2 Starter: What do you know about catalysts? Substance that changes the speed/rate of a chemical reaction, but is not used up in the reaction.Speed up reactions by providing an alternative route with a lower activation energy(Do not alter ∆G, ∆H or ∆S for a reaction, only the activation energy).Speed the rate at which an equilibrium is reached by speeding up the forward and reverse reactionThey do not alter the position of the equilibrium (or the value of Kc –the equilibrium constant).Catalysts that slow down reactions are called inhibitors (or negative catalysts).
3 IntroductionCatalyst = substance that speeds up reaction without being used up.They provide an alternative mechanism with lower activation energy.For an equilibrium, they speed up both reactions equally and so do not effect equilibrium position.Heterogeneous catalyst is in different phase to the reactants.Homogeneous catalyst is in the same phase as the reactants.
5 Common Features of Catalysts They can be very specific and only catalyse a certain reaction or class of similar reactionsSize and surface area: Often only needed in very small amounts (e.g., 2 g of Pt can catalyse the decomposition of 1000m3 of H2O2).Solid catalysts more effective when they are a powder or as a thin layer – their surface area to volume ratio is very large.
6 Heterogeneous Catalysts Heterogeneous catalyst is in different phase to the reactants.Usually a solid catalyst and gaseous reactants.Most industrial catalysts are like this(e.g. Haber Process – Fe; Contact Process – V2O5).
7 Heterogeneous Catalysts 1) Reactants adsorbed onto surface (onto active sites).weakens bondsbrings molecules closermore favourable orientation2) Reaction takes place.The reaction happens at the surface of the catalystThe places where the reactions happen are called “active sites”e.g., in the Haber Process, nitrogen and hydrogen react together in the presence of an iron catalyst to produce ammoniaThe N2 and H2 molecules are adsorbed on to the surface of the ironThey react together then ammonia desorbs from the surface3) Products are desorbed (leave the surface).
11 Importance of Variable Oxidation States Variable oxidation states very important in heterogeneous catalysisE.g., use of V2O5 in the Contact ProcessIn the second stage of H2SO4 manufacture, sulfur(IV) oxide is oxidised to sulfur(VI) oxide (2-step mechanism):1. The vanadium(V) oxide catalyst oxidises sulfur(IV) oxide to sulfur(VI) oxide.SO2(g) + V2O5(s) → SO3(g) + V2O4(s)2. The vanadium(IV) oxide reacts with oxygen and the Vanadium(V) oxide is regeneratedV2O4(s) + 1/2O2(g) → V2O5(s)Overall:SO2(g) + 1/2O2(g) ↔ SO3(g)Step 1: Ox state of V decreases from +5 to +4V2O5 catalyst
13 Heterogeneous Catalysts Nature of catalystLarge surface area.Spread thinly over ceramic honeycomb.To make catalysts more efficient ...They can be expensive so more efficient catalysts mean that costs will be minimised.Increase their SA the larger the SA, the better the efficiencySpread the catalyst onto an inert support medium (or impregnate it into one). Increases the surface-to –mass ratio (a little goes a long way). This is usually for more expensive catalysts - e.g., catalytic converter in car finely divided Rh and Pt on a ceramic material.(Possibly mention PdENCat – worked on during MChem ... Polymer encapsulated Pd catalyst)
14 Heterogeneous Catalysts 1) Reactants adsorbed onto surface (onto active sites).weakens bondsbrings molecules closermore favourable orientation2) Reaction takes place.The reaction happens at the surface of the catalystThe places where the reactions happen are called “active sites”e.g., in the Haber Process, nitrogen and hydrogen react together in the presence of an iron catalyst to produce ammoniaThe N2 and H2 molecules are adsorbed on to the surface of the ironThey react together then ammonia desorbs from the surface3) Products are desorbed (leave the surface).
15 Homogeneous Catalysts Homogeneous catalyst is in the same phase as the reactants.Most examples involve reactions in solution with catalyst in solution.Some gas phase examples.
16 Homogeneous Catalysts Catalyst reacts with a reactant to form intermediate.2) Intermediate reacts to form product faster than original reactant (and regenerates catalyst).e.g. acid catalystX + Y productsX + H+ → HX+2) HX+ + Y → products + H+
17 A reaction of the S2O82- ion 2 I- + S2O82- → I SO42-Why would this be considered a slow reaction?
18 A reaction of the S2O82- ion Fe2+2 I- + S2O82- → I SO42-
19 Homogeneous Catalysts 2 Fe2+ + S2O82- → 2 Fe SO42-2) 2 Fe I- → 2 Fe2+ + I2Fe2+ is regenerated, so the two steps can repeat continuously.
20 Can also happen the other way around… 2 Fe I- → 2 Fe2+ + I22 Fe2+ + S2O82- → 2 Fe SO42-Fe3+ is regenerated, so the two steps can repeat continuously.
21 Redox TitrationsMay want to measure the conc of an oxidising or reducing agentDo a redox titrationSimilar to acid-base titration (find out how much acid is needed to react with a certain volume of base [or vice versa])C2O42- react with MnO4- acidified with excess dilute sulfuric acid (H+)
22 Homogeneous Catalysts: Autocatalysis The reaction is catalysed by one of the products (Mn2+) – this is called autocatalysis.Such a reaction starts slowly – uncatalysed rateAs conc of the products (also the catalyst) builds up, the reaction speeds up to catalysed rate.Then behaves as normal reaction ... Slows down as reactants are used upe.g. transition metal catalyst: works by metal varying oxidation stateOxidation of ethanedioic acid by manganate(VII) ionsDiscuss the reaction and draw the curve on p232 of textbookMn2+2MnO4- (aq)+16 H+ (aq) + 5 C2O42- (aq) → 2Mn2+ (aq) + 8 H2O (aq) +10 CO2 (aq)