Catalysts

• A catalyst is a substance that increases the rate of reaction without being used up in the reaction. It works by providing an alternative reaction path that requires a lower activation energy for the reaction to occur.
• When using a catalyst, the Maxwell-Boltzmann distribution curve is unchanged in shape, however the position of the activation energy is shifted to the left, so a greater proportion if molecules have sufficient energy to react.
• Catalysts are used in industry because they lower the energy costs of the reaction process. They allow lower temperatures and pressures to be used, whilst still achieving the same rate of reaction. They can also give a higher atom economy.
• During catalysed reactions, an intermediate will be formed. Therefore, there is a dip in the reaction profile of a catalysed reaction, since the intermediate is less stable (has higher energy) than the reactants and products.
• Heterogenous catalysts are catalysts that are in a different phase/state to the species in the reaction.
• An example of a heterogenous catalyst is the solid iron catalyst used in the Haber process, used to speed up the reaction between nitrogen and hydrogen gases.
• Transition metals make good catalysts because they have variable oxidation states. Electrons are transferred to produce a reactive intermediate and speed up reaction rate.
• An example of a transition metal being used as an effective catalyst is vanadium oxide in the Contact Process, helping to speed up the conversion of sulfur dioxide to sulfur trioxide.
• Solid catalysts work by adsorbing molecules onto an active site on the surface of the catalyst. These active sites increase the proximity of molecules and weaken the covalent bonds in the molecules so that reactions occur more easily and the rate is increased. These catalysts are used in industry to give a surface for the reaction to occur on.