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Rates and reaction mechanism
► The
reaction mechanism is the sequence of
individual reaction steps that together
complete the transformation of reactants to
products
► The single reaction described by chemical
equation may involve several sequential
elementary steps
Reduction of NO2 by CO
►
►
Overall reaction is
NO2 + CO = NO + CO2
Two elementary steps
NO2 + NO2 = NO + NO3
NO3 + CO = NO2 + CO2
Reaction intermediates
►
►
►
Reaction intermediate: an entity created
during one of the elementary steps, but
consumed during a subsequent step
Intermediates are not transition states
Intermediate does not appear in overall
reaction equation
Intermediates are neither product
nor reactant
► Overall
reaction is
NO2 + CO = NO + CO2
 NO3 is product of first step:
NO2 + NO2 = NO + NO3
 NO3 is consumed in second step
NO3 + CO = NO2 + CO2
NO3 is an intermediate
NO2 and CO are reactants
NO and CO2 are products
Molecularity
► The
number of molecules that participate in
the step
 Unimolecular involves one molecule
 Bimolecular involves two molecules
 Termolecular (rare) involves three molecules
Laws and mechanism
► The
rate law for an elementary (single) reaction
step follows directly from the molecularity
Rate laws for overall reactions
► If
overall reaction occurs in single step, rate
law is obviously the same as for that step
► For the single-step reaction
CH3Br + OH- = Br- + CH3OH
CH 3 Br 
Rate  
 k[CH 3 Br ][OH  ]
t
Slow steps, bottlenecks and rate
equations
►
►
For two or more steps, one step will be
slower than the rest. This is rate
limiting. The reaction rate will not be
faster than the slowest step
In the reaction
NO2 + CO = NO + CO2
1. NO2 + NO2 = NO + NO3 slow
2. NO3 + CO = NO2 + CO2 fast
The molecules in the fast step do not participate in the
rate equation
NO2 
2
Rate  
t
 k[NO2 ]
Rate-limiting step has highest Ea
Rate laws inform about reaction
mechanism
►A
logical sequence for determining reaction
mechanisms from observed kinetic data
Some like it hot: The Arrhenius
equation
► It
is well known that all chemical processes
go faster as temperature increases
► This suggests
 Not all collisions result in product
 Fraction of collisions that results in products
increases with T
Collision theory
► Reactions
only occur when molecules collide
► Not all collisions result in reactions
 Not all collisions have the right orientation
 Not all collisions have sufficient energy
► Why
is that?
Transition state and activation
energy
► The
approach of A towards BC increases energy
due to repulsions between electrons
E
Maximum energy is transition state
► Approach
of A causes weakening of B – C
and formation of A – B
► At transition state, the energy is a maximum
E
Over the barrier
► The
B – C bond weakens and C leaves
► The A – B bond forms completely
► The energy lowers
► Final state compared to initial state depends
on relative strengths of the A – B and B – C
bonds
E
The activation energy and transition
state
The problem of energy
►
►
►
►
►
►
Not all molecules have same energy
Distribution in energy is broad
Distribution shifts to higher energy as T increases
Only small fraction have sufficient energy to leap barrier
Fraction increases with T
Therefore rate increases with T
The exponent
►
Fraction of molecules with sufficient energy is
given by
E
f e
►

a
RT
f is very, very small: activation energy of 50
kJ/mol yields value of
f = 10-9
► Increasing T by 10ºC causes f to increase by a
factor of 3
► Collision rate is much less sensitive to T change
► T-dependence of rate is determined by fraction
of collisions with correct energy, not total
number of collisions
The problem of orientation
► Not
all (molecular) orientations are equal
► Fraction of collisions that has correct
orientation is the steric factor – p
Reaction
No reaction
Putting it all together:
The Arrhenius equation
► Collision
► Rate
rate
Z [ A][ B]
pfZ[ A][ B]  k[ A][ B]
law
► But
f e
k  pZe
Orientation
factor


Ea
Ea
RT
Collision
factor
RT
 Ae

Ea
RT
Energy
factor
Experimental determination of Ea
► Plot
of ln k vs 1/T is linear
Ea
ln k  ln A 
RT
Ea = -R(slope)
► From two temperatures:
 k2   Ea  1 1 
ln        
 k1   R  T2 T1 
Graphical determination of Ea
► Exploring
factors that affect rate
Ea
ln k  ln A 
RT
Ea
d ln k

R
d1
T
Ways over (or around) the barrier
► Temperature
increases reaction rate by
increasing fraction of molecules with
sufficient energy to jump barrier
►A
catalyst lowers the barrier. A catalyst acts
to increase the reaction rate, but is not
consumed itself during the reaction
Catalysis
► Catalysis
is a fascinating field. It deals with
processes which may provide solutions for many of
the key problems we face. It provides food
through the ammonia synthesis. It creates
flexibility in the complex matrix of energy sources,
energy carriers and conversion and it contributes
in minimizing pollution.
J. Rostrup-Nielsen, Catalysis Today, 111 (2006) 4 11
Catalysts modify the pathway
► Addition
of chlorine catalyst increases rate of
decomposition of ozone into O2 – reason for the
destruction of ozone layer by CFCs
► Although two barriers are present, both are
smaller than the one without the catalyst, and
reaction rate is higher
Without catalysts, there would be no
life at all, from microbes to humans
► ENZYMES
are biological catalysts
► Most enzymes are proteins – large
molecules
► Have correct shape to bring reactant
molecules together in correct orientation
Enzyme