Chapter 13
Exothermic release energy into the surroundings
temp of surroundings increases
products have less energy - ΔE
R

P
Endothermic absorb energy from the surroundings
temp of surroundings decreases
products have more energy +Δ E
P

R
• A + B ----> C + D + 30 joules
• A + B ----> C + D ΔE = -30 joules
• Forward reaction is exothermic
• Reverse reaction is endothermic
• 30 joules + C + D ----> A + B
• C + D ----> A + B ΔE = + 30 joules
Reactions involve
• Bond breaking –requires energy +ΔE
• Bond formation –releases energy -ΔE
• If more energy is needed to break bonds than is
released
+ΔE ( endothermic)
• If more energy is released in bond formation
than is absorbed -ΔE (exothermic)
• Reaction mechanism- steps by which a
reaction occurs
• Each step of a reaction mechanism involves
the collision of two molecules.
Colliding molecules need
• 1. enough energy to break the bonds
• Activation Energy ( A.E) or threshold energy
• 2. the proper geometry or orientation (correct
angle)
Energy
(P.E.)
reaction coordinate
c
Factors affecting rate of reaction
•
•
•
•
Nature of reactants (number of bonds)
Surface area (solids and liquids)
Temperature
Concentration of reactants (solutions and
gases)
• Catalyst
• Rate =(# of coll/time) (fract with A.E.) (fract
with orientation)
• If orientation factor is 1 (orientation does not
matter)
• TEMPERATURE
• affects fraction with A.E.
Temp - measure of the average K.E.
At a given temp all molecules do not have the same K.E.
# of molecules
Kinetic energy
threshold energy (A.E)
• Rate =(# of coll/time) (fract with A.E.) (fract
with orientation)
• Concentration
• affects the # of collisions
At a given temp
Rate = k (# of collisions /time)
Rate = k (# of collisions /time)
• k large - fast rxn
• k small - slow rxn
• Concentration of reactants raised to some
power (order)
• Rate = (k ) [A]x [B]y
• X and Y (orders)
• Orders found only by experiment
Rate = (k ) [A]x [B]y
• If concentration of a reactant is doubled and
the rate doubles the order is 1
• If concentration of a reactant is doubled and
the rate quadruples the order is 2
• If changing concentration of a reactant has no
effect on the rate of reaction it is not included
in the rate law
• Sum of all orders is the order of the reaction
Catalysts and rate
Potential energy
Reaction coordinate
A catalyst does not change the K.E of the molecules
# of molecules
Kinetic energy

catalyst

no catalyst
• Catalyst -changes the pathway (steps) Steps
require less energy
• Changes the orientation requirement- more
molecules have required orientation
• Not consumed in the reaction
Reaction mechanism and Rate Law
•
•
•
•
Each step involves the collision of two molecules
–adding steps gives net reaction
Each step has its own rate law.
In the steps the coefficients are the orders of the
rate law
• Slowest step determines the overall rate -rate
determining step
• The rate law for the slowest step is the rate law
for the reaction.
Reaction Mechanism
• A + B- C + 2D
• D+ B  DB
• DB + D  F
• A + 2B - C + F
Determining Reaction Mechanism
• 1. Do an experiment to determine the rate
law. Experimental rate law
• 2. Postulate possible steps by which the
reaction could take place. Reaction
Mechanism
A2+ B2 
•
•
•
•
•
•
2AB
rate = k [A2]2 (from experiment)
Possible mechanism 1
A2+ B2  2AB
rate = k [ A2] [ B2]
Predicted does not match experimental
Not the reaction mechanism
A2+ B2 -> 2AB
•
•
•
•
•
rate = k [A2]2 (from experiment)
Possible mechanism 2
Step 1
A 2 + A2  2A + A2
Step 2
A + B2  AB2
Step 3
AB2 + A  2AB
•
•
•
•
Step 1 rate = k [A2][A2] or rate =k [A2]2
Step 2 rate = k [A][B2]
Step 3 rate = k [AB2] [A]
Do any rate laws match the experiment rate law?
rate = k [ A2]2
2NO + O2 2NO2
Experiment
Initial
concentration
[NO]
Initial
concentration
[O2]
Rate of formation
of NO2 (M/s)
1
0 .015 M
0 .015 M
0.048
2
0 .030 M
0 .015 M
0.192
3
0 .015 M
0 .030 M
0.096
4
0 .030 M
0 .030 M
0.384