1 Chemistry 2C Lecture 22: May 21 th , 2010 1) Arrhenius Equation 2) Transition State Theory 3) Molecularity 4) Rate limiting steps 5) Reaction mechanisms 6) Catalysis 7) Nuclear Introduction Lecture 22: Kinetics
1Chemistry 2C Lecture 22: May 21th, 2010
1) Arrhenius Equation2) Transition State Theory
3) Molecularity4) Rate limiting steps
5) Reaction mechanisms6) Catalysis
7) Nuclear Introduction
Lecture 22: Kinetics
2Chemistry 2C Lecture 22: May 21th, 2010
Temperature Effects on Reaction Rate
“Normal” Catalyzed reaction(inactivation of catalyst)
rare Chain reaction(Explosion)
Not every reaction follows the Arrhenius equation. But most!
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Activated Complex-Transition States
The potential energy of the system increases at the transition state because: 1) The approaching reactant molecules must overcome the mutual repulsive
forces between the outer shell electrons of their constituent atoms 2) Atoms must be separated from each other as bonds are broken
In the transition state theory, the mechanism of interaction of reactants is not considered; the important criterion is that colliding molecules must have
sufficient energy to overcome a potential energy barrier (the activation energy) to react.
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Activated Complex-Transition StatesThis increase in potential energy corresponds to an energy barrier
over which the reactant molecules must pass if the reaction is to
proceed. The transition state occurs at the maximum of this energy
barrier.
The transition state is an unstable transitory combination
of reactant molecules that occurs at a potential energy maximum
The combination can either go on to form products or fall apart
to return to the unchanged reactants.
The energy difference between the reactants and the potential energy maximum is referred to as the activation energy: Ea or
G‡
5Chemistry 2C Lecture 22: May 21th, 2010
Energy barrier from reactants to
products (forward direction)
Reaction Profile
Energy barrier from products to reactants
(reverse direction)
G for reaction
Since G is negative, this is a spontaneous reaction, although its timescale for occurring is dictated by the energy barrier
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Arrhenius EquationsHow is Ea measured?
The rate constant is a function of temperature, but Ea is considered to be a constant and depends only on thermodynamics
k=Ae-Ea/RT
This is the form of a line!
RT
EAk
Aek
A
RT
EA
lnln
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Arrhenius EquationsHow is Ea measured (measure two rates at two different temperatures?
ln k2 = –Ea/RT2 + ln A
@ T1
@ T2
ln k1 = –Ea/RT1 + ln A
Subtract the two equations:
ln k1 - ln k2 = –Ea/RT1 –Ea/RT2
ln (k1 / k2) = –Ea/R (1/T1 -1/T2)
Can just substitute into this equation! Make sure temperature is in Kelvin!
8Chemistry 2C Lecture 22: May 21th, 2010
Chemical Kinetics Molecularity of a Reaction
The reaction order refers to the concentration dependence of the reaction rate and can be an integer or a non-integer and even negative!
The molecularity of a reaction refers to a definite molecular encounter during the course of the reaction. The molecularity has to be an integer
(there are no partial atoms/molecules!)
Unimolecular Reactions: One reactant molecule undergoes transformation into the product(s). Examples are racemizations, thermal
decomposition, or isomerizations.
Bimolecular Reactions: Two reactant molecules collide in one elementary step. Most common type of reaction molecularity.
Termolecular reactions: Three reactants have to collide to lead to an reaction. Extremely rare.
No reactions of molecularities higher than 3 are known.
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Reaction MechanismA mechanism describes in detail exactly what takes place at each stage of a chemical transformation. It describes the transition state and which bonds
are broken and in what order, which bonds are formed and in what order, and what the relative rates of the steps are. A complete mechanism must also
account for all reactants used, the function of a catalyst, stereochemistry, all products formed and the amount each.
For example:
CO + NO2 → CO2 + NO
In this reaction, it has been experimentally determined that this reaction takes place according to the rate law R = k[NO2]2. Therefore, a possible
mechanism by which this reaction takes place is:
2 NO2 → NO3 + NO (slow)
NO3 + CO → NO2 + CO2 (fast)
Each step is called an elementary step, and each has its own rate law and molecularity. All of the elementary steps must add up to the original reaction.
10Chemistry 2C Lecture 22: May 21th, 2010
Intermediates & Rate Determining StepsReactions may have intermediate species that don’t show up in the final
reaction equation, but play a role in the mechanism. Kindah like electrons in REDOX reactions.
For example: 2 NO2 (g) + F2 (g) → 2NO2F (g)
The first step is slow (k1<<1) and the second step is fast (k2>>1), the rate determining step
NO2 + F2 → 2NO2F + F
Overall Rx.
Step one: k1
F + NO2 → NO2F Step two: k2
rate = k[NO2][F2]
The experimentally determined rate law is:
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NO2 + F2 NO2F + F
F + NO2 NO2F
k1
k2
Slow
Fast
2NO2 + F2 2NO2F
rate = k[NO2][F2] according to the elementary reaction, with a rate limiting step
Intermediates & Rate Determining Steps
Each of these steps is an elementary process. That means that those two species must collide for a reaction to occur. In this
example, each step is bimolecular.
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Energy Diagram for a two-step Reaction
Reactants -> transition state -> intermediateIntermediate -> transition state -> product
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Chemical Kinetics: Enzymes and Catalysis
General principles of catalysis:
• A catalyst works by lowering the Gibbs energy of activation. This enhances the rate of forward and backward reaction.
• The catalyst forms an intermediate with the reactant(s) in the initial step of the reaction and is released in during product formation.
• A catalyst can not affect the enthalpies or the Gibbs energies of the reactants and products.It increases the rate of the approach to equilibrium, but can not change the change the equilibrium constant.
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Chemical Kinetics: Enzymes and Catalysis
General principles of catalysis:
Uncatalyzed Reaction Catalyzed Reaction
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Types of Catalysis
• Homogeneous Catalysis– The interaction of reactants and catalysts in the
same phase.– e.g., CFC’s (gas/gas)
• Heterogeneous Catalysis– The interaction of reactants and catalysts in
different phases.– e.g., catalytic converters (solid/gas)
Enzymes in the body are biological catalysts!
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Reaction Profile for Enzyme
The catalysed reaction pathway goes through the transition states TSc1, TSc2 and TSc3, with an energy barrier Gc*, whereas the uncatalysed reaction goes through the transition state TSu
with a barrier of Gu*. In this example the rate limiting step would be the conversion of ES into EP.
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Reaction Profiles
The catalysis of H2O2 decomposition by Br-
2H2O2 -> 2H20 + O2 Overall net Rx.
When Br- is added, the reaction goes
2Br- + H2O2 +2H+ -> Br2 + 2H20
Br2 + H2O2 -> 2Br- + 2H+ + O2
Step 1
Step 2
Br2 is an intermediate because it is produced and then consumed!
Br- is a catalyst because it speeds up the reaction and is neither and is unchanged during the net reaction!