23.5 Transition metal catalysts

A catalyst is something that speeds up a chemical reaction by providing an alternative reaction pathway with a lower activation energy, and it doesn’t get used up

They can change oxidation states easily by gaining or losing electrons in their d orbitals, so they can transfer electrons to speed up reactions

It is a catalyst that is in a different phase to the reactants, eg is in a different physical state The contact process with V2O5 The Haber process with Fe

. Increase their surface area: The larger the surface area, the better the efficiency . Spread catalyst onto support medium so increase surface to mass ratio so a little goes a long way. Do that with expensive catalyst

It is used industrially to make sulfuric acid Vanadium (V) oxide V2O5 catalyses this reaction

SO2 is made into SO3 You can react SO3 with water (H2O) to make H2SO4

. Vanadium (v) oxide is the catalyst used, and it occurs in two steps: • First the V2O5 oxidises SO2 to SO3 in the reaction, and itself is reduced : V2O5 + SO2 —> SO3 + V2O4 This reduced catalyst is then oxidised back to V2O5 by oxygen to start the reaction again V2O4 + 1/2 O2 —> V2O5

It is a heterogeneous catalyst because it is in a different phase to the reactants eg a different physical state Because it has the right variable oxidation state needed

. They generally aren’t used to catalyse redox reactions because they don’t have an incomplete 3d sub shell, and don’t have variable oxidation states

SO2(g) + 1/2O2(g) ———-> SO3(g) With V2O5 catalyst used, over the arrow

Ammonia is made from reacting nitrogen and hydrogen

. Nitrogen and Hydrogen are reacted together . The iron catalyst Fe(s) is used in the shape as lots of pea sized lumps to. increase surface area Reaction: N2(g) + 3H2(g) —> 2NH3(g) With iron over the arrow

The catalyst in both is a solid whilst the reactants are gases So the reaction occurs on the surface of the catalyst So increasing surface area of catalyst increases the number of molecules that can react at the same time, increasing the rate of the reaction

Rhodium is the catalyst in them They clean up emissions from car engines, by converting nasty waste gases into less harmful products . They are shaped as a ceramic lattice coated with a thin layer of rhodium on top to increase surface area of catalyst to minimise cost of it because only a small coat is needed

During a reaction, reactants are adsorbed onto the surface of heterogeneous catalysts Impurities in the mixture may also bind to the catalyst surface, blocking the reactants from being absorbed So it reduces the surface area available for the reactants so slows down the reaction It can increase cost as less product is made with the certain amount of energy

There are sulfur impurities because the hydrogen used is from methane which is obtained from natural gas So that sulfur can be adsorbed into the iron forming iron sulfide

. These are catalysts in the same physical state as the reactants. Eg an aqueous catalyst for a reaction between two aqueous solutions

It forms an intermediate species by the reactants combining with the catalyst This intermediate species then reacts to form the products and reform the catalyst So these two reactions make the enthalpy profile for a homogeneously catalysed reaction have two humps in it The energy needed to form the intermediates (and to form the products from the intermediates) is less than the energy needed to create the products directly from the reactants

. Fe2+ catalysing the reaction between S2O8 2- ions and I- ions . Mn2+ autocatalysing the reaction between MnO4- and C2O4 2-

The redox reaction is catalysed by iron (Fe2+) The overall reaction is shown S2O8 2- + 2I- —> I2 + 2SO42-

Because the reactants are both negatively charged so will repel each other so it is unlikely they will collide and react

Because it is positively charged so the reaction will now involve a positive and negative ion First the Fe2+ ions are oxidised to Fe3+ ions by the S2O8 2- ions: S2O8 2-(aq) + 2Fe2+ —> 2Fe3+ + 2SO4 2- The newly formed intermediate Fe3+ ions can now easily oxidise the I- ions to iodine, and the catalyst is regenerated 2Fe3+ + 2I- —->I2 + 2Fe2+

Mn2+ is the catalyst in this reaction, and it is an autocatalysis reaction because Mn2+ is a product and also acts as a catalyst This means that as the reaction progresses, more product is formed so it speeds up over time

2MnO4- + 5C2O42- + 16H+ —–> 2Mn2+ + 8H2O + 10CO2 The activation energy at the beginning is very high as there is no catalyst There are two negative ions as reactants which repel eachother