Transition Metal Catalysts

Catalysis is one of the most useful properties of the transition metals. A catalyst is a substance that speeds up a chemical reaction without being used up itself, by providing an alternative reaction pathway with a lower activation energy. Transition metals and their compounds are especially good at this, and you will meet them again and again across industrial chemistry and A Level reaction mechanisms.

Two properties of transition metals explain their catalytic behaviour, both of which come from their electron structure (covered in more detail in Introduction to transition metals and the d-block).

• Variable oxidation states. Transition metals can gain and lose electrons to switch between several stable oxidation states. This lets a metal ion accept electrons from one reactant and then hand them to another, cycling back to its original oxidation state at the end.
• Available d orbitals. Partially filled d orbitals let transition metal atoms and ions form weak, temporary bonds with reactant molecules, either by adsorbing them onto a solid surface or by forming complex ions in solution (see Ligands and complex ions). This holds the reactants in a favourable position and weakens their bonds, lowering the activation energy of the reaction.

Because a catalyst is chemically unchanged at the end of the reaction, it does not appear in the overall balanced equation, even though it is essential to the mechanism.

A heterogeneous catalyst is in a different physical state from the reactants, most commonly a solid catalyst working on gaseous or liquid reactants. Reaction happens at active sites on the catalyst's surface, so the process is often described in three stages.

Classic examples include:

• Iron in the Haber process, catalysing the reaction between nitrogen and hydrogen to make ammonia.
• Vanadium(V) oxide in the Contact process, converting sulfur dioxide to sulfur trioxide (related vanadium chemistry is explored in Reactions of vanadium and other transition metals).
• Platinum, palladium and rhodium in catalytic converters, converting toxic carbon monoxide and unburned hydrocarbons into carbon dioxide and water, and reducing nitrogen oxides back to nitrogen.
• Nickel in the hydrogenation of alkenes, used to harden vegetable oils into margarine.

A key practical issue with heterogeneous catalysts is catalyst poisoning: impurities such as sulfur compounds can bind permanently to active sites, blocking them and reducing the catalyst's effectiveness over time.

A homogeneous catalyst is in the same phase as the reactants, usually all dissolved in aqueous solution. Instead of adsorbing molecules onto a surface, the metal ion reacts with one of the reactants to form an intermediate species, briefly changing its own oxidation state, before reacting again to regenerate the original catalyst and release the product.

A well known example is the iron-catalysed reaction between iodide ions and peroxodisulfate(VI) ions. This reaction is slow when uncatalysed, because both ions are negatively charged and repel each other. Iron ions get around this by reacting with one anion at a time:

Step 1: \( 2Fe^{3+} + 2I^- \)→\( 2Fe^{2+} + I_2 \)

Step 2: \( 2Fe^{2+} + S_2O_8^{2-} \)→\( 2Fe^{3+} + 2SO_4^{2-} \)

Adding the two steps together and cancelling the iron species gives the overall equation:

\( 2I^- + S_2O_8^{2-} \)→\( I_2 + 2SO_4^{2-} \)

Notice that iron can act as the catalyst here whether it starts as \(Fe^{2+}\) or \(Fe^{3+}\), because the cycle just runs through the two steps in a different order. Many homogeneous catalytic cycles like this involve the metal ion forming a temporary complex with a ligand along the way, which is why it is worth also reviewing Ligand exchange reactions.

Sometimes one of the products of a reaction is itself a catalyst for that same reaction, a phenomenon called autocatalysis. A classic example is the reaction between manganate(VII) ions and ethanedioate (oxalate) ions:

\( 2MnO_4^- + 5C_2O_4^{2-} + 16H^+ \)→\( 2Mn^{2+} + 10CO_2 + 8H_2O \)

At the start, the reaction is slow because the two reacting ions are both negatively charged. As \(Mn^{2+}\) ions build up, however, they begin to catalyse the reaction, offering a lower activation energy pathway similar to the iron cycle above. The rate therefore increases even as reactant concentration falls, giving a characteristic slow start followed by a speeding up, before the rate finally drops again as the reactants run out.

• Phase: heterogeneous catalysts are in a different phase to the reactants; homogeneous catalysts are in the same phase.
• Mechanism: heterogeneous catalysis relies on surface adsorption at active sites; homogeneous catalysis relies on the metal ion temporarily changing oxidation state or forming an intermediate complex.
• Separation: heterogeneous catalysts are easy to filter out of a reaction mixture afterwards; homogeneous catalysts are dissolved in the mixture and much harder to recover.

Understanding both types gives you a complete picture of how transition metals control reaction rates, from industrial gas-phase reactions to reactions in aqueous solution.

Related lessons

• Uk A Level Chemistry