Transcript Document 7928417

Kinetics
How fast does a reaction (event) occur?
Reaction rates are controlled by:
•Nature of reactants
•Ability of reactants to meet
•Concentration of reactants
•Temperature
•Presence of a catalyst
Rate of pay = €10/hour
(changein concentration)
d [ A]


(changein time)
dt
UNITS: mol/L x 1/s =mol.L-1.s-1 or M.s-1
Kinetics
Change of reaction rate with time
Concentration and rate
A + B products
In general it is found that:
rate[A]m[B]n
The values of the exponents, m and
n, must be determined empirically
(by experiment).
We can replace  by = if we
introduce a rate constant, k.
Rate = k [A]m[B]n
This expression is the rate law
Rate Laws
Example: H2SeO3 + 6I- + 4H+ Se + 2I3- + 3H2O
Rate = k[H2SeO3]x[I-]y[H+]z
Experimentally found that x=1, y=3, z=2
Rate = k[H2SeO3][I-]3[H+]2
At 0C, k=5.0 x 105 L5 mol-5 s-1
(units of rate constant are such that the rate has units of mol.L-1.s-1)
Notice that exponents in rate law frequently are unrelated to reaction stoichiometry.
Sometimes they are the same, but we cannot predict this without experimental data!
Exponents in the rate law are used to describe the order of the reaction
with respect to each reactant. The overall order of a reaction is the sum
of the orders with respect to each reactant.
Determining exponents in a rate law
One way to do this is to study how changes in initial concentrations
affect the initial rate of the reaction
A + B products
Initial Concs Initial rate
[A]
[B] (mol L-1 s-1)
0.10
0.20
0.30
0.30
0.30
0.10
0.10
0.10
0.20
0.30
0.20
0.40
0.60
2.40
5.40
Rate = k [A]m[B]n
1-3:
[B] is constant.
Rate changes due only to [A]
m must be 1
3-5:
[A] is constant. When [B] is
doubled, rate increases by
factor of 4 (=22). When [B] is
tripled, rate increases by
factor of 9 (=32).
n must be 2
Concentration and Time-1st order reactions
Rate = -k[A]
Integrated rate law

A0
ln
At
 kt
Half-life: time required for
half of initial concentration of
reactant to disappear.
t1/2 = ln2/k
Concentration and Time-2nd order reactions
Simplest 2nd order: 2A  B
Rate = k[A]2
Integrated rate law
1
1

 kt
At A0
Half-life
t1/2 = 1/k[A]0
Half-life depends on initial concentration
Temperature dependence of reaction rates
Transition-state theory is used to explain what happens
when reactants collide. Most often a change in
momentum or direction simply occurs. Sometimes a
reaction occurs. Only combining molecules that have
kinetic energies at least as large as the activation
energy can surmount the barrier and produce products.
The difference in potential energy of products to
reactants is the heat of reaction (exothermic in this
case). PE decreases, therefore KE and T increase!
The Arrhenius equation relates the
activation energy to the rate constant
k  Ae Ea / RT
ln k  ln A  Ea / RT
Catalysis
Ozone depletion
Radio-activity
Unstable atomic nuclei may decay by emitting particles that are detected with
special counters. Alpha, beta, and gamma emission are common types of
radioactivity. In beta decay the emitted particles are electrons; in alpha decay they
are helium nuclei, and in gamma decay they are high energy photons.
Counters can be sensitive to either alpha, beta, or gamma-ray particles. The
rubidium isotope 37Rb87 decays by beta emission to 38Sr87, a stable strontium
nucleus:
37Rb
87
 38Sr87 + b.
From the following experimental data, calculate (a) the rate constant and (b) the
half-life of the rubidium isotope. From a 1.00 g sample of RbCl which is 27.85%
37Rb , an activity of 478 beta counts per second was found. The molecular weight
87
of RbCl is 120.9 g mole-1.