Activation energy
Review of Exothermic
• Reactants Ep is higher
than Products Ep.
• Now, we must consider
the activation energy
(the energy needed so
that the reactants bonds
will break and reform to
make product)
Review of Endothermic
• Reactants Ep is lower
than Products Ep.
• Need to add more energy
to the system for the
forward reaction to take
place.
• Still need to consider
activation energy
Activated complex is an unstable chemical species
containing partially formed bonds representing the maximum
potential energy point in the change
-also known as the transition state
Activated Complex
• Is the short-lived, unstable structure formed
during a successful collision between
reactant particles.
• Old bonds of the reactants are in the process
of breaking, and new products are forming
• Ea is the minimum energy required for the
activation complex to form and for a
successful reaction to occur.
Fast and slow reactions
• The smaller the activation energy, the faster
the reaction will occur regardless if
exothermic or endothermic.
• If there is a large activation energy needed,
that means that more energy (and therefore,
time) is being used up for the successful
collisions to take place.
Reaction Mechanism
Reaction Mechanisms
The overall progress of a chemical reaction can be represented
at the molecular level by a series of simple elementary steps
or elementary reactions.
The sequence of elementary steps that leads to product
formation is the reaction mechanism.
2NO (g) + O2 (g)
2NO2 (g)
N2O2 is detected during the reaction!
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
Reaction Intermediates
Intermediates are species that appear in a reaction
mechanism but not in the overall balanced equation.
An intermediate is always formed in an early elementary step
and consumed in a later elementary step.
Elementary step:
NO + NO
N 2O 2
+ Elementary step:
N2O2 + O2
2NO2
Overall reaction:
2NO + O2
2NO2
Rate Laws and Rate Determining Steps
Writing plausible reaction mechanisms:
•
The sum of the elementary steps must give the overall
balanced equation for the reaction.
•
The rate-determining step should predict the same rate
law that is determined experimentally.
The rate-determining step is the
slowest step in the sequence of
steps leading to product formation.
Molecularity
• The number of molecules that participate as
reactants in an elementary step
• Unimolecular: a single molecule is
involved.
– Ex: CH3NC (can be rearranged)
– Radioactive decay
– Its rate law is 1st order with respect to that
reactant
• Bimolecular: Involves the collision of
two molecules (that form a transition
state that can not be isolated)
– Ex: NO + O3  NO2 + O2
– It’s rate law is 1st order with respect to each
reactant and therefore is 2nd order overall.
– Rate =k[NO][O3]
Termolecular: simultaneous collision of three
molecules. Far less probable.
• Some possible mechanisms for the reaction;
2A + B  C + D
Rate Laws and Elementary Steps
Unimolecular reaction
A
products
rate = k [A]
Bimolecular reaction
A+B
products
rate = k [A][B]
Bimolecular reaction
A+A
products
rate = k [A]2
A catalyst is a substance that increases the rate of a
chemical reaction without itself being consumed.
Ea
k
uncatalyzed
catalyzed
ratecatalyzed > rateuncatalyzed
Energy Diagrams
Exothermic
Endothermic
(a) Activation energy (Ea) for the forward reaction
50 kJ/mol
300 kJ/mol
(b) Activation energy (Ea) for the reverse reaction
150 kJ/mol
100 kJ/mol
(c) Delta H
-100 kJ/mol
+200 kJ/mol
The experimental rate law for the reaction between NO2
and CO to produce NO and CO2 is rate = k[NO2]2. The
reaction is believed to occur via two steps:
Step 1:
NO2 + NO2
NO + NO3
Step 2:
NO3 + CO
NO2 + CO2
What is the equation for the overall reaction?
NO2+ CO
NO + CO2
What is the intermediate? Catalyst?
NO3
NO2
What can you say about the relative rates of steps 1 and 2?
rate = k[NO2]2 is the rate law for step 1 so
step 1 must be slower than step 2
Write the rate law for this reaction.
Rate = k [HBr] [O2]
List all intermediates in this reaction.
HOOBr, HOBr
List all catalysts in this reaction.
None
Reaction mechanisms are only educated guesses at the
behaviour of molecules, but there are three rules that
must be followed in proposing a mechanisms:
1. each step must be elementary
2. The slowest step must be consistent
with the rate equation
3. The elementary steps must add up to
the overall equation
Temperature Dependence of the Rate Constant
k = A e -Ea/RT
(Arrhenius equation)
Ea is the activation energy (J/mol)
R is the gas constant (8.314 J/K•mol)
T is the absolute temperature
A is the frequency factor
-Ea 1
ln k = + lnA
R T
Note: Changes to both the activation energy and temperature
have exponential effects on the value k, hence rate of reaction
• The Arrhenius equation show the effect of temperature on
the rate constant, k
• It indicates that k depends exponentially on temperature
Arrhenius equation:
k = A e-Ea/RT
Ea – activation energy
R – gas constant, 8.3145 J mol-1K-1
T - Kelvin temperature
A – Arrhenius constant (depends on collision rate and shape
of molecule)
k=A
-Ea/RT
e
• As T increases, the negative exponent
becomes smaller, so that value of k
becomes larger, which means that the rate
increases.
Higher T  Larger k  Increased rate
ln k and 1/T is linear
• With R known, we can find Ea graphically
from a series of k values at different
temperatures.
• ln k2 = - Ea ( 1/T2 – 1/T1)
k1
R
• Ea = -R (ln k2 / k1) ( 1/T2 – 1/T1) -1
Problem
• The decomposition of hydrogen iodide has
rate constants of 9.51 x10-9 L/mol.s at 500.0
K and 1.10 x 10-5 L/mol.s at 600.0 K. Find
Ea.
Ea = - (8.314)( ln 1.10 x10-5/ 9.51 x109)(1/600.00 – 1/500.0)
= 1.76 x 105 J/mol
• If rearrange the equation and convert it to:
lnk = - Ea . 1 + lnA
R T
• A graph of ln k against 1/T will be linear with
a slope/gradient of –Ea/R and an intercept on
the y-axis of lnA
y
lnk = - Ea . 1 + lnA
R T
= m . x +b
Plot lnk vs. 1/T = straight line
This is known as an Arrhenius plot