Reaction Rates
Chapter 17
Honors Chemistry
Reaction Rates
Red  Blue
Reactions can be…

FAST!
Liquid hydrogen and oxygen
reacting to launch a shuttle
..or

S
L
O
W
Concrete
hardening
..or

S
L
O
W
Watching paint dry
Expressing rxn rates in quantitative terms
quantity
Average rate 
t
Example: Reaction data for the reaction between butyl chloride
(C4H9Cl) and water is given below. Calculate the average reaction
rate over this time period expressed as moles of C4H9Cl consumed
per liter per second.
quantity
Average rate 
t

[C4H9Cl]final - [C4H9Cl]initial
tfinal  tinitial
mol
(0.100 mol
0.220
L
L )

(4.00 s - 0.00 s)
Table 17-1: Molar Concentration
[C4H9Cl] at
t=0.00 s
[C4H9Cl] at
t=4.00 s
0.220 M
0.100 M
-0.120 mol
L

4.00 s
 -0.0300
mol
L s
Collision Theory



Atoms, ions and molecules must
collide in order to react.
Reacting substances must collide with
the correct orientation.
Reacting substances must collide with
sufficient energy to form the activated
complex.
Orientation and the activated complex

Analogy: if you start with
two separate paperclips
(reactants) and you wish
to link them together
(products), not only must
the paperclips come into
contact, but they also
must collide with a
specific orientation.
Orientation and the activated complex

Biological example:
ENZYMES

You’ll learn much more
about this next year
Activation energy and reaction

Only collisions with enough energy to
react form products
Activation energy and reaction
Activation energy and reaction
Another example
reactants
Activated complex (also
called transition state)
products
Does ∆G tell us anything about rxn rate?
 NO


If a reaction is spontaneous, it does not
follow that it is fast or slow.
Thus, a new branch of chemistry…
kinetics
Factors affecting reaction rates
1)
2)
3)
4)
5)
6)
The nature of the reactants
Concentration
Pressure (gases only)
Surface area
Temperature
Catalysts
NATURE OF REACTANTS

Some elements/compounds are more
reactive than others
FAST
sodium in water (alkali metals are VERY reactive)
NATURE OF REACTANTS

Some elements/compounds are more
reactive than others
SLOW
Rusting of iron (it takes time for moisture in the air to oxidize the metal… process
can be sped up if salt is present, but will still not react as fast as sodium and water)
CONCENTRATION

As concentration ↑, frequency of
collisions ↑, and therefore rxn rate ↑
PRESSURE (gases)

For gases,
increasing
pressure creates
the same effect
as increasing
concentration
SURFACE AREA

As surface area ↑, rxn rate ↑
← slow
fast
TEMPERATURE:

Generally, ↑ temp = ↑ rate


Why?
Higher temp = faster molecular motion
= more collisions and more energy
Analogy: imagine that you
are baby-sitting a bunch of
per collision = faster rxn
6 year olds. You put them
in a yard and you let them
run around. Every now and
then a couple of kids will
run into each other. Now
imagine that you decide to
feed them some sugar.
What happens? They run
around faster and of course
there are many more
collisions. Not only that, the
collisions are likely to be a
lot harder/more intense.
CATALYST

Catalyst: a substance that speeds up the rate
of a reaction without being consumed in the
reaction.



Provides an easier
way to react
Lowers the
activation energy
Enzyme =
biological catalyst
CATALYST

Adding a catalyst speeds up the rxn by
lowering the activation energy
Reaction Rate Laws

The equation that expresses the mathematical
relationship between the rate of a chemical
reaction and the concentration of reactants is
a rate law.
For the reaction : A  B
Rate  k[A]
*where k is a constant
specific to this reaction
Reaction Order

The reaction order for a reactant
defines how the rate is affected by the
concentration of that reactant.

