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1. REACTION KINETICS (AS)
 1.Rate of reaction = change in concentration
of reactant or product over time
 Rate of reaction = [reactant]/time OR
[product]/time

2. 2.Concentration –time graphs
time
Conc of
a reactant Conc of reactant decreases
with time

3. time
Conc of
product
After certain time ,conc of
products becomes
constant
Conc of product
increases with
time

4.  a. Rate of reaction at time , t :
(instantaneous rate)
 draw a tangent to the concentration
vs time curve at time t
 the gradient of tangent = rate of
reaction

5. Example
time
[reactant ]
t
y
x
Gradient = y/x =
rate of reaction
at time , t
Unit : mol dm-3 s-1 or
mol dm-3 min-1

6.  Note :
 i)Average rate : rate measured over a
period of time
 Eg : rate = change in [reactant]/ t2 – t1
 ii)Initial rate : rate at almost t=0
 b. Rate of rxn is proportional to
concentration of most reactants
 Concentration increases, rate increases

7. Note : Rate is independent of
concentration of a reactant
 Concentration changes but rate is constant
 Zero order reaction
time
Conc of
reactant
Conc decreases with time
Constant gradient
Rate is constant

8. THEORIES OF REACTION
RATES
 1. Collision theory :
 a. reactions occur due to collision of
reactant particles
 b. not all collisions results in reaction
 effective collisions : collisions
between reacting particles that
results in a reaction

9.  c.Characteristics of effective collisions
:
 i) have favourable orientation
 eg C – C – C – C –Br + OH-
 C – C – C – C –OH + Br-
collision of an OH- with the
bromoethane molecule is unlikely to
result in a reaction if it hits the end of
the molecule away from the Br

10.  ii) possess a minimum energy = Ea
 (1) Definition : Activation energy ,Ea,
is the minimum energy required for a
reaction to take place
 High Ea  slow reaction
 (2) Ea is used to enable bonds in the
reactants to stretch and break as new
bonds form in the products

11.  2. Transition state theory :
 a. reactions takes place via transition
state in which reactants come together
 b. bond making and breaking occur
continuously and simultaneously
 In the transition state, bonds are in the
process of making and breaking.

12.  A-B + C  A + B-C
 A B C
transition state Bond forming
Bond breaking

13.  c. reaction profile/enthalpy diagram :
 Note :
 (1) Transition state is the highest
point in the reaction profile
 (2) Energy gap between reactants and
transition state = Ea
 (3) Ea forward rxn ≠ Ea reverse rxn

14. Reaction profile or energy / enthalpy
diagram for uncatalysed reactions
 exothermic reversible reaction
Extent of reaction
Energy
Products
Reactants
Transition state
Ea forward rxn
Ea reverse rxn
H

15. endothermic reversible reaction
Extent of rxn
Energy
Reactants
Products
Transition state
Ea reverse rxn
Ea forward rxn
H

16. d. Multi step reaction
 Reaction that takes place via an
intermediate
 Mechanism of rxn involves a multi
step reaction
 The intermediate will occur at a
minimum on the graph
 One minimum = one intermediate

17. Eg :
Step 1 : Reactants  Intermediate ,
H = positive
Step 2 : Intermediate  Products ,
H = negative
Overall : Reactants  Products ,
H = negative

18. Energy
Extent of rxn
Reactants
Products
Transition state 1
Transition state 2
Intermediate
Ea(1) Ea (2)
Overall H

19.  e. Reacting particles must possess
energy greater than or equal to the Ea
before they can react

20. FACTORS AFFECTING
RATE OF REACTION
 Concentration
 Temperature
 Catalyst

21.  I. Concentration of reactants
 1. conc increases , rate normally
increases
 ( exception : zero order )
 2. as concentration increases :
 frequency of collisions increases
 no of effective collisions increases
 rate of reaction increases

22.  3. Expt to show effect of concentration on
rate of reaction :
 Eg:
 Na2S2O3(aq) + 2HCl(aq)  2 NaCl(aq) +
H2O(l) + SO2(g) + S(s)
 a. Effect of [S2O3
2-] on rate of reaction
 b. Sulphur appears as particles of solid
 c. Measure time taken to block view of
cross/words under conical flask

23.  Experiment to show effect of concentration
on rate of reaction :
 Eg Na2S2O3 (aq) + 2HCl (aq) 
2 NaCl(aq) + H2O(l) + SO2(g) + S(s)
a. Effect of conc of S2O3
2- on rate of rxn
b. Sulphur appears as small particles of
solid
c. Measure time taken for enough sulphur to
form to block view of the cross/words
under conical flask

24.  d. Use different volumes of S2O3
2- but
keep volume of HCl constant
 e. H2O used to keep total volume of all
mixtures constant
 Hence volume of S2O3
2- used  conc
S2O3
2-
 eg : volume doubles , conc doubles

25. Mixture Volume of
S2O3
2-/cm3
Volume of
HCl/cm3
Volume of
H2O/cm3
Time/s
1 10 20 30
2 20 20 20
3 40 20 0

26.  Rate of reaction α 1/time
 From expt ,
 As volume of S2O3
2- increases,
 [S2O3
2-] increases , time taken
decreases
 Rate of reaction increases

