Kinetics

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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 UNITS mol/L x 1/s
mol.L-1.s-1 or M.s-1
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Kinetics
Change of reaction rate with time
Concentration and rate A B ?products In general
it is found that rate?AmBn 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
AmBn This expression is the rate law
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Rate Laws
Example H2SeO3 6I- 4H ?Se 2I3-
3H2O Rate kH2SeO3xI-yHz Experimentally
found that x1, y3, z2 Rate
kH2SeO3I-3H2 At 0?C, k5.0 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 (6th order in example above).
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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 Rate k AmBn 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
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Concentration and Time-1st order reactions
Rate kA Integrated rate law
We can show that A plot of lnAt versus t is a
straight line y mx c with slope -k and y
intercept lnA0.
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Concentration and Time-1st order reactions
A plot of lnAt versus t is a straight line
with slope -k and y intercept lnA0.
Half-life time required for half of initial
concentration of reactant to disappear. Set At
½A0 t1/2 ln2/k
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Concentration and Time-2nd order reactions
Simplest 2nd order 2A ? B Rate
kA2 Integrated rate law
Half-life t1/2 1/kA0 Half-life depends on
initial concentration
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Temperature dependence of reaction rates
Activation Energy In order to form products,
bonds must be broken in the reactants. Bond
breakage requires energy.
The Arrhenius equation relates the activation
energy to the rate constant
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• Activation Energy
• Consider the reaction between Cl and NOCl
• If the Cl collides with the Cl of NOCl then the
  products are Cl2 and NO.
• If the Cl collided with the O of NOCl then no
  products are formed.

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Arrhenius

• Arrhenius discovered most reaction-rate data
  obeyed the equation

k is the rate constant, Ea is the activation
energy, R is the gas constant (8.314 J/mol-K) and
T is the temperature in K. A is called the
frequency factor. A is a measure of the
probability of a favorable collision. Both A and
Ea are specific to a given reaction.
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Catalysis
A catalyst provides a reaction with an alternate
pathway that has a lower energy of activation. A
catalyst is not consumed in a reaction.
Enzymes are biological catalysts.
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Nerve Agents-Inhibition of Acetylcholinesterase
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Ozone depletion
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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 37Rb87 ? 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 37Rb87, an activity of 478
beta counts per second was found. The molecular
weight of RbCl is 120.9 g mole-1.
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Summary Of Decay Types
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94Pu244

• 94Pu239 is used for nuclear weapons and for
  energy

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