The overall reaction order of a chemical
reaction is the sum of the orders for the
individual reactants in the rate law.
Reaction Order


The rate law for most reactions has the
general form….
Rate = k[reactant 1]m[reactant 2]n…
The exponents m and n are called
reaction orders. Their sum (m + n) is
called the overall reaction order.
Reaction Order

For the reaction aA + bB products
Rate = k[A]m[B]n
where m and n are the reaction orders for A and B respectively


Only if the rxn between A and B happens in a
single step (with a single activated complex…
which is unlikely) does m=a and n=b.
Thus, the values of m and n must be
determined experimentally!!!
Reaction Order

For the reaction aA + bB products
Rate = k[A]m[B]n
where m and n are the reaction orders for A and B respectively

Rate laws cannot be predicted by looking at a
balanced chemical equation.
Finding the rate law


The most common method for
experimentally determining the differential
rate law is the method of initial rates.
In this method several experiments are run
at different initial concentrations and the
instantaneous rates are determined for each
at the same value of time (as near t = 0 as
possible)
Using Initial Rates to Determine
the Form of the Rate Law
A + B  C
Exp #
[A]
[B]
Initial Rate (M/s)
1
.100M
.100M
4x10-5
2
.100M
.200M
4x10-5
3
.200M
.100M
16x10-5
From this data, find the form of the rate law..
Rate = k[A]m[B]n
Exp #
[A]
[B]
Initial Rate (M/s)
1
.100M
.100M
4x10-5
2
.100M
.200M
4x10-5
3
.200M
.100M
16x10-5
Rate = k[A]m[B]n
Rate 2 4  10 5 k [.100]m [.200]n



5
m
Rate 1 4  10
k [.100] [.100]n
[.200]n
1
[.100]n
12
n
n=0
Exp #
[A]
[B]
Initial Rate (M/s)
1
.100M
.100M
4x10-5
2
.100M
.200M
4x10-5
3
.200M
.100M
16x10-5
Rate = k[A]m[B]n
Rate 3 16  10 5 k [.200]m [.100]n



5
m
Rate 1
4  10
k [.100] [.100]n
[.200]m
4
m
[.100]
4  2m
m=2
Exp #
[A]
[B]
1
.100M
.100M
4x10-5
2
.100M
.200M
4x10-5
3
.200M
.100M
16x10-5
Rate = k[A]m[B]n
Rate = k[A]2[B]0
Rate = k [A]2
Initial Rate (M/s)
Exp #
[A]
[B]
1
.100M
.100M
4x10-5
2
.100M
.200M
4x10-5
3
.200M
.100M
16x10-5
Now, solve for k…
rate
k 
[A]2 [B ]0
Initial Rate (M/s)
Rate = k [A]2
4  105
3


4

10
[.100]2
Rate = 4x10-3 [A]2
Knowing rate laws and rate orders helps us
predict how the reaction will proceed over time

Application:



Radioactive decay
is a first order
reaction
Half life is constant
over time
Allows us to date
fossils, etc.
Reaction Mechanisms



Most chemical reactions consist of a
sequence of two or more steps (or
simpler reactions). These add together
to create the overall reaction equation.
Generally, some steps will be fast and
others will be slow.
The slow step is the rate determining
step.
Example:
The reaction between NO2 and CO has the
overall reaction:
NO2 + CO  NO + CO2
A study of the kinetics of this reaction revealed the
rate law for the reaction is.
Rate = k[NO2]2
This Rate Law requires that the slow step of
the reaction involves a collision between two NO2
molecules. How can this be a step in the seemingly
simple reaction above?
NO2 + CO  NO + CO2
Further study of this reaction established that
two NO2 molecules can react as follows…
NO2 + NO2  NO3 + NO
NO3 is a highly reactive material which is capable
of transferring an oxygen atom.
NO3 + CO  NO2 + CO2
2NO2 + NO3 + CO  NO2 + NO3 + NO + CO2
NO2 + CO  NO + CO2
NO2 + CO  NO + CO2
The first equations sets the rate law, so it must be
the slow one.
NO2 + NO2  NO3 + NO
slow
NO3 + CO  NO2 + CO2
fast
NO2 + CO  NO + CO2
Remember, the rate law only provides information
about the slowest reaction in the mechanism.