Catalysts (Transition Metals as Catalysts)

Role of Transition Metals as Catalysts

Transition metals are widely used as catalysts because they can speed up chemical reactions without being consumed in the process. They do this by providing an alternative reaction pathway with a lower activation energy, which means the reaction can happen faster at the same temperature.

Key points:

• Transition metals increase the rate of reaction.
• They are not used up or permanently changed during the reaction.
• They provide an alternative pathway with lower activation energy.

This lowering of activation energy means more particles have enough energy to react when they collide, increasing the reaction rate.

For instance, in the presence of a transition metal catalyst, fewer collisions need to be very energetic for the reaction to proceed, so the reaction speeds up.

Examples of Transition Metal Catalysts

Transition metals are used in many important industrial and environmental processes as catalysts. Here are some key examples:

• Iron in the Haber Process: Iron is used as a catalyst to speed up the reaction between nitrogen and hydrogen to make ammonia. This process is essential for producing fertilisers.
• Nickel in Hydrogenation: Nickel catalyses the addition of hydrogen to unsaturated vegetable oils to make margarine. This process converts liquid oils into solid fats.
• Platinum in Catalytic Converters: Platinum helps convert harmful gases from car exhausts (carbon monoxide, nitrogen oxides, and unburnt hydrocarbons) into less harmful substances like carbon dioxide, nitrogen, and water.

These examples show how transition metals are vital in both industry and environmental protection.

For example, in catalytic converters, platinum provides a surface where harmful gases adsorb and react more easily, reducing pollution.

Surface Properties of Transition Metals

The ability of transition metals to act as catalysts is closely linked to their surface properties:

• Large surface area: Transition metals often have a large surface area when finely divided, providing more sites for reactions to occur.
• Adsorption of reactants: Reactant molecules stick (adsorb) onto the metal surface, which weakens their bonds and makes them more reactive.
• Active sites: Specific areas on the metal surface, called active sites, are where the chemical reactions happen.

The adsorption of reactants onto the catalyst surface helps break bonds and form new ones more easily, speeding up the reaction.

For example, in the hydrogenation of vegetable oils, hydrogen and the oil molecules adsorb onto the nickel surface, allowing hydrogen atoms to add across the double bonds of the oil.

Learning Example: Calculating Activation Energy Reduction

Suppose a reaction has an activation energy of 100 kJ/mol without a catalyst. When a transition metal catalyst is added, the activation energy drops to 60 kJ/mol. Calculate the percentage decrease in activation energy.

Percentage decrease = $\frac{\text{Original} - \text{New}}{\text{Original}} \times 100$

= $\frac{100 - 60}{100} \times 100 = 40\%$

So, the catalyst lowers the activation energy by 40%, making the reaction much faster.