Chemical Kinetics. Chemical kinetics - speed or rate at which a reaction occurs How are rates of...
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Chemistry 232 Chemical Kinetics
Chemical Kinetics. Chemical kinetics - speed or rate at which a reaction occurs How are rates of reactions affected by Reactant concentration? Temperature?
Chemical kinetics - speed or rate at which a reaction occurs
How are rates of reactions affected by Reactant concentration?
Temperature? Reactant states? Catalysts?
Slide 3
The Instantaneous Reaction Rate Consider the following reaction
A + B C Define the instantaneous rate of consumption of reactant A,
A
Slide 4
Reaction Rates and Reaction Stoichiometry Look at the reaction
O 3 (g) + NO(g) NO 2 (g) + O 2 (g)
Slide 5
Another Example 2 NOCl (g) 2 NO + 1 Cl 2 (g) WHY? 2 moles of
NOCl disappear for every 1 mole Cl 2 formed.
Slide 6
The General Case a A + b B c C + d D rate = -1 d[A] = -1 d[B] =
+1 d[C] = +1 d[D] a dt b dt c dt d dt Why do we define our rate in
this way? Removes ambiguity in the measurement of reaction rates
Obtain a single rate for the entire equation,
Slide 7
Alternative Definition of the Rate Rate of conversion related
to the advancement of the reaction, . V = solution volume
Slide 8
An Example Examine the following reaction 2 N 2 O 5 (g) 4 NO 2
(g) + O 2 (g) N2O5N2O5 NO 2 O2O2 Initial n 00 Change -2 +4 ++ Final
n - 2 44
Slide 9
The N 2 O 5 Decomposition Note constant volume system
Slide 10
The Rate Law Relates rate of the reaction to the reactant
concentrations and rate constant For a general reaction a A + b B +
c C d D + e E
Slide 11
Rate Laws (Contd) The only way that we can determine the
superscripts (x, y, and z) for a non- elementary chemical reaction
is by experimentation. Use the isolation method (see first year
textbooks).
Slide 12
For a general reaction x + y + z = reaction order e.g. X = 1; Y
= 1; Z = 0 2nd order reaction (x + y + z = 2) X = 0; Y = 0; Z = 1
(1st order reaction) X = 2; Y = 0; Z = 0 (2nd order)
Slide 13
Integrated Rate Laws The rate law gives us information about
how the concentration of the reactant varies with time How much
reactant remains after specified period of time? Use the integrated
rate laws.
Slide 14
First Order Reaction A product Rate = v = - d[A]/d t = k[A] How
does the concentration of the reactant depend on time? k has units
of s -1
Slide 15
The Half-life of a First Order Reaction For a first order
reaction, the half-life t 1/2 is calculated as follows.
Slide 16
Radioactive Decay Radioactive Samples decay according to first
order kinetics. This is the half-life of samples containing e.g. 14
C, 239 Pu, 99 Tc. Example
Slide 17
Second Order Reaction A productsv = k[A] 2 A + B products v =
k[A][B] Case 1 is 2 nd order in A Case 2 is 1 st order in A and B
and 2 nd order overall
Slide 18
The Dependence of Concentration on Time For a second order
process where v = k[A] 2
Slide 19
Half-life for This Second Order Reaction. [A] at t = t = [A]
0
Slide 20
Other Second Order Reactions Examine the Case 2 from above A +
B productsv = k[A][B]
Slide 21
Reactions Approaching Equilibrium Examine the concentration
profiles for the following simple process. A B
Slide 22
Approaching Equilibrium Calculate the amounts of A and B at
equilibrium.
Slide 23
The Equilibrium Condition At equilibrium, v A,eq = v B,eq.
Slide 24
Temperature Dependence of Reaction Rates Reaction rates
generally increase with increasing temperature. Arrhenius Equation
A = pre-exponential factor E a = the activation energy
Slide 25
Rate Laws for Multistep Processes Chemical reactions generally
proceed via a large number of elementary steps - the reaction
mechanism The slowest elementary step the rate determining step
(rds).
Slide 26
Elementary steps and the Molecularity Kinetics of the
elementary step depends only on the number of reactant molecules in
that step! Molecularity the number of reactant molecules that
participate in elementary steps
Slide 27
The Kinetics of Elementary Steps For the elementary step
unimolecular step For elementary steps involving more than one
reactant a bimolecular step
Slide 28
For the step a termolecular (three molecule) step. Termolecular
(and higher) steps are not that common in reaction mechanisms.
Slide 29
The Steady-State Approximation Examine the following simple
reaction mechanism Rate of product formation, v p, is proportional
to the concentration of an intermediate.
Slide 30
What Is an Intermediate? B is an intermediate in the above
reqction sequence. A species formed in one elementary step of a
reaction mechanism and consumed in one or more later steps.
Intermediates generally small, indeterminate concentrations.
Slide 31
Applying the Steady State Approximation (SSA) Look for the
intermediate in the mechanism. Step 1 B is produced. Reverse of
Step 1 B is consumed. Step 2 B is consumed.
Slide 32
The SSA (Contd) The SSA applied to the intermediate B.
Slide 33
SSA The Final Step Substitute the expression for the
concentration of B into the rate law v p.
Slide 34
Thermodynamic Formulation of Transition State Theory Activated
complex theory Reactant molecules proceed through transition state
TS falls apart unimolecularly to form products
Slide 35
The Equilibrium Constant Assume an equilibrium between the
activated complex and the reactants
Slide 36
The Gibbs Energy of Activation
Slide 37
Relating E a and H For a bimolecular, gas-phase reaction For
unimolecular, or solution-phase reactions
Slide 38
Comparing The Two k 2 Expressions
Slide 39
Know your ABCs For all solution-phase reactions For gas-phase
reactions