Transcript ppt

Chemical Kinetics
How rapidly reactions proceed - rate of reaction
Details of process from reactants to products - mechanism
Thermodynamics determines the direction in which reactions
proceed spontaneously and the conditions at equilibrium,
but not the rate at which equilibrium is reached.
For a complete picture of a chemical reaction need both the
thermodynamics and kinetics of a reaction.
N2(g) + 3H2(g)  2NH3(g)
At 298 K
DGo = -33.0 kJ
DHo = -92 kJ; exothermic reaction
K = 6.0 x 105 ; thermodynamically favored at 298 K,
Rate is slow at 298K
Commercial production of NH3 is carried out at temperatures
of 800 to 900 K, because the rate is faster even though K is
smaller.
Thermodynamic functions are state functions (DG, DH, DE)
Thermodynamic functions do not depend on the mechanism
of the reaction.
The rate of the reaction is very dependent on the path of the
process or path between reactants and products.
Kinetics reveals information on the rate of the reaction and
the mechanism (path) of the reaction.
Thermodynamics vs Kinetics
A+BC+D
K1
A+BE+F
K2
If K1 > K2 - products C & D are thermodynamically favored
over E & F.
If products observed are C & D: reaction is
thermodynamically controlled
If products observed are E & F: reaction is kinetically
controlled
Zn2+(aq) + S2-(aq)  ZnS(s)
K = 1/Ksp = 2.2 x 1023
Fe2+(aq) + S2-(aq)  FeS(s)
K = 1/Ksp = 2.7 x 1018
On addition of S2- to an aqueous solution containing both
Zn2+ and Fe2+, ZnS precipitates first - reaction is
thermodynamically controlled.
(1) 2NO(g) + O2(g)  2NO2(g)
(2) 2CO(g) + O2(g)  2CO2(g)
Both have large values of K; both are thermodynamically
favored in the forward direction
Reaction (1) is very fast; reaction (2) slow
Reactions are kinetically controlled
Result is brown color of air due NO2 and buildup of CO in
the air
Rates of Reactions
Zn(s) + 2 H+(aq)  Zn2+(aq) + H2(g)
Rate of a reaction: change in concentration per unit time
average reaction rate =
change in concentration
change in time
Units of rate: concentration / time
NO2(g) + CO(g)  NO(g) + CO2 (g)
average reaction rate =
[NO]final - [NO]initial
tfinal - tinitial
NO2(g) + CO(g)  NO(g) + CO2 (g)
Time (s)
0
50
100
150
200
[NO] mol L-1
0
0.0160
0.0240
0.0288
0.0320
Average rate 1st 50 seconds = 3.2 x 10-4 mol L-1
Average rate 2nd 50 seconds = 1.6 x 10-4 mol L-1
Average rate 3rd 50 seconds = 9.6 x 10-5 mol L-1
Instantaneous Rate - rate at a particular moment in time
NO2(g) + CO(g)  NO(g) + CO2 (g)
rate =
D[NO]
D[CO2] - D[NO2] - D[CO]
=
=
=
Dt
Dt
Dt
Dt
For an infinitesimally small changes, the instantaneous rate
rate =
d[NO]
d[CO2]
d[NO2] - d[CO]
=
dt =
dt =
dt
dt
For a general reaction:
aA + bB  xC + yD
1 d[D] =
1 d[C]
rate = x d t = y d t
-
1 d[A] = a
dt
1 d[B]
b dt
initial rate (t = 0)
Rate (at 5 weeks) = -6.3 x 10-3
Rate (at 10 weeks) = -2.6 x 10-3
Factors affecting rates of reactions
a) Nature of reactants
2NO(g) + O2(g)  2NO2(g)
fast
2CO(g) + O2(g)  2CO2(g)
slow
b) Concentration of reactants: reactions proceed by
collisions between reactants
c) Temperature: In general, as T increases, rate increases
d) Catalyst: increases rate of reaction
e) Surface: larger surface area increases rate of reaction
f) Nature of solvent
Effect of concentration
Effect of temperature
Effect of surface area
Rate Laws and Rate Constant
Initial reaction rate: instantaneous rate of change in
concentration of a species at the instant the reaction begins.
(products present later in the reaction may affect the rate)
2 N2O5 (g)  4 NO2 (g) + O2 (g)
Perform a series of experiments with different initial
concentrations of N2O5 (g).
For each, monitor the concentration of N2O5 (g) versus time
Determine the initial rate of reaction at t = 0
2 N2O5 (g)  4 NO2 (g) + O2 (g)
For this reaction, experiments indicate:
rate  initial concentration of N2O5 (g)
rate = k x initial concentration of N2O5 (g)
Rate = - d[N2O5] / dt = k [N2O5] Rate law
k = 5.2 x 10-3 s-1
k is the specific rate constant
Units of k depends on the rate law
Rate laws are determined experimentally
2 NO2(g)  2 NO(g) + O2 (g)
d[NO2]
rate =
= k [NO2]2
dt
k = 0.54 (mol NO2)-1 s-1
Rate Law
2 NO2(g)  2 NO(g) + O2 (g)
Rate = k [NO2]2
For a general reaction:
aA + bB  cC + dD
1 d[D] =
1 d[C]
rate = c
= d
dt
dt
-
1 d[A] = a
dt
rate = k [A]m [B]n
1 d[B]
b dt
Rate Law
For a reaction k has a specific value; k for the reaction
changes with temperature
Note: m need not equal a; n need not equal b
Order of a Reaction
rate = k [A]m [B]n
Reaction is mth order in A and nth order in B
Overall reaction order = m + n
The reaction order is determined by the experimentally
determined rate law
N2O5(g)  N2O4(g) + 1/2 O2(g)
Rate = k [N2O5]
First order reaction
For a 1st order reaction, units of k: time-1
C2H6(g)  2 CH3(g)
rate = k [C2H6]2
second order reaction
2NO2(g)  2NO(g) + O2 (g)
Rate = k [NO2]2
second order reaction
For 2nd order reactions, units of k: concentration-1 time-1
S2O82-(aq) + 3 I-(aq)  2 SO42-(aq) + I3-(aq)
(S2O82- persulfate)
Rate of disappearance of S2O82- = k [S2O82-(aq)] [I-(aq)]
Overall reaction order = 1 + 1 = 2
1st order in S2O82-(aq) and 1st order in I-(aq)
If the above reaction was carried out at a high concentration
S2O82-(aq), and experimentally:
Rate of disappearance of S2O82- = k’ [I-(aq)]
pseudo first-order reaction
2 NH3(g)  N2 (g) + 3 H2(g)
Rate of disappearance of NH3 = k
zero order reaction