27. [S2O3
2- ]
1 / time
Rate of reaction α [S2O3
2-]

28. II.Temperature
 1. When temperature increases :
 average speed of reacting particles
increases
 particles collide more frequently and
with greater energy
 no of particles with energy ≥ Ea
increases
 no of effective collisions increases
 rate of reaction increases

29.  2. Why does rate increase with
temperature?
 Molecules in a gas does not all have the
same speed.
 Their speeds and therefore their
energies are distributed according to the
Maxwell Boltzmann distribution curve

30.  Maxwell Boltzmann distribution curve
Energy/speed
Fraction or no of
molecules with
energy E
Most probable energy

31.  a. Shape : at a temp T , molecules in
a sample of gas have different
speed/energy
 Most probable speed/energy
corresponds to the maximum of the
curve.
 b. Area under the curve = total no of
molecules in the sample

32.  c. As temp increases ,
 curve flattens ( have a lower peak )
 more spread out ( moves to the right )
 however total no of molecules =
areas under the curves remains the
same

33.  Effect on Maxwell Boltzmann distribution curve
Energy/speed
No of molecules
with energy E Lower T
Higher T
Ea

34.  d. Shaded area = no of molecules with
energy ≥ Ea
 As temp increases ,
 Size of shaded area increases
 More molecules with energy ≥ Ea
 No of effective collisions increases
 Rate of reaction increases

35.  Note : At temp T and ( T + 10 K ) ,
 Size of shaded area doubles
 No of molecules with energy ≥ Ea
doubles
 Rate of reaction doubles

36.  e. Reactions with larger Ea are slower
but rise in temp has more
significant increase on the rate of
reaction with a higher Ea

37. III.Effect of catalyst ( catalysis )
 1.Catalysts are substances that affects the
rate of a chemical reaction without being
chemically changed themselves
 They are not consumed and are
regenerated at the end of the reaction

38.  Properties of catalyst:
 increase the rate of reaction
 amount of catalyst used affects the rate
which is proportional to the amount used
 required in small amount
 chemically unchanged after the reaction
 do not affect H

39.  2. Two types of catalyst :
 a. positive catalyst : increases rate of
reaction
eg ferum in Haber process
 b. negative catalyst / inhibitor : slows
down a reaction
eg glycerine or phosphoric acid
inhibits decomposition of hydrogen
peroxide

40.  3. Action of positive catalyst
 Provides alternative pathway with a
lower Ea
 More molecules with energy ≥ Ea
 No of effective collisions increases
 Rate of reaction increases
 Note : different catalyst can affect a
similar reaction differently

41. 4. Diagrams :
a. Enthalpy diagram or energy profile :
eg exothermic rxn
Reaction pathway
Energy
Reactants
Products
Ea catalysed rxn(lower)
Ea uncatalysed rxn

42. b. Maxwell Boltzmann distribution curve
( at a certain temp T )
Energy
No of molecules
with energy E
Ea uncatalysed
Ea catalysed (lower)

43.  For catalysed reaction :
 Size of shaded area increases
 No of molecules with energy ≥ Ea
increases
 No of effective collisions increases
 Rate of reaction increases
 Note : another factor affecting rate is
surface area ( higher surface area ,
faster reaction )

44. 5. Types of catalyst : 3 types
 a. Heterogeneous catalyst : catalyst is in a different
phase compared to reactants .
 Examples :
 Reaction Catalyst
 N2(g) + 3H2(g)  2NH3(g) ferum (s)
( Haber process )
2SO2(g) + O2(g)  2SO3(g) V2O5 (s)
( Contact process )
C2H4(g) + H2(g)  C2H6(g) Ni (s)
( Hydrogenation of alkenes in
manufacture of margarine )

45.  b. Homogeneous catalyst : catalyst is present in the
same phase as the reactants.
 Examples:
 Reaction Catalyst
 CH3COOH(aq) + C2H5OH(aq) H+ (aq)
 CH3COOC2H5(l) + H2O (l)
S2O8
2- (aq) + 2I- (aq) Fe2+(aq)
 2SO4
2- (aq) + I2 (aq) or Fe3+ (aq)

46.  c. Biological catalyst ( enzymes ):
 Proteins which catalyses chemical reactions in living
systems
 Are extremely specific , one enzyme normally
catalyses one reaction
 Example: amylase found in saliva. It is used to break
carbohydrates into simpler molecules.

47. Autocatalysis
 1. One of the product is a catalyst for the
reaction
 2. Reaction proceeds slowly at first at
uncatalysed rate
 until a significant amount of the product (
also the catalyst ) is established
 3. Then reaction will speed up to catalysed
rate
 Reaction will stop when reactants are
exhausted

48.  Eg :
 2 MnO4
- + 16 H+ + 5 C2O4
2- 
2 Mn2+ + 8 H2O + 10 CO2
catalyst

49. time
[ MnO4
- ]
Fast decrease in
conc
Faster reaction
Catalysed rate
Slow decrease in conc
Slow reaction
Uncatalysed rate

50. time
rate
Slow
Uncatalysed
rate
Fast
Catalysed rate