Transcript Slide 1

Topic 6. Kinetics
aka Reaction Rates
1
IB Topic 6: Kinetics
6.1: Rates of Reaction
6.1.1
Define the term rate of reaction
6.1.2
Describe suitable experimental
procedures for measuring rates of
reactions
6.1.3
Analyse data from rate
experiments
2
2
6.1.1
Define the term rate of reaction
Rate of Reaction
Some reactions occur rapidly
 Inflation of an airbag
 Explosion of nitrogen triiodide
Some reactions occur slowly
 Reaction of pigments in paintings with light & pollutants
 Tarnishing of silver/Iron rusting
3
3
6.1.1
Define the term rate of reaction
Rate of Reaction
Use this definition:
An increase in concentration of one of the products per
unit time
OR
A decrease in concentration of one of the reactants per
unit time
How fast reactants are being converted to products during a
chemical reaction –

Either the rate of formation of a product or the rate of consumption of a
reactant may be used (they are related to each other by dividing the
corresponding coeffficient in the stoichiometric equation)
4
4
6.1.1
Define the term rate of reaction
Units of rate of reaction:
moles per dm3 per second
We will use mol dm-3 s-1
5
5
6.1.1 Define the term rate of reaction
Rates of reactions can be determined by
monitoring the change in concentration of
either reactants or products as a function
of time. [A] vs t or [B] vs t
6
6
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions
How fast does a reaction happen? How would we measure that?
 We need an indicator that a reaction is happening
 Color change
 Gas evolution
 Precipitate formation
 Heat and light
 In many cases its easiest to measure how fast the product is
appearing
 Example: MgCO3(s) → MgO(s) +CO2 (g)



The easiest thing to measure here would probably be the CO2
We could collect the CO2 in a syringe and measure how the
volume changed with time. The rate would be:
 ∆ Volume/time
We could also measure how the reactant is disappearing
7
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions
Measuring the Rate of Reaction








There are many ways to measure rate:
∆ Mass / time
∆ Volume / time
∆ Concentration / time
∆ Temperature / time
∆ pH / time
∆ Color / time (Use of a spectrophotometer or colorimeter)
∆ Conductivity / time
8
8
6.1.2 Describe suitable experimental
procedures for measuring rates of reactions




Rate = Δ[P] = -Δ[R]
ΔT
ΔT
[ ] = molar concentration mol dm-3
Δ[ ] = concentration @ T2 –
concentration @ T1
#rate value will vary according to
the number of moles of the
substance (coefficients in a
balanced equation)
9
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data
from rate experiments
Measuring the Rate of Reaction

The graph shows the
volume of CO2
produced against time
when excess CaCO3 is
added to HCl
Measuring the Rate of Reaction

The graph shows the
concentration of a
reactant against time
10
10
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from
rate experiments
State how the rate of formation of carbon dioxide
changes with time. Explain your answer in terms of
collision theory
Volume of CO 2
Time
11
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from
rate experiments
Time: 0 s As the reaction begins, the reactant concentration is
high, there is no product
Time: 5 s Reactant is being used up at about the same rate as
the product is being made
Time: 10 s The rates are slowing down (slopes are decreasing)
Time: 60 s No net change in concentration. Reaction has
finished.
Did all of the reactant react??
Concentration
Red: Reactant
Blue: Product
0s
10 s
20 s
30 s
40s
50s
60 s
12
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from
rate experiments
So… What is happening to the reaction rate of the
reaction shown below as time progresses?
Concentration
Red: Reactant
Blue: Product
0s
10 s
20 s
30 s
40s
50s
60 s
13
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from rate
experiments
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
[C4H9Cl] M
In this reaction,
the concentration
of butyl chloride,
C4H9Cl, was
measured at
various times, t.
14
1
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from rate
experiments
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
Average Rate, M/s
The average rate of
the reaction over
each interval is the
change in
concentration divided
by the change in
time:
15
1
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from rate
experiments
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.
16
1
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from rate
experiments
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.
17
1
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from
rate experiments
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)

The reaction slows
down with time
because the
concentration of the
reactants
decreases.
1818
6.1.2 Describe suitable experimental procedures for
measuring rates of reactions and 6.1.3 Analyse data from
rate experiments
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)


