Post on 06-Jul-2018
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REACTION KINETICS (AS)
1.Rate of reaction = change in concentration
of reactant or product over time Rate of reaction = [reactant]/time OR
[product]/time
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2.Concentration
–
time graphs
time
Conc of
a reactant Conc of reactant decreaseswith time
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time
Conc of
productAfter certain time ,conc of
products becomes
constant
Conc of productincreases with
time
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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
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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
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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
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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
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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 thatresults in a reaction
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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 toresult in a reaction if it hits the end ofthe molecule away from the Br
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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
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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 theprocess of making and breaking.
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A-B + C A + B-C
A B C
transition stateBond formingBond breaking
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c. reaction profile/enthalpy diagram :
Note :
(1) Transition state is the highestpoint in the reaction profile
(2) Energy gap between reactants and
transition state = Ea
(3) Ea forward rxn ≠ Ea reverse rxn
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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
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endothermic reversible reaction
Extent of rxn
Energy
Reactants
Products
Transition state
Ea reverse rxnEa forward rxn
H
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d. Multi step reaction
Reaction that takes place via anintermediate
Mechanism of rxn involves a multistep reaction
The intermediate will occur at aminimum on the graph
One minimum = one intermediate
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Eg :Step 1 : Reactants Intermediate ,
H = positive
Step 2 : Intermediate Products ,H = negative
Overall : Reactants
Products ,H = negative
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Energy
Extent of rxn
Reactants
Products
Transition state 1
Transition state 2
Intermediate
Ea(1) Ea (2)
Overall H
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e. Reacting particles must possess
energy greater than or equal to the Ea
before they can react
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FACTORS AFFECTING
RATE OF REACTION
Concentration
Temperature
Catalyst
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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
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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
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Experiment to show effect of concentrationon 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 toform to block view of the cross/words
under conical flask
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d. Use different volumes of S2O32- but
keep volume of HCl constant
e. H2O used to keep total volume of allmixtures constant
Hence volume of S2O32- used conc
S2O32-
eg : volume doubles , conc doubles
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Mixture Volume of
S2O32-/cm3
Volume of
HCl/cm3Volume of
H2O/cm3
Time/s
1 10 20 30
2 20 20 20
3 40 20 0
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Rate of reaction α 1/time
From expt ,
As volume of S2O32- increases, [S2O3
2-] increases , time taken
decreases
Rate of reaction increases
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[S2O32-
]
1 / time
Rate of reaction α [S2O32-]
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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 ≥ Eaincreases
no of effective collisions increases rate of reaction increases
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2. Why does rate increase with
temperature?
Molecules in a gas does not all have thesame speed.
Their speeds and therefore their
energies are distributed according to the
Maxwell Boltzmann distribution curve
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Maxwell Boltzmann distribution curve
Energy/speed
Fraction or no of
molecules with
energy E
Most probable energy
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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 ofmolecules in the sample
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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
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Effect on Maxwell Boltzmann distribution curve
Energy/speed
No of molecules
with energy E Lower T
Higher T
Ea
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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
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Note : At temp T and ( T + 10 K ) ,
Size of shaded area doubles
No of molecules with energy ≥ Eadoubles
Rate of reaction doubles
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e. Reactions with larger Ea are slower
but rise in temp has more
significant increase on the rate ofreaction with a higher Ea
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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
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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
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2. Two types of catalyst :
a. positive catalyst : increases rate of
reaction
eg ferum in Haber process
b. negative catalyst / inhibitor : slowsdown a reaction
eg glycerine or phosphoric acid
inhibits decomposition of hydrogenperoxide
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3. Action of positive catalyst
Provides alternative pathway with alower Ea
More molecules with energy ≥ Ea No of effective collisions increases
Rate of reaction increases
Note : different catalyst can affect asimilar reaction differently
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4. Diagrams :
a. Enthalpy diagram or energy profile :
eg exothermic rxn
Reaction pathway
Energy
Reactants
Products
Ea catalysed rxn(lower)
Ea uncatalysed rxn
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b. Maxwell Boltzmann distribution curve
( at a certain temp T )
Energy
No of molecules
with energy E
Ea uncatalysed
Ea catalysed (lower)
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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 issurface area ( higher surface area ,
faster reaction )
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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 )
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b. Homogeneous catalyst : catalyst is present in thesame phase as the reactants.
Examples:
Reaction Catalyst
CH3COOH(aq) + C2H5OH(aq) H+ (aq)
CH3COOC2H5(l) + H2O (l)
S2O82-
(aq) + 2I-
(aq) Fe2+
(aq) 2SO4
2- (aq) + I2 (aq) or Fe3+ (aq)
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c. Biological catalyst ( enzymes ):
Proteins which catalyses chemical reactions in living
systems
Are extremely specific , one enzyme normallycatalyses one reaction
Example: amylase found in saliva. It is used to break
carbohydrates into simpler molecules.
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Autocatalysis
1. One of the product is a catalyst for thereaction
2. Reaction proceeds slowly at first atuncatalysed 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 areexhausted
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Eg :
2 MnO4- + 16 H+ + 5 C2O4
2-
2 Mn2+ + 8 H2O + 10 CO2
catalyst
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time
[ MnO4- ]
Fast decrease in
conc
Faster reaction
Catalysed rate
Slow decrease in conc
Slow reactionUncatalysed rate
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time
rate
Slow
Uncatalysed
rate
Fast
Catalysed rate