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.Concentration –time graphs

time

Conc of

a reactantConc of reactant decreases

with time

time

Conc of

productAfter certain time ,conc of

products becomes

constant

Conc of product

increases with

time

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

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

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

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

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

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

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

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.

A-B + C A + B-C

A B C

transition stateBond formingBond breaking

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

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

endothermic reversible reaction

Extent of rxn

Energy

Reactants

Products

Transition state

Ea reverse rxnEa forward rxn

H

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

Eg :

Step 1 : Reactants Intermediate ,

H = positive

Step 2 : Intermediate Products ,

H = negative

Overall : Reactants Products ,

H = negative

Energy

Extent of rxn

Reactants

Products

Transition state 1

Transition state 2

Intermediate

Ea(1) Ea (2)

Overall H

e. Reacting particles must possess

energy greater than or equal to the Ea

before they can react

FACTORS AFFECTING

RATE OF REACTION

Concentration

Temperature

Catalyst

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

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 [S2O32-] on rate of reaction

b. Sulphur appears as particles of solid

c. Measure time taken to block view of

cross/words under conical flask

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 S2O32- 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

d. Use different volumes of S2O32- but

keep volume of HCl constant

e. H2O used to keep total volume of all

mixtures constant

Hence volume of S2O32- used conc

S2O32-

eg : volume doubles , conc doubles

Mixture Volume of

S2O32-/cm3

Volume of

HCl/cm3

Volume of

H2O/cm3

Time/s

1 10 20 30

2 20 20 20

3 40 20 0

Rate of reaction α 1/time

From expt ,

As volume of S2O32- increases,

[S2O32-] increases , time taken

decreases

Rate of reaction increases

[S2O32- ]

1 / time

Rate of reaction α [S2O32-]

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

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

Maxwell Boltzmann distribution curve

Energy/speed

Fraction or no of

molecules with

energy E

Most probable energy

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

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

Effect on Maxwell Boltzmann distribution curve

Energy/speed

No of molecules

with energy E Lower T

Higher T

Ea

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

Note : At temp T and ( T + 10 K ) ,

Size of shaded area doubles

No of molecules with energy ≥ Ea

doubles

Rate of reaction doubles

e. Reactions with larger Ea are slower

but rise in temp has more

significant increase on the rate of

reaction with a higher Ea

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

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

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

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

4. Diagrams :

a. Enthalpy diagram or energy profile :

eg exothermic rxn

Reaction pathway

Energy

Reactants

Products

Ea catalysed rxn(lower)

Ea uncatalysed rxn

b. Maxwell Boltzmann distribution curve

( at a certain temp T )

Energy

No of molecules

with energy E

Ea uncatalysed

Ea catalysed (lower)

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 )

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 )

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)

S2O82- (aq) + 2I- (aq) Fe2+(aq)

2SO42- (aq) + I2 (aq) or Fe3+ (aq)

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.

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 catalysedrate

Reaction will stop when reactants are exhausted

Eg :

2 MnO4- + 16 H+ + 5 C2O4

2-

2 Mn2+ + 8 H2O + 10 CO2

catalyst

time

[ MnO4- ]

Fast decrease in

conc

Faster reaction

Catalysed rate

Slow decrease in conc

Slow reaction

Uncatalysed rate

time

rate

Slow

Uncatalysed

rate

Fast

Catalysed rate