inorganic - transition metals 5 Flashcards
catalysts
catalysts are substances that affect rate of reaction by providing an alternative route with a lower activation energy, without being used up or chemically altered
there are two types:
- heterogenous catalysts
- homogenous catalysts
advantages of using catalysts in industries
- allows reactions to proceed at lower temperatures and pressures; saving valuable resources and reducing costs
heterogenous catalysts
- a heterogeneous catalyst is in a different phase (solid, liquid or gas) from the reactants
- they are usually present as solids while the reactants are either liquid or gas
- their catalytic reaction occurs at active sites on solid surface
how to make heterogenous catalysts more efficient
increase their surface area; the larger the surface area, the greater the efficiency - this can be done by finely dividing them
heterogenous catalyst - the Haber process
- ammonia is made by the reaction of nitrogen and hydrogen
- the catalyst is IRON - present in small pea sized lumps to increase the surface area
iron catalyst
N2(g) + 3H2(g) ⇋ 2NH3(g)
heterogenous catalyst - the Contact process
- sulphuric acid produced
- reactions catalysed by Vanadium (V) oxide V2O5
1) V2O5 oxidises sulphur dioxide to sulphur trioxide and then is reduced to V2O4
SO2 + V2O5 –> SO3 + V2O4
2) V2O4 is then re-oxidised by oxygen back to V2O5
2V2O4 + O2 –> 2V2O5
homogenous catalysts
- a homogeneous catalyst is in the same phase as the reactants, the reaction proceeds through an intermediate species
homogenous catalyst - Fe2+ with iodide and peroxidesulfate ions
- peroxidesulfate ions S2O8^2- will oxidise iodide ions to iodine
- this reaction is catalysed by Fe2+
1) peroxidesulfate ions oxidises iron(ii) to
iron (iii)
S2O8^2-(aq) + 2Fe2+(aq) –> 2SO4^2- (aq) + 2Fe3+ (aq)
2) Fe3+ oxidises iodide to iodine, regenerating Fe2+ ions
2Fe3+ (aq) + 2I-(aq) –> 2Fe2+(aq) + I2(aq)
overall reaction:
S2O8^2-(aq) + 2I- –> SO4^2-(aq) + I2(aq)
implications of heterogenous catalysts
- they may get poisoned by impurities which block the active site
- and consequently reduce efficiency
- buying a new catalyst to replace it may be costly
why do transition metals make good catalysts
- they can exist in variable oxidation states, therefore they can easily provide alternative routes
- they have partly filled d-orbitals which can be used to form weak chemical bonds with the reactants - this has two effects:
1) weakening bonds within the reactants
2) holding reactants close together on the solid surface in the correct orientation for reaction
autocatalysts
- the reaction is catalysed by one of the reaction’s products
- the reaction starts out slowly, but as the concentration of the product that catalyses the reaction increases, the reaction speeds up
- it then behaves like a normal reaction, gradually slowing down as reactants are getting used up
C2O42- and MnO4- catalysis
- both of the species are negatively charged so they repel each other and are unlikely to collide. this makes the reaction very slow
- the reaction produces Mn2+ ions which catalyse the reaction
- as a result the rate of reaction increases over time as more product, and therefore catalyst, is produced.
equations involved
- Mn2+ ions reduce MnO4- to Mn3+
- Mn3+ is an intermediate species
1) 4Mn2+(aq) + MnO4-(aq) +8H+(aq)
- -> 5Mn3+(aq) + 8H2O(l) - Mn3+ ions then oxidise the ethandioate ions
2) 2Mn3+(aq) + C2O4-(aq)
- -> 2Mn2+(aq) + 2CO2(g) - Mn2+ reformed
overall reaction
2MNO4-(aq) + 16H+(aq) + 5C2O4 2- (aq)
- -> 2Mn2+(aq) + 8H2O(l) + 10CO2 (g)
- Mn2+ and Mn3+ are excluded from the overall reaction since they exist on both sides
