product formed catalyses the reaction.

Example

When  and $$C_2O_4 \^{2-}$$$$MnO_4 \ ^{2-}$$ ions react, the rate of the uncatalysed reaction is very slow as both reactants are negatively charged and repel each other.

$$2MnO_4 \ ^- (aq) \space + \space 16H^+(aq) \space + \space 5C_2O_4 \ ^{2-} (aq) \rightarrow 2Mn^{2+}(aq) \space + \space 8H_2O(l) \space + \space10CO_2(g)$$​​

The $$Mn^{2+}$$​ ions formed can catalyse the reaction. The $$Mn^{2+}$$ ions react with the $$MnO_4 \ ^-$$ ions to form $$Mn^{3+}$$ ions.

​$$MnO_4 \ ^- (aq) \space + \space 4Mn^{2+}(aq)+ \space 8H^+(aq) \rightarrow 5Mn^{3+}(aq) \space + \space 4H_2O(l)$$​​

The $$Mn^{3+}$$ ions can then react with the $$C_2O_4 \ ^{2+}$$ ions to produce carbon dioxide. The $$Mn^{2+}$$ catalyst ions reform.

​$$2Mn^{3+} \space (aq) \space + \space C_2O_4 \space^{2-} (aq) \rightarrow 2Mn^{2+}(aq) \space + \space 2CO_2(g)$$​​

Impurities 

Impurities in reaction mixtures can bind to the active sites of catalysts and block reactants from binding. This is referred to as catalyst poisoning.

Catalyst poisoning reduces the surface area the reactants can bind to and therefore reduces the rate of reaction. Catalyst poisoning increases the cost of production significantly as more energy is required to produce a given quantity of product and sometimes the catalyst may need to be replaced. 

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Example 

Methane is used to produce hydrogen for the Haber Process. Methane is extracted from natural gas. Natural gas contains sulfur compounds as well as other impurities. When sulfur is adsorbed onto the iron catalyst during the Haber Process, iron sulfide forms which reduces the efficiency of the catalyst.