Catalyst

• A catalyst is a substances that speeds up reactions by providing an alternative route with a lower activation energy
• Transition are good catalyst because:

  • They show variable oxidation states. This allows them to act as intermediates in exchange of electrons between reacting species
  • They provide a surface for reactions to occur. The metal forms weak bonds to the reacting species, holding them in place
• There are two types of catalyst:

  • Homogeneous catalyst
  • Heterogenous catalyst
• Homogeneous catalyst are in the same phase as the reaction species
• Heterogenous catalyst are in a different phase to the reaction species and the reaction occurs at the active site of the surface
• Heterogenous catalyst act as a surface for a reaction to occur on, providing a reaction route with a lower activation energy

  • As the reactants adsorb to the surface of the catalyst (weakly bonds) at an active site
  • Then a reaction takes place on the surface
  • Lastly, products desorb from the surface
• Catalyst are very expensive so to maximise the efficiency of the catalyst minimises the cost. A method to do this is to increase surface area of catalyst by using a support medium e.g. platinum on a catalytic converter
• Impurities can bind onto the active site of the catalyst and stays on there (poisons it) which would decrease the efficiency of the catalyst so it would have to be changed eventually
• Haber process

  • Used to make ammonia (NH3)
  • Equation is N2 (g) + 3H2 (g) ⇌ 2NH3 (g)
  • Requires an iron catalyst (heterogenous)
  • Impurities can poison the catalyst
• Contact process

  • Used to produce sulfuric acid
  • Requires a Vanadium (V) oxide which is heterogenous
  • First step in equation is SO2 + V2O5 --> SO3 +V2O4
  • Second equation is 2V2O4 +O2 ---> 2V2O5
  • Overall equation is 2SO2 + O2 ⇌ 2SO3
  • Then the SO3 reacts with water to form sulfuric acid
• Manufacture of methanol

  Stage 1 (uncatalysed)

  • CH4 + H2O ---> CO + 3H2

  Stage 2 (with chromium || oxide which is a heterogenous catalyst)

  • CO + 2H2 --> CH3OH
• Homogenous catalyst

  This catalyst often involves a change in oxidation state of transition metal ions.

  • The transition metal ion forms an intermediate, then a further reaction occurs to regenerate the original transition metal ion
  • So there would be two humps on the reaction profile diagram
• Reaction of I2 and S2O8 2-

  • The uncatalysed reaction between iodine and peroxodisulfate ion is very slow because the reactants are both negatively charged so they repel from each other
  • Equation : S2O8 2- + 2I-(aq) --> I2 + 2SO4 2-
  • Adding aqueous Fe2+ ions (catalyst) provides an alternative reaction pathway which is much faster
  • Equations: S2O8 2- + 2Fe2+ --> 2SO4 2- + 2Fe3+
  • Which reacts with I- ions forming the equation 2Fe3+ + 2I- --> 2Fe2+ + I2
• Autocatalysis is when one of the products of a reaction acts as a catalyst for the reaction
• Initially the rate of an autocatalysed reaction is very slow, but as the product increases the reaction rate increases
• Ethanedioic acid is oxidised by acidified potassium magnate (V||) ions with the Mn2+ ions acting as a catalyst.
• Ethandioic acid and magnate ions

  First the magnate ions react with Ethanedioic acid.

  Equation:

  • 2MnO4 - + 5C2O4 2- + 16H+ --> 2Mn 2+ + 10CO2 + 8H2O

  Second the Mn2+ formed would provide an alternative faster reaction pathway

  Equations:

  • 4Mn 2+ + MnO4 - + 8H+ --> 5Mn 3+ + 4H2O
  • 2Mn 3+ +C2O4 2- --> 2CO2 + 2Mn 2+
  • Both of these reaction can happen either way so both Mn2+ and Mn3+ act as catalyst