10 Catalysis

Transition metals and their compounds can act as heterogeneous and homogeneous catalysts Catalysis A heterogeneous catalyst is in a different phase from the reactants A homogeneous catalyst is in the same phase as the reactants Surface area: Increasing the surface area of a solid catalyst will improve its effectiveness. A support medium is often used to maximise the surface area and minimise the cost (e.g. Rh on a ceramic support in catalytic converters). Heterogeneous catalysts are usually solids whereas the reactants are gaseous or in solution. The reaction occurs at the surface of the catalyst. Heterogeneous catalysis Adsorption of reactants at active sites on the surface may lead to catalytic action. The active site is the place where the reactants adsorb on to the surface of the catalyst. This can result in the bonds within the reactant molecules becoming weaker, or the molecules being held in a more reactive configuration. There will also be a higher concentration of reactants at the solid surface so leading to a higher collision frequency. Transition Metals can use the 3d and 4s e- of atoms on the metal surface to form weak bonds to the reactants. Steps in Heterogeneous Catalysis 1. Reactants form bonds with atoms at active sites on the surface of the catalyst (adsorbed onto the surface) 2. As a result bonds in the reactants are weakened and break 3. New bonds form between the reactants held close together on catalyst surface. 4. This in turn weakens bonds between product and catalyst and product leaves (desorbs). Strength of adsorption The strength of adsorption helps to determine the effectiveness of the catalytic activity. Some metals e.g. W have too strong adsorption and so the products cannot be released. Some metals e.g. Ag have too weak adsorption, and the reactants do not adsorb in high enough concentration. Ni and Pt have about the right strength and are most useful as catalysts. Catalysts increase reaction rates without getting used up. They do this by providing an alternative route with a lower activation energy. V2O5 is used as a catalyst in the Contact Process. Overall equation : 2SO2 + O2  2SO3 step 1 SO2 + V2O5  SO3 + V2O4 step 2 2V2O4 + O2  2V2O5 Examples of heterogeneous catalysts Learn the equations for this mechanism. Note the oxidation number of the vanadium changes and then changes back. It is still classed as a catalyst as it returns to its original form. Cr2O3 catalyst is used in the manufacture of methanol from carbon monoxide and hydrogen. CO + 2H2  CH3OH Fe is used as a catalyst in the Haber Process N2 + 3H2  2NH3 Poisoning Catalysts Catalysts can become poisoned by impurities and consequently have reduced efficiency. Poisoning has a cost implication e.g. poisoning by sulphur in the Haber Process and by lead in catalytic converters in cars means that catalysts lose their efficiency and may need to be replaced. Leaded petrol cannot be used in cars fitted with a catalytic converter since lead strongly adsorbs onto the surface of the catalyst. It is important to ensure the purity of the reactants if poisoning can occur. When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species. Homogeneous catalysis The intermediate will often have a different oxidation state to the original transition metal. At the end of the reaction the original oxidation state will reoccur. This illustrates importance of variable oxidation states of transition metals in catalysis. N Goalby chemrevise.org Examples of homogeneous catalysts The reaction between I- and S2O8 2- catalysed by Fe2+ overall S2O8 2- + 2I-  2SO4 2- + I2 Catalysed alternative route stage 1 S2O8 2- + 2Fe2+  2SO4 2- + 2Fe3+ stage2 2I- + 2Fe3+  2Fe2+ + I2 The uncatalysed reaction is very slow because the reaction needs a collision between two negative ions. Repulsion between the ions is going to hinder this – meaning high activation energy. Learn these 2 examples and equations carefully Both of the individual stages in the catalysed mechanism involve collision between positive and negative ions and will have lower activation energies. The autocatalysis by Mn2+ in titrations of C2O4 2- with MnO4 – overall 2 MnO4 – + 5 C2O4 2- + 16 H+  2Mn2+ + 10 CO2 + 8 H2O Catalysed alternative route Step 1 4Mn2+ + MnO4 – + 8 H+  5Mn3+ + 4 H2O Step 2 2Mn3+ + C2O4 2-  2Mn2+ + 2 CO2 The initial uncatalysed reaction is slow because the reaction is a collision between two negative ions which repel each other leading to a high activation energy. The Mn2+ ions produced act as an autocatalyst and therefore the reaction starts to speed up because they bring about the alternative reaction route with lower activation energy. The reaction eventually slows as the MnO4 – concentration drops. This is an example of autocatalysis where one of the products of the reaction can catalyse the reaction. Mixture with only reactants showing effect of autocatalysis Mixture with added catalyst Mn2+ at start Following the reaction rate This can be done by removing samples at set times and titrating to work out the concentration of MnO4 – . It could also be done by use of a spectrometer measuring the intensity of the purple colour. This method has the advantage that it does not disrupt the reaction mixture, using up the reactants and it leads to a much quicker determination of concentration. Autocatalytic Reaction between Ethanedioate and Manganate ions Reaction between iodide and persulphate ions Fe3+ ions can also act as the catalyst because the two steps in the catalysed mechanism can occur in any order. 9 Constructing a catalysed mechanism for a reaction Example The following reaction is catalysed by Co2+ ions in an acidic solution. SO3 2– + ½ O2  SO4 2– . Write a mechanism for the catalysed reaction by writing two equations involving Co 2+ and Co3+ ions. Split the full equation into its two half equations SO3 2– + ½ O2  SO4 2– SO3 2– + H2O  SO4 2– + 2H+ + 2e- ½ O2 + 2H+ + 2e-  H2O Add in cobalt to make two new redox equations. Making sure the oxidised cobalt equation is combined with the original reduced half equation and vice versa Co2+  Co3+ + e- Co3+ + e-  Co2+ ½ O2 + 2H+ + 2Co2+  H2O + 2Co3+ 2Co3+ + SO3 2– + H2O  SO4 2– + 2H+ + 2Co2+ Check your two mechanism equations add up to the original full non-catalysed equation.