Rate =
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.
-[C4H9Cl]
=
t
[C4H9OH]
t
1919
hydrogen produced (cm3)
How can the rate of reaction be calculated from a graph?
70
60
x
50
rate of reaction = y
x
40
30
y
20
10
0
0
10
20
30
40
50
time (seconds)
The gradient of the graph is equal to the initial rate of
reaction at that time
rate of reaction = 45 cm3 rate of reaction = 2.25 cm3/s
20
20 s
Slower and slower!
Reactions do not proceed at a steady rate. They start off at a
certain speed, then get slower and slower until they stop.
As the reaction progresses, the concentration of reactants
decreases.
This reduces the frequency of collisions between particles
and so the reaction slows down.
0%
25%
reactants
product
50%
75%
100%
percentage completion of reaction
21
6.1.3
Analyse data from rate experiments
Consider the hypothetical aqueous reaction
A(aq)  B(aq)
A 100.0 mL flask initially contains 0.065 mol of A and
0 mol B. The following data are collected:
Time (min)
0
10
20
30
40
Moles of A
0.065
0.051
0.042
0.036
0.031
a) Calculate the number of moles of B at each time in the table.
b) Graph the change in amount of A & B over time.
c) Calculate the average rate of disappearance of A for each 10-minute
interval in terms of mol s-1
d) Between t=10 min & t=30 min, what is the average rate of
appearance of B in units of mol dm-3 s-1
22
22
6.1.3
Analyse data from rate experiments
a) Calculate the number of moles of B at each time in the table.
Time (min)
0
10
20
30
40
Moles of A
0.065
0.051
0.042
0.036
0.031
Moles of B
0.000
0.014
0.023
0.029
0.034
23
23
6.1.3
Analyse data from rate experiments
b) Graph the change in amount of A & B over time.
Reaction Rate A --> B
0.07
0.06
Moles
0.05
0.04
A
0.03
B
0.02
0.01
0
0
10
20
30
40
50
Time (min)
24
24
6.1.3
Analyse data from rate experiments.
c) Calculate the average rate of disappearance of A for
each 10-minute interval in terms of mol s-1
0
10
20
30
Time (min)
0
10
20
30
40
Moles of A
0.065
0.051
0.042
0.036
0.031
min10 min: (.051-.065)/600 = 2.3 x 10-5 mol s-1
min20 min: (.042-.051)/600 = 1.5 x 10-5 mol s-1
min30 min: (.036-.042)/600 = 1.0 x 10-5 mol s-1
min40 min: (.031-.036)/600 = 0.8 x 10-5 mol s-1
25
25
6.1.3
Analyse data from rate experiments
d) Between t=10 min & t=30 min, what is the average
rate of appearance of B in units of mol dm-3 s-1
Time (min)
0
10
20
30
40
Moles of B
0.000
0.014
0.023
0.029
0.034
[B] at 10 min = 0.014 mol/.100 dm3 = .14 M
[B] at 30 min = 0.029 mol/.100 dm3 = .29 M
Rate = (.29-.14)/1200 = 1.3 x 10-4 mol dm-3 s-1
26
26
IB Topic 6: Kinetics
6.2: Collision Theory
6.2.1 Describe the kinetic theory in terms of the movement of
particles whose average energy is proportional to temperature in
Kelvins.
6.2.2 Define the term activation energy, Ea.
6.2.3 Describe the collision theory.
6.2.4 Predict and explain, using the collision theory, the
qualitative effects of particle size, temperature, concentration
and pressure on the rate of a reaction.
6.2.5 Sketch and explain qualitatively the Maxwell-Boltzman
energy distribution curve for a fixed amount of gas at different
temperatures and its consequences for changes in reaction rate.
6.2.6 Describe the effect of a catalyst on a chemical reaction.
6.2.7 Sketch and explain Maxwell-Boltzmann curves for reactions
with and without catalysts.
27
27
6.2.1 Describe the kinetic theory in terms of the
movement of particles whose average energy is
proportional to temperature in Kelvins.



Kinetic Theory
Tiny particles in all forms of matter are in constant motion.
The energy an object has because of motion is called kinetic
energy.
The particles in any collection of atoms or molecules at a given
temperature have a wide range of kinetic energies, from very
low to very high. Overall, there is an average kinetic
energy.
28
28
6.2.1 Describe the kinetic theory in terms of the
movement of particles whose average energy is
proportional to temperature in Kelvins.



What happens when a substance is heated?
Particles absorb energy, some of which is stored within the
particles (potential energy). This does not affect temperature.
The remaining energy speeds up particles (increases the
average kinetic energy) which increases temperature.
We’ll look at this distribution in our last asmt stmt…
29
29
6.2.1 Describe the kinetic theory in terms of the
movement of particles whose average energy is
proportional to temperature in Kelvins.


Kelvin Temperature
As a substance cools, the
average kinetic energy
declines and the particles
move more slowly.
At some point, the
temperature will be low
enough so the particles will
stop moving and have no
kinetic energy. That point
is called absolute zero (0 K
, -273.15 oC)
30
30
6.2.1 Describe the kinetic theory in terms of the
movement of particles whose average energy is
proportional to temperature in Kelvins.


Proportionality
Kelvin temperature of a
substance is directly
proportional to the
average kinetic energy
of the particles.
For example: The particles
in helium gas at 200 K
have twice the kinetic
energy average as the
particles in helium gas at
100K
31
31
6.2.1 Describe the kinetic theory in terms of the
movement of particles whose average energy is
proportional to temperature in Kelvins.
32
32
6.2.2 Define the term activation energy, Ea.


