AH Chemistry – Unit 1
Kinetics
Kinetics
How fast does it go?
Thermodynamics
Is the reaction feasible?
How far will the reaction go?
Thermodynamics is about start and
finish
Kinetics is about what happens along
the way.
Kinetics
• Studies the rate at which a chemical process
occurs.
• Besides information about the speed at which
reactions occur, kinetics also sheds light on
the reaction mechanism (exactly how the
reaction occurs).
Measuring reaction rates
Properties suitable for measuring rate
•The mass of the apparatus
(when a gas is being released)
•The volume of a gaseous product
•The pH of a solution
(when H+ or OH- ions are used or produced)
•The conductivity of a solution
(when the number or nature of the ions present changes)
•The colour
(when products and reactants have different colours)
Measuring reaction rates
Concentration of
reactants
mol l-1
Time/s
Measuring reaction rates
Concentration of
products
mol l-1
Time/s
Reaction Rate
We will study 2 examples – one on the PPT and one in your notes
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
• 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.
Reaction Rates
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.
Reaction Rates and Stoichiometry
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
• 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 C4H9OH.
Rate =
-[C4H9Cl]
=
t
[C4H9OH]
t
Reaction Rates and Stoichiometry
• What if the ratio is not 1:1?
H2(g) + I2(g)  2 HI(g)
• Only 1/2 HI is made for each H2 used.
Reaction Rates and Stoichiometry
• To generalize, for the reaction
aA + bB
Reactants (decrease)
cC + dD
Products (increase)
The Rate Law
aA
+
bB
→
cC
+
dD
Rate 
m
n
Rate = k[A] .[B]
m
n
[A] .[B]
Where…
k = rate constant
m and n = orders of reaction with respect to A and B
Rate Laws
• A rate law shows the relationship between the reaction rate and
the concentrations of reactants.
– For gas-phase reactants use PA instead of [A].
• k is a constant that has a specific value for each reaction.
• The value of k is determined experimentally.
“Constant” is relative herek is unique for each reaction
k changes with temperature.
order of a reaction
The power to which the concentration of a particular
reactant is raised in the rate equation
overall order of a reaction
The sum of the powers to which the concentrations of all
reactants in the rate equation are raised in the rate
equation
Rate equation
0
Order of reaction
0
1
1
2
2
Rate = k [A]
Rate = k [A]
Rate = k [A]
1
Rate = k [A]
1
Rate = k [A]
1
2
2
3
[B]
[B]
Order for the reaction is the power to which the conc. of that
reactant is raised in the rate equation
Units for k depend on the overall rate of the reaction.
Small integral values refer to the number of molecules in the
single step of the reaction which controls the overall reaction rate
– the RATE DETERMINING STEP.
These do not necessarily refer to the
number of molecules in the balanced
equation.
Order cannot be deduced from the balanced equation.
It can only be determined experimentally
H2O2 + 2HI  H2O + I 2
Rate equation (from experiment)
Rate
[H2O2][HI]
Overall order of reaction = 2
NOTE: Balanced equation has 3 reactant molecules
Order of reaction has been experimentally determined
ZERO ORDER
Changing the concentration has no effect on the rate of the
reaction.
0
Rate = k [reactant]  rate = k
FIRST ORDER
Doubling the concentration doubles the rate.
Tripling the concentration triples the rate.
1
Rate = k [reactant]  rate = k [R]
SECOND ORDER
Doubling the concentration quadruples (4x) the rate.
2
Rate = k [reactant]  rate = k [R]
2
k – the value for k is found experimentally
Its value depends on the experimental conditions e.g.
varies with temperature
A + B + C  D (g)
Determining Order
Expt. no
[A]
[B]
[C]
(mol l-1 s-1)
Initial rate
(ml of D s-1)
1
1.0
1.0
1.0
20
2
2.0
1.0
1.0
40
3
1.0
2.0
1.0
20
4
1.0
1.0
2.0
80
Rate Law/equation  Rate = k [A] [B]0 [C]2
Rate Law/equation  Rate = k [A] [C]2
Orders of reaction
• Reactions usually occur in a series of steps.
• The integers in the rate equation refer to the
actual number of particles involved in the
slowest step of the reaction.
• This is the rate determining step.
Calculating the Rate Constant
• Can be calculated from a series of
experiments in which the starting
concentrations of reactants are varied…
Example
Reaction rates and
reaction mechanisms