Kinetics = The study of the rate at
which a reaction occurs.
KINETICS
For a chemical reaction to occur,
molecules of the reactants must
come together so that atoms can
be exchanged or rearranged.
Atoms and molecules are mobile
in the gas phase or in solution,
and so reactions are often carried
out using a mixture of gases or
using solutions of reactants.
Reaction Rates
• Three “types” of rates:
initial rate
average rate
instantaneous rate
Reaction Rate = The change in
concentration of a reactant or
product per unit time
Therefore, when studying reaction
rates we most commonly observe changes
in:
1.
Concentration
(for solutions)
2.
Pressure
(for gaseous)
3. Absorbance
(for solutions)
Average Rate vs.
Instantaneous Rate
Average rate is the change in
concentration, [ ], of a reactant or
product per unit time (M/s).
For example, a trip to Wal-mart had
an average rate of travel equal to 35
mi/hr.
2 NO2  2 NO + O2
Instantaneous rate, referred to
simply as rate, is the rate of a
reaction at some particular time
(M/s).
-Δ[NO2]
+ Δ[NO]
+ Δ[O2]
For example, when getting pulled
over, the police officer measured
your instantaneous speed to
determine if you should get a
ticket
Observe:
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
•
In this
reaction, the
concentration of
butyl chloride,
C4H9Cl, was
measured at
various times.
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
•
The average rate of the reaction over each
interval is the change in concentration divided
by the change in time:
[C H Cl]
Average rate =
4
t
9
1. Using the data in table 14.1 and figure 14.4 on
pg 561, calculate the average rate for the
reaction of C4H9Cl with water from t = 200.0 s
to t = 500.0 s.
• Note that the average rate decreases as the
reaction proceeds.
• This is because as the reaction goes forward,
there are fewer collisions between reactant
molecules.
The average rate of a chemical reaction
can easily be determined by measuring the
concentration of either reactants or products at
two times and dividing by the total elapsed time
of reaction using the following formula as we
have seen in the above problem:
Average rate =
2 NO2  2 NO + O2
- Δ[NO2]
+ Δ[NO]
+ Δ[O2]
[C4H9Cl]
t
Therefore,
Δ[NO2] is
negative
However, to determine the rate of the
reaction at some time, t, we must analyze the
graph of concentration vs. time.
A line drawn tangent to the curve of
concentration, [X], vs. time at any given
time, t, has the equation:
y = mx + b
[X] = mt + b
Where y is the concentration of the
reactant or product species, x is some time
t, b is the y intercept and m is the slope of
the line, describing the rate at some time t.
m = Δy / Δ x = rate = Δ[ ] / Δ t
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
• All reactions slow
down over time.
• Therefore, the
best indicator of
the rate of a
reaction is the
instantaneous rate
near the
beginning.
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
• A plot of
concentration vs.
time for this
reaction yields a
curve like this.
• The slope of a line
tangent to the curve
at any point is the
instantaneous rate
at that time.
2. Using the data
in table 14.1
and figure 14.4
on pg 561,
estimate the
rate of the
reaction at 600
sec.
Reaction Rates and Stoichiometry
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
For any chemical reaction, the rate
of disappearance of reactants is directly
related to the appearance of products,
described by the reaction coefficients.
• In this reaction,
the ratio of
C4H9Cl to
C4H9OH is 1:1.
• Thus, the rate of
disappearance of
C4H9Cl is the
same as the rate of
appearance of
Rate =
C4H9OH.
• What if the ratio is not 1:1?
Example:
-[C4H9Cl]
=
t
[C4H9OH]
t
• To generalize, then, for the reaction
aA + bB
cC + dD
2 HI(g)  H2(g) + I2(g)
•The coefficients can be described as
ratios (fractions) of the reaction rate:
Rate = −
1 [A]
1 [B]
1 [C]
1 [D]
=
=
=−
a t
b t
c t
d t
Rate = − 1 [HI] = [I2]
2 t
t
For the Reaction:
2 NO2  2 NO + O2
Describe:
Δ[NO2]
Δ[NO]
Δ[O2]
In our first example, for every 1 NO2
molecule reacted, 1 NO and ½ O2 molecules
are produced
Because only ½ mole O2 is produced,
twice the amount of NO2 and NO are
reacted and produced respectively
Or,
Rate  
1 [ NO2 ] 1 [ NO ] 1 [O2 ]


2 t
2 t
1 t
Notice, negative NO2; why?
3. If the rate of decomposition of
N2O5 at a particular instant in a
reaction vessel is 4.2 x 10-7 M/s,
what is the rate of appearance of
a) NO2?
b) O2?
c) What is the Rate of the reaction?
Nature of Reactants
 Some compounds are more reactive
than others
Factors Affecting Rates
1. nature of reactants
2. concentration
3. temperature
4. surface area
5. presence of a catalyst
Concentration of
Reactants
Rate with 0.3 M HCl
 As the concentration
of reactants
increases, so des the
likelihood that
reactant molecules
will collide.
Rate with 6.0 M HCl
Temperature
 At higher temperatures,
reactant molecules have
more kinetic energy,
move faster, and collide
more often and with
greater energy.
Physical State of the Reactants
 In order to react, molecules must come
in contact with each other.
 The more homogeneous the mixture of
reactants, the faster the molecules can
react.
Gas
vs.
Gas Vapors
demo
Presence of a Catalyst
 Catalysts speed up reactions
by changing the mechanism
of the reaction.
 Catalysts are not consumed
during the course of the
reaction.