Transcript From rate law to reaction mechanism

From rate law to reaction mechanism
Products of a reaction can never be produced faster than the
rate of the slowest elementary reaction - rate determining
step
Experimental data for the reaction between NO2 and F2
indicate a second-order rate
Overall reaction: 2 NO2(g) + F2(g)  2FNO2(g)
Rate = k [NO2] [F2]
From the rate law : reaction does not take place in one step
How can a mechanism be deduced from the rate law?
Rate determining step must involve NO2 and F2 in 1:1 ratio
Possible mechanism
Step 1
NO2(g) + F2(g)  FNO2(g) +F(g)
slow
Step 2
NO2(g) + F(g) -> FNO2(g)
fast
Overall: 2 NO2(g) + F2(g)  2FNO2(g)
Rate for step 1 = k1 [NO2] [F2]
Rate for step 2 = k2 [NO2] [F]
rate determining step
For the reaction:
2 H2(g) + 2NO(g)  N2(g) + H2O(g)
The observed rate expression is: rate = k[NO]2 [H2]
The following mechanisms have been proposed. Based on
the rate law can any mechanism be ruled out?
k1
Mechanism I 2 H2(g) + 2NO(g)  N2(g) + H2O(g)
Mechanism II
Mechanism III
k2
H2(g) + NO(g)  N(g) + H2O(g)
k3
NO(g) + N(g)  N2(g) + O(g)
k4
O(g) + H2 (g)  H2O(g)
slow
fast
fast
k5
H2(g) + 2NO(g)  N2O(g) + H2O(g) slow
k6
H2(g) + N2O(g)  N2(g) + H2O(g)
fast
Mechanism I
rate = k1[H2]2 [NO]2
not possible
Mechanism II
rate = k2[H2] [NO]
not possible
Mechanism III
rate = k5[H2] [NO]2
possible
If mechanism III is a possible mechanism, try to detect N2O
experimentally to confirm mechanism.
Mechanism involving an initial fast reaction that produces an
intermediate, followed by a slower second step.
Rate law of an elementary step must be written with respect to
the reactants only; an intermediate cannot appear in the rate
law
2 NO(g) + O2(g)  2 NO2(g)
rate = k [NO]2 [O2]
Experiments indicate that an intermediate is involved in this
reaction
Step 1
k1
NO(g) + O2 (g) 

OONO(g)
fast, equilibrium
k-1
k2
Step 2
OONO(g) + NO(g)
 2 NO2(g)
slow
Rate law for second step = k2 [NO] [OONO]
Cannot be compared with experimental rate law since OONO
is an intermediate
Rate of production of OONO = k1 [NO] [O2]
Rate of consumption of OONO to NO and O2 = k-1 [OONO]
Rate of forward and reverse of 1 > rate of 2; equilibrium is
established in 1 before significant amount of OONO
consumed to form NO2
At equilibrium
k1 [NO] [O2] = k-1 [OONO]
k1 [OONO]

K
k-1 [NO][O 2 ]
Rate = k2 [NO] [OONO] = k2 [NO] (K [NO] [O2])
= k2 K [NO]2 [O2] = k’ [NO]2 [O2]
Stratospheric Ozone
Chapman mechanism (1930’s)
(1) O2 + light  2O.
(2) O. + O2 + M  O3 + M
(3) O3 + light  O. + O2
(4) O. + O3  2 O2
For O3 concentration to be in balance:
Rate of production of O3 by reactions (1) and (2) must equal
rate at which is O3 consumed in (3) and (4)
In the 1960’s data published showed that reaction (4) was
much too slow to balance levels of O3
Paul Crutzen realized that rate constants could not explain
measured distribution of ozone in the stratosphere.
Using experimental data, Crutzen presented a model:
NO + O3  NO2 + O2
NO2 + O.  NO + O2
Overall: O. + O3  2O2
Contributed to reaction (4) in Chapman mechanism
1995 Nobel Prize in Chemistry
Kinetics and Equilibrium
For a reaction which occurs in a single elementary step
NO(g) + O3 (g)
k1
NO2 (g) + O2 (g)
k-1
Rate of forward reaction = k1 [NO] [O3]
Rate of reverse reaction = k-1 [NO2] [O2]
At equilibrium:
rate of forward reaction = rate of reverse reaction
k1 [NO]eq [O3]eq = k-1 [NO2]eq [O2]eq
where the eq denotes equilibrium concentrations
[NO2 ]eq [O2 ]eq k1

K
[NO] eq[O3 ]eq k 1
where K is the equilibrium constant
For the reaction: 2NO(g) + 2H2(g)
N2(g) + 2H2O(g)
K = ([N2] [ H2O]2) /([ NO]2 [H2]2)
The reaction occurs through three elementary reactions
NO(g) + NO(g)
k1
k-1
N2O2(g) + H2(g)
N2O(g) + H2(g)
k2
k-2
k3
k-3
N2O2(g)
N2O(g) + H2O(g)
N2(g) + H2O(g)
At equilibrium
k1 [NO]2eq = k-1 [N2O2]eq
k2 [N2O2]eq [H2]eq = k-2 [N2O]eq [H2O]eq
k3 [N2O]eq [H2]eq = k-3 [N2]eq [H2O]eq
K1 
K2 
[N2O2 ]eq

k1
k 1
[NO]2eq
[N2O]eq [H2O]eq
[N2O2 ]eq [H2 ]eq

k2
k 2
[N2 ]eq [H2O]eq k 3
K3 

[N2O]eq [H2 ]eq k 3
K  K 1K 2K 3 
[H2O]2eq [N2 ]eq
k 1k 2k 3

k 1k 2k 3 [NO]2eq [H2 ]2eq
Chain reactions
Reaction which proceeds through a series of elementary
steps, some of which are repeated many times.
A highly reactive intermediate
reacts to produce another
highly reactive intermediate
Chain reactions include reactions in:
Atmosphere - ozone depletion
Explosions
Polymerization
Nuclear Fission
Anti-oxidants
Many chain reactions involve free-radicals - atoms or
molecules with one or more unpaired electron; formed by
homolytic cleavage of a covalent bond
Overall: H2(g) + Cl2(g)  2 HCl(g)
Cl2(g) + light  2 Cl(g)
+ H2 (g)  HCl(g) + H(g)
propagation
+ Cl2(g)  HCl(g) + Cl(g)
propagation
Cl(g)
H(g)
+ Cl(g)  Cl2(g)
termination
+ H(g)  H2(g)
termination
+ Cl(g)  HCl(g)
termination
Cl(g)
H(g)
H(g)
initiation
Stratospheric Ozone depletion by chlorofluorocarbons
(CFC’s)
CCl2F2 + hn  CF2Cl + Cl
CCl2F2 + O  CF2Cl + ClO
Cl
reactions which generate
free radicals
+ O3  ClO + O2
ClO
+ O  Cl + O2
Overall Reaction
O
+ O3 -> 2O2
“Ozone Depletion”
1995 Nobel Prize in Chemistry
Polymerization
initiation
propagation
termination