Chemical Kinetics Terminology (CIE A Level Chemistry)

• The rate of reaction refers to the change in the amount or concentration of a reactant or product per unit time and can be found by:
• Measuring the decrease in the concentration of a reactant OR
• Measuring the increase in the concentration of a product over time
  • The units of rate of reaction are mol dm^-3 s^-1 

Rate equation

• The following general reaction will be used as an example to study the rate of reaction

D (aq) → E (aq) + F (g) 

• The rate of reaction at different concentrations of D is measured and tabulated

  • Note that the remains constant for all values

Rate of reactions table

[D] (mol dm-3)	Rate (mol dm-3 s-1)	(s-1)
3.00	2.00 x 10-3	6.67 x 10-4
2.00	1.33 x 10-3	6.67 x 10-4
1.00	6.60 x 10-4	6.67 x 10-4

• A directly proportional relationship between the rate of the reaction and the concentration of D is observed when a graph is plotted

The gradient of the graph corresponds to  and remains constant throughout

• Rate equations can only be determined experimentally and cannot be found using the stoichiometric equation

Rate of reaction = k [A]^m [B]ⁿ

• • Where [A] and [B] = concentrations of reactants
  • m and n = orders of the reaction
• For example, the formation of nitrogen gas (N₂) from nitrogen oxide (NO) and hydrogen (H₂): 

2NO (g) + 2H₂ (g) → N₂ (g) + 2H₂O (g)

• • The rate equation for this reaction is:

rate = k [NO]² [H₂]

• As mentioned before, the rate equation of the reaction above cannot be deduced from the stoichiometric equation but can only experimentally be determined by:
  • Changing the concentration of NO and determining how it affects the rate while keeping [H₂] constant
    • This shows that the rate is proportional to the square of [NO]

Rate = k₁ [NO]²

• • Then, changing the [H₂] and determining how it affects the rate while keeping [NO] constant
    • This shows that the rate is proportional to [H₂]

Rate = k₂ [H₂]

• • Combining the two equations gives the overall rate equation (where k = k₁ + k₂)

Rate = k [NO]² [H₂]

Order of reaction

• The order of reaction shows how the concentration of a reactant affects the rate of reaction
  • It is the power to which the concentration of that reactant is raised in the rate equation
  • The order of reaction can be 0, 1,2 or 3
  • When the order of reaction of a reactant is 0, its concentration is ignored
• The overall order of reaction is the sum of the powers of the reactants in a rate equation
• For example, in the following rate equation, the reaction is:

Rate = k [NO₂]² [H₂]

• • Second-order with respect to NO
  • First-order with respect to H₂
  • Third-order overall since 2 + 1 = 3

Half-life

• The half-life (t_1/2) is the time taken for the concentration of a limiting reactant to become half of its initial value

Rate-determining step & intermediates

• The rate-determining step is the slowest step in a reaction
• If a reactant appears in the rate-determining step, then the concentration of that reactant will also appear in the rate equation
• For example, the rate equation for the reaction below is rate = k [CH₃Br] [OH^–]

CH₃Br + OH^– → CH₃OH + Br^–

• This suggests that both CH₃Br and OH^– take part in the slow rate-determining step
• This reaction is, therefore, a bimolecular reaction
  • Unimolecular: one species involved in the rate-determining step
  • Bimolecular: two species involved in the rate-determining step
• The intermediate is derived from substances that react together to form it in the rate-determining step
  • For example, for the reaction above the intermediate would consist of CH₃Br and OH^–

The intermediate formed during the reaction of CH₃Br and hydroxide ions

The intermediate is formed from the species that are involved in the rate-determining step (and thus appear in the rate equation)