Reaction Kinetics

  • Rate of reaction is the change in concentration of products/reactants per unit time

R = Unit: mol dm-3 s-1



Order of Reaction

  • For zero order, rate of reaction not dependant on conc. of that reactant but it is need for completion of reaction. Reaction continues until all of that reactant is used up.
  • Recognize shapes above in graphs given to work out order of reactant and construct rate equation. Calculate kk using data from graph(s).

Constructing rate equation using conc. data:

Using following data from the experiments construct a rate equation and calculate rate constant.


Relationship of Temperature and k

  • Increasing the temperature, increases the value of k
  • When temp. is increased, the k.e. of reacting molecules increases resulting in more successful collisions
  • Reactants change faster to products therefore conc. of reactants decreases.
  • Using fraction above, numerator increases and denominator gets smaller therefore k increases.


Reaction Mechanism

Rate-determining step

  •  the slowest step in a reaction mechanism
  • many reactions consist of a series of separate stages
  • each of these stages has its own rate and hence its own rate constant
  • the slowest step is known as the rate determining step


In a multistep reaction:

  • The rate of reaction is dependent on the slowest step that needs the highest activation energy.
  • Rate equation includes only reactants that are present in the rate-determining step.
  • The orders with respect to the reactants are the moles of the reactants in the rate determining step

Example 1

H2O2 + 2H3O+ + 2I¯ ——> I2 + 4H2O

Step 1 H2O2 + I¯ ——> IO¯ + H2O slow

Step 2 IO¯ + H30+ ——> HIO + H2O fast

Step 3 HIO + H30+ + I¯ ——> I2 + 2H2O fast

The rate determining step is STEP 1 as it is the slowest


Example 2

The reaction 2N2O5 —> 4NO2 + O2

Step 1     N2O5 —> NO2 + NO3

Step 2     NO2 + NO3 —> NO + NO2 + O2

Step 3     NO + NO3 —> 2NO2


The rate determining step is STEP 1 as it is the slowest. The rate equation for the reaction is rate = k [N2O5]


Measuring Reaction Rates

Sampling method that involves taking small sample of a reaction mixture at various times and then carrying out chemical analysis on sample.

Chemical Analysis

C4H9Br + OH C4H9OH + Br

  • Sampled removed at various times and quenched – stopping/slowing down reaction (e.g. cooling in ice)
  • OH conc. can be found using titration
  • Plot graph and calculate rate of reaction



Method that involves monitoring a physical property over a period of time

Change in Volume of Gas Produced

Mg + 2HCl MgCl2 + H2

  • Measure change in volume of gas using a gas syringe
  • Take down readings at regular intervals
  • Plot graph and calculate rate of reaction


Changes in Colour


  • I2 starts brown, fades through orange to yellow to colourless as iodine used up.
  • Colorimeter measures amount of light absorbed as it passes through solution; recorded as absorbance.
  • ‎Before experiment, create calibration curve by finding absorbance of different conc. I2 and plot a graph of concentration against absorbance
  • During experiment, measure absorbance from meter at regular intervals, and use calibration curve to convert values into concentrations



They work by providing an alternative reaction pathway with a lower Activation Energy

Lowering Ea results in there being a greater area under the curve showing that more molecules have energies in excess of the Activation Energy. Catalysts remain chemically unchanged at the end of the reaction.


  1. Homogeneous Catalysts
  2. Heterogeneous Catalysts


They do not affect the position of the equilibrium but they do affect the rate at which equilibrium is attained.


Homogeneous Catalysis

  • Catalyst and reactants in same physical state
  • Catalyst takes part in reaction:
    • Forms intermediate with reactant
    • Intermediate breaks down giving product
  • Rate of reaction dependant on conc. of catalyst

Examples of Homogenous Catalysis


Iodine-Peroxydisulfate Reaction

S2O82- + 2I > 2SO42- + I2

  • Mechanism of the catalysed reaction:
    • Reaction 1: reduction of Fe3+ to Fe2+ ions by I ions

      2I + 2Fe3+ > I2 + 2Fe2+

    • Reaction 2: oxidation of Fe2+ to Fe3+ ions by S2O82- ions
    • S2O82- + 2Fe2+ > 2SO42- + 2Fe3+

Oxides of Nitrogen and Acid Rain

  • SO3 can then react with water to form H2SO4


Heterogeneous Catalysis

  • Catalyst and reactants in different physical state
  • Catalyst provides surface on which reaction occurs
    • Reactant particles adsorbed on surface on collision with catalyst
    • Molecular rearrangement occurs – bonds in reactants break and new bonds in product formed
    • Product molecules desorbed from surface
  • Rate of reaction dependant on surface area of catalyst


Catalysis is thought to work in three stages

  1. Adsorption – formation of bonds with surface weakens bonds in gas molecules makes a subsequent reaction easier
  2. Reaction – adsorbed gases may be held on the surface increases chances of favourable collisions
  3. Desorption – the products are then released from the active sites (catalyst (diffusion away from surface)


Hetero = Adsorption + Reaction + Desorption


Examples of catalysts

  • Metals
    • Ni, Pt hydrogenation
    • reactions Fe Haber Process
  • Oxides
    • AlO3 dehydration
    • reactions V2O5 Contact Process


  • Honeycomb structure containing small beads coated with Pt, Pd and Rh


Possible catalytic process

  • Adsorption of NOx and CO
  • Weakening of covalent bonds within NOx and CO
  • Formation of new bonds between
    • Adjacent N atoms form N2
    • CO and O atoms form CO2
  • Desorption of N2 and CO2



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