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Reactants must collide with proper
orientation and sufficient energy
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Explains what happens once colliding
particles react
◦ Transition state is the in-between state when
reactants are being converted to products
◦ Kinetic energy is converted to potential energy
(think about bouncing a basketball)
Kinetic Energy
Potential Energy
Negative
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Exothermic reactions- give off heat (products
more stable, less potential energy)
Endothermic reactions require heat (products
less stable than reactants, higher PE)
Exothermic Reaction
Endothermic Reaction
Draw a PE diagram. Include: axes
labels, the transition state, activated
complex, Ea (forward and reverse),
and ΔE.
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CO reacts with NO2 to form CO2 and NO. The
activation energy of the forward reaction is
134 kJ and the ΔE is -226 kJ.
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Chemical reactions typically occur as a series
of steps. This series of steps that make up
the overall reaction is called the reaction
mechanism.
Elementary reactions are a single step in the
overall reaction mechanism.
◦ Singular molecular event, such as a simple collision
of atoms, molecules or ions.
◦ Cannot be broken down into further simpler steps.
For example, the reaction
2NO(g) + O2(g)  2NO2(g)
involves a two-step reaction mechanism:
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◦ Step 1: NO(g) + O2(g)  NO3(g)
◦ Step 2: NO3(g) + NO(g)  2NO2(g)
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Each step is an elementary reaction, both
steps together give the overall reaction
mechanism.
Notice NO3(g)
◦ Not a product or reactant of overall reaction
◦ Produced then consumed = reaction intermediate
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Describes the number of reactant particles in
an elementary step
◦ Unimolecular = one reactant
 (CH3)3CBr(aq)  (CH3)3C+ + Br-
◦ Bimolecular = 2 reactants come together. Ex:
 Step 1: NO(g) + O2(g)  NO3(g)
 Step 2: NO3(g) + NO(g)  2NO2(g)
◦ Termolecular = 3 reactants
 Extremely rare!!! Why?
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One step in a reaction mechanism is always
much slower than the others
◦ Since it is so much slower, it determines the rate of
the overall reaction.
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Hence, this slow step is called the rate
determining step
◦ Consider the process of making toast:
slow
fast
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Activation energy for rds is always higher
◦ 2 STEP REACTION MEANS 2 TRANSITION STATES
AND 1 INTERMEDIATE
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Increases the rate of a reaction
◦ Homogeneous catalyst: Same phase as reactants
◦ Heterogeneous catalyst: Different phase as
reactants
Pd
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Not consumed!
◦ There in beginning and end of reaction
◦ Works by lowering the activation energy
 Therefore, greater number of collisions have sufficient
energy to react
 Provide alternative reaction mechanism
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Reaction Rate
Concentration
Surface Area
Catalyst
Reactivity
Temperature
Collision Theory
Orientation
Rate Expression
Exothermic
Endothermic
Homogeneous
Heterogeneous
Kinetic Energy
Unimolecular
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Rate Constant
Elementary Reaction
Mechanism
Catalyst
Rate Law
Reaction Order
Molecularity
Activation Energy
Reverse Reaction
Potential Energy
Transition State
Activated Complex
Rate Determining
Step
Biomolecular
Termolecular