STPM Chemistry Form 6 Notes – Reaction Kinetics (Part 4 – Final)

by BerryBerryTeacher

in Berry Reference (Notes)

In this final part of the STPM Form 6 Chapter on Reaction Kinetics from Berry Berry Easy, you’ll be expose to the concepts of catalysis (homogeneous and homogeneous) and the effects of temperature of reaction rates and their rate constants. They will be explained using kinetics theory, Arhenius equations and Clausius-Clapeyron equation. It might seemed difficult at first go due to the relatively long formula, but it is exceptionally easy once you understand the underlying concept of it. So keep on trying and trying until you understand, it is worth the trouble as this will be used all the way to university level.

(Tips: The examples given are not exercises but you can still test yourself with the concepts. Try and do a Q&A with a partner and test your understanding of the topic subject. More importantly, understand and memorise all the processes shown below, they are berry berry important)

STPM Chemistry Form 6 Notes – Reaction Kinetics (Part 4 – Final)


  • only small amount of catalyst is needed for the reaction
  • catalysts lower the activation energy of the reaction
  • catalysts speed up the rate of reactions and has no effect on the yield of a reaction
  • catalysts are specific for one reaction

Homogeneous catalysis

  • catalyst that exists in the same phase as the reactant
  • an intermediate species is produced in the reaction

Example 1 (homogeneous catalyst in gas – NO)

Step 1: NO(g) + ½ O2(g) –> NO2(g) where NO2 is an intermediate compound
Step 2: NO2 (g) –> NO(g) + O (g) with sunlight
Step 3: O(g) + O2(g) –> O3(g)
Overall reaction / uncatalysed reaction: 3/2O2(g) –> O3(g)

Example 2 (homogeneous catalyst in aqueous solution – acid catalysis)

CH3COOCH3(aq) + H2O(l) –> CH3COOH(aq) + CH3OH(aq)

Example 3 (homogeneous catalyst in aqueous solution – Fe3+)

Step 1: 2Fe3+(aq) + 2I-(aq) –> 2Fe2+(aq) + I2(aq) where Fe2+ is an intermediate ion
Step 2: 2Fe2+(aq) + S2O82-(aq) –> 2Fe3+(aq) + 2SO42-(aq)
Overall reaction / uncatalysed reaction: S2O82-(aq) + 2I-(aq) –> 2SO42-(aq) + I2(aq)

Example 4 (homogeneous catalyst in aqueous solution – Br2)

Step 1: Br2(aq) + H2O2(aq) –> 2Br-(aq) + 2H+(aq) + O2(g)
Step 2: 2Br-(aq) + H2O2(aq) + 2H+(aq) –> Br2(aq) + 2H2O(l)
Overall equation / uncatalysed reaction: 2H2O2(aq) –> 2H2O(l) + O2(g)

Example 4 (homogeneous catalyst in aqueous solution – NaI)

Step 1: I-(aq) + H2O2(aq) –> IO-(aq) + H2O(l)
Step 2: IO-(aq) + H2O2(aq) –> I-(aq) + O2(g) + H2O(l)
Overall equation / uncatalysed reaction: 2H2O2(aq) –> 2H2O(l) + O2(g)

Heterogeneous catalysis

  • catalyst that exists in a different phase from the reactant
  • they are transition metals, the oxides of transition metals and the oxides of aluminium and silicon
  • adsorption process which involve formation of bonds between the reactant molecules and the atoms on the surface (active sites) of the catalyst (solid metal)
  • 4 steps in heterogeneous catalysis: reactant molecules are adsorbed on the surface, reactant molecules diffuse along the surface, reactant molecules react to form product molecules and molecules of product desorb from the surface

Example 1 (heterogeneous catalyst – gold, Au)

Decomposition of N2O
N2O(g) –Au–> N2(g) + 1/2O2(g)

Example 2 (heterogeneous catalyst – nickel, Ni)

Hydrogenation of food oils or semisolid fats

Example 3 (heterogeneous catalyst – platinum metal mixed with rhodium)

Oxidation of carbon monoxide and unburned hydrocarbon such as benzene
2CO(g) + O2(g) –Pt–> 2CO2(g)
C6H6(l) + 15/2O2(g) –Pt–> 6CO2(g) + 3H2O(l)
2NO(g) –Rd–> N2(g) + O2(g)

Example 4 (heterogeneous catalyst – V2O5)

Contact process
2SO2(g) + O2(g) –V2O5–> 2SO3(g)

Example 5 (heterogeneous catalyst – Iron)

Haber process
N2(g) + 3H2(g) –Fe–> 2NH3(g)

Example 6 (heterogeneous catalyst – Platinum)

Ostwald process
4NH3(g) + 5O2(g) –Pt–> 4NO(g) + 6H2O(g)

Example 7 (heterogeneous catalyst – Nickel, Palladium, Platinum)

Hydrogenation of alkenes (manufacture of margarine)
CH2=CH2(g) + H2(g) –> CH3CH3(g)

Effect of temperature on reaction rates and rate constants

  • Temperature increases causes the reaction rates increase
  • Reaction rate is doubled for every 10˚C rise in temperature

Kinetics Theory

  • Kinetic theory = mv2/2
  • Raising the temperature greatly increases the fraction of molecules having very high velocity and high kinetics energy.
  • Most molecules are likely to collide to react
  • Increase in the frequency of collision, this causes the fraction of effective collision increases (kinetic energy equal to or greater than the activation energy, Ea).

Arrhenius equation,

  • k = A e –Ea/RT
  • In natural logarithm, ln k = ln A – (Ea/RT)
  • The graph of ln k versus (1/T) should be linear and the slope of straight line is -(Ea/R)

Clausius-Clapeyron equation (two different temperatures, T2 and T1)

  • In natural logarithm, ln k2 – ln k1 = – (Ea/R) [(1/T2) – (1/T1)]
  • ln (k2/k1) = (Ea/R) [(1/T2) – (1/T1)]

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February 5, 2011
February 9, 2011

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