This document discusses chemical kinetics and reaction rates. It begins by explaining that reaction rate is a measure of how fast a chemical reaction occurs and can be affected by factors like the physical state and concentration of reactants, temperature, and presence of catalysts. It then discusses these factors in more detail and how they influence the collision and orientation of reactant molecules. The document also covers concepts like reaction order, rate laws, activation energy, reaction mechanisms, and the effects of temperature on reaction rates based on the Arrhenius equation. In addition, it distinguishes between elementary reactions, reaction intermediates, transition states, and multistep reaction mechanisms.

2. Learning Outcomes
 Explain importance of chemical kinetics
 Apply principles of order of reaction kinetics
 Describe the order of reactions with suitable working
illustrations
 Describe the experimental methods of chemical
kinetics
 Describe the various types of reaction and their
chemical kinetics applied to electrochemistry
 Understand the applications of chemical kinetics and
reaction dynamics in electrochemistry

3. Introduction
 Reaction rate
 Measures the speed of a chemical reaction occurs
 Can be affected by multiple factors
○ Physical state of the reactants
○ Reactant concentration
○ Reaction temperature
○ Presence of catalyst

4. Factors that affect reaction rates
 Physical state of reactants
 Reactants collide to react
 The more readily reactant molecules collide, the faster the rate
 Depends on the phases of the reactants
 Solid with more surface area react faster
 E.g. Medicine in fine powder vs medicine in tablet
○ Fine powder – higher surface area, dissolves in stomach faster,
enters blood quicker
 Reactant concentrations
 Rate of reaction increases with reactant concentrations
 Reactant concentration increases, the frequency of reactant molecules
collide increases

5.  Reaction temperature
 Rate of reaction increases with temperature
 Temperature of reaction will affect the kinetic energies of
molecules
○ Molecules collide more frequently
○ Increase in rate of reaction
 Presence of catalyst
 Catalyst – agents which increase reaction rate
 Affect the collision of molecules
 Collision will provide sufficient energy in bond breaking

6. Reaction rates
 Speed of the reaction occurs – change in the
concentration of reactants or products per unit of time
 Reactant concentrations decrease
 Product concentrations increase
 Rate of reaction is usually expressed in molarity per
second (M/s)

7. A →B
Rate of reaction =
Change in concentration of A
=
- (Δ[A])
Change in time Δt
Rate of reaction =
Change in concentration of B
=
+ (Δ[B])
Change in time Δt

8. Question:
Calculate the rate of reaction at which A disappears over the time
interval from 10 to 70 s

9. Instantaneous rate
 Instantaneous rate
 Rate at a particular instant during the reaction
 Determined from the slope of the curve at a
particular point in time
 To determine the instantaneous rate at
specific t.
 Draw tangent line which passed through the
point at t
 Example, determination of instantaneous rate
at t = 600 s

10. Initial rate
 Initial rate
 Instantaneous rate at the moment the
reactants are mixed
 t = 0
 Product concentration is negligible
 To determine the initial rate (t = 0)
 Draw a tangent line to the curve at t =
0
 Example, determination of
instantaneous rate at t = 0 s

11. Reaction rates and stoichiometry
 Rate of reactants disappearance = Rate of products appearance
 Example: aA + bB → cC + dD
 Use stoichiometry of the reactants and products to determine
the rate of reaction
 Reactants: -ve sign → disappearance of reactants
 Products: +ve sign → appearance of products

12. Rate law
 For a general reaction,
aA + bB → cC + dD
 The rate law generally has the form
Rate = k [A]m[B]n
 The constant k = the rate constant.
 Magnitude of k is affected by the temperature
 Exponents m & n = reaction orders.
 If m & n =1, the rate is first order in A and is also first order in
B.
 Overall reaction order is the sum of the orders with respect to
each reactant in the rate law.
 Overall reaction order for the reaction above is 2, second order
overall

13. First-order reactions
 First-order reaction – rate depends on the concentration
of single reactant
 E.g. A → product
 Rate law is
 Integrated the rate law for first-order reaction
Can be rearranged as:

14.  When the first order rate equation is expressed as y = mx +c
 Therefore, a graph of ln[A] vs. time → straight line graph
 slope, m = -k
 c = ln[A]0

15. Second order reactions
 Second order reaction – rate depends either on a
reactant concentration raised to the second power ([x]2),
or on the concentration of 2 reactants each raised to the
first power ([y]1[x]1)
 E.g. 2A → product
 Differentiate the rate law for second-order reaction

16.  When the second order rate equation is expressed as y = mx +c
 Therefore, a graph of 1/[A]t vs. time → straight line graph
 slope, m = k
 c = 1/[A]0
y = mx + c

17. Zero-order reaction
 Zero-order reaction – the rate of disappearance of A is independent
of [A]
 The rate law for zero-order reaction is
 Integrated rate law for zero-order reaction is

18.  When the zero-order rate equation is expressed as y = mx +c
 Therefore, a graph of [A]t vs. time → straight line graph
 slope, m = -k
 c = [A]0
y = mx + c

19. Half life
 Half life of a reaction , t½ - the time required for the concentration
of a reactant to reach half its initial value,
 Substitute and t½ = t into first order reaction

20. Summary of reaction order

22. The collision model
 Rate of reaction is affected by
 Temperature
 Concentration of reactants
 The collision model – based on kinetic-molecular theory
 Molecules must collide to react
 The greater the no. of collisions per seconds, the greater
the reaction rate
 As reactant concentration increases, no. of collision
increases
 As temperature increases, molecules will move faster, and
hence more collisions

23. The orientation factor
 Molecules must be oriented in a certain way during collision for
a reaction to occur
 Relative orientation – atoms are suitably positioned to form new
bonds
 E.g. Cl + NOCl → NO + Cl2

24. Activation energy
 Molecules must possess a certain minimum amount of energy
to react
 Energy comes from kinetic energy of the colliding molecules
 Kinetic energy is used to change the potential energy of the
molecule - stretch, bend and ultimately break bond leading to
chemical reaction
 Activation energy, Ea - Minimum energy required to initiate a
chemical reaction
 Ea needs to overcome the barrier through collision with other
molecules in order for a reaction to occur

26. Arrhenius Equation
 The rate constant (k) of a reaction is dependent on temperature.
 Three factors are incorporated into Arrhenius equation:
 The fraction of molecules possessing an activating energy of
Ea.
 The number of collisions occurring per second
 The fraction of collisions that have the appropriate orientation
 All 3 factors incorporated into Arrhenius equation:
k = Ae-Ea/RT
k = rate constant Ea = activation energy
R = gas constant, 8.314 J/mol-K
T = absolute temperature
A = frequency factor

27.  Determination of activation energy
Arrhenius equation: k = Ae-Ea/RT
ln of Arrhenium equation:
Graph of ln k vs 1/T → Straight line
Slope, m = -Ea/R
y-intercept, b = ln A

28. Non-graphical calculation
 At two different temperatures, T1 and T2, a reaction has rate
constants of k1 and k2.
 Subtract ln k2 from ln k1:
 Rearrange the equation:

29. The order in which bonds are broken and formed and the
changes in relative positions of the atoms in the reaction

30. Elementary reactions
 Elementary reactions – reactions which occur in a single step
 E.g.
 Molecularity – no. of molecules that participate as reactants in an
elementary reaction
 Unimolecular – 1 molecule of reactant
N2O5 → 2NO2 + ½ O2
 Bimolecular – 2 molecules of reactants; involving collision
CH3COOC2H5 + H2O → CH3COOH + C2H5OH
 Termolecular – 3 molecules of reactants; simultaneous
collision
○ Less probable than unimolecular and bimolecular
2NO + O2 → 2NO2

31. Multistep mechanism
 Multistep mechanism – sequence of elementary reactions
 E.g. reaction of NO2 and CO
 The reaction could not occur in only single step
 Undergoes multisteps - 2 elementary steps (2 step mechanism)
 1st step: 2 NO2 collide and an O atom transferred from one to
another
 2nd step: resultant NO3 collides with CO molecule, an O atom
from NO3 transferred to CO
 Intermediate - Molecules which formed in 1 elementary reaction
and consumed in another elementary reaction. E.g. NO3
 Usually stable – can be identified and isolated

32.  Transition state – unstable and cannot be isolated
 E.g. Stabilisation of von Willebrand factor A2 by Ca
 vWF-A2 : serum glycoprotein that mediates platelet adhesion to injured vascular
endothelium

33. Thermodynamic vs Kinetic control
 A reactive intermediate may sometimes form
different products depending on conditions as
shown by energy diagram
 Conversion of A → C goes through 1 intermediate
(B) and 2 transition states
 Conversion of A → D also goes through B and 2
transition states
 Formation of product C or D
 Determined by reaction conditions
 Since both products form from B, only 2nd step of reaction
will be considered

34. Question:
 Which product forms faster?
 Which product is more stable?

35. Rate laws for elementary reaction
 If a reaction is elementary, its rate law is based directly on its molecularity
 E.g.
Rate law:

36. Rate-determining step for a
multistep mechanism
 Rate-determining step – slowest step in a multistep reaction
which limits the overall reaction rate
 Which step is the rate-determining step?
 Rate law of the overall reaction will be the slow, rate-determining
step

37. Question:
Gas phase reaction of nitric oxide (NO) with bromine (Br2) happened
in 2 elementary steps
1. What is the overall reaction equation?
2. The experimental law for the overall reaction is
What can you say about the relative rates of the two steps of the
mechanism?

38. Mechanism with fast initial step
 As discussed earlier, it is known that Step 2 of the reaction below is
the rate determining step
And the overall reaction is
However, NOBr2 is an intermediate and will not accumulate, therefore
the concentration of NOBr2 is unknown
So [NOBr2] needs to be substituted

39.  Reverse step 2
 As in a equilibrium reaction, rate of forward = rate of reverse
Equate the rate law of forward and reverse reaction of Step 1
Rearrange the equation,
Therefore the rate law of overall reaction is

40. Catalysis
 Catalyst – substance that changes the speed of a chemical reaction
without undergoing a permanent chemical change itself
 Catalysts increase the rate of a reaction by decreasing the
activation energy of the reaction
 E.g. decomposition of H2O2 with the addition of NaBr catalyst

41.  Catalyst – will not be consumed in the reactions
 Homogenous catalyst – catalyst that is present in the same phase as the
reactant in the reaction mixture
 codissolved in a solvent with the reactants
 Heterogeneous catalyst – catalyst that is present in different phase
from the reactant molecules
 offers the advantage that products are readily separated from the catalyst
 Initial step of cataysis: Adsorption (binding of catalyst to reactants)

42.  Heterogeneous catalyst
C2H4 (g) + H2 (g) → C2H6 (g)
Catalyst: solid nickel, palladium, platinium