Activation Energy: The minimum amount of energy
required for reaction: to happen is called the
activation energy, Ea.
Just as a ball cannot get over a hill if it does not roll
up the hill with enough energy, a reaction cannot
occur unless the molecules possess sufficient energy
to get over the activation energy barrier.
33
33
6.2.2
Define the term activation energy, Ea.
Activation energy is the
minimum amount of kinetic
energy that must be given
to the reactants before they
will react.
(a) Activation energy (Ea) of
forward reaction
(b) Activation energy (Ea) of
reverse reaction
(c) Heat of reaction (ΔH)
What is the heat of reaction for
each of the graphs? Which is
exothermic and which is
endothermic?
34
34
6.2.3 Describe the collision theory.



In a chemical reaction, bonds are broken
and new bonds are formed.
There are a few criteria that need to occur
in order for particles to undergo this
process
This establishes the 3 parts of the Collision
Theory
35
35
6.2.3 Describe the collision theory.

COLLISION THEORY: Know this!!!
1. In order for a reaction to occur,
reactants (atoms, ions, radicals or
molecules) must collide
2. Molecules must collide in the correct
orientation so that the reactive parts
of each of the two particles come into
contact with each other
3. Reactants must collide with sufficient
kinetic energy to bring about the
reaction (at least the activation energy,
Ea).
36
6.2.3 Describe the collision theory.


Not all collisions result in a reaction.
The likelihood of these happening determines
the rate.
37
6.2.4 Changing the rate of reactions
Anything that increases the number of successful collisions
between reactant particles will speed up a reaction.
What factors increase the rate of reactions?
 increased temperature
 increased concentration of
dissolved reactants, and increased
pressure of gaseous reactants
 increased surface area of solid
reactants (decrease particle size)
 use of a suitable catalyst.
38
6.2.4 Effect of temperature on rate
Increasing the temperature will make the
particles move faster, so there will be more
collisions.
At a higher temperature, many
more particles will have the
necessary activation energy. The
ratio of successful collision to
unsuccessful collisions will
increase.
Which one will contribute more
towards increasing the rate of
reaction?
39
6.2.4
Predict and explain, using the collision theory, the qualitative
effects of particle size, temperature, concentration and pressure on
the rate of a reaction.

Temperature: greater T = more
collisions with the required Ea

Increasing the temperature increases
the frequency of collisions but, more
importantly, the proportion of
molecules with E ≥ Ea increases
40
Effect of concentration on rate of reaction
The higher the concentration of a dissolved reactant, the
faster the rate of a reaction.
Why does increased concentration increase the rate of
reaction?
At a higher concentration, there are more particles in the
same amount of space. This means that the particles are
more likely to collide and therefore more likely to react.
The ratio of successful collisions to unsuccessful collisions
will stay the same, but there will be more successful
collisions.
lower concentration
higher concentration 41
Effect of pressure on rate of reaction
Why does increasing the pressure of gaseous reactants
increase the rate of reaction?
As the pressure increases, the space in which the gas
particles are moving becomes smaller.
The gas particles become closer together, increasing the
frequency of collisions. This means that the particles are more
likely to react.
lower pressure
higher pressure
42
6.2.4
Predict and explain, using the collision theory, the qualitative
effects of particle size, temperature, concentration and pressure on
the rate of a reaction.

Concentration: higher conc = more
particles in given volume = more
collisions per unit of time (greater
frequency of collisions)
43
Effect of surface area on rate of reaction
Any reaction involving a solid can only take place at the
surface of the solid.
If the solid is split into several pieces, the surface area
increases. What effect will this have on rate of reaction?
low surface area
high surface area
This means that there is an increased area for the reactant
particles to collide with.
The smaller the pieces, the larger the surface area. This
means more collisions and a greater chance of reaction.
44
6.2.4
Predict and explain, using the collision theory, the qualitative
effects of particle size, temperature, concentration and pressure on
the rate of a reaction.

Particle size: greater surface area =
greater # of collisions per unit of
time (greater frequency of
collisions)
45
6.2.6 Describe the effect of a catalyst on a
chemical reaction.

A catalyst is a substance that increases
the rate of a chemical reaction without
itself being chemically changed at the
end of the reaction.
46
46
What are catalysts?
Catalysts are substances that change the rate of a reaction
without being used up in the reaction.
Catalysts never produce more product – they just
produce the same amount more quickly.
energy (kJ)
Ea without
catalyst
Different catalysts work in
different ways, but most
provide an alternative path
with lower activation
energy (Ea).
Ea with
catalyst
reaction (time)
47
6.2.4
Predict and explain, using the collision theory, the qualitative
effects of particle size, temperature, concentration and pressure on
the rate of a reaction.
Catalysts: provides alternative
mechanism w/ lower Ea = more
collisions with sufficient Ea
**You must not say “the number of
collisions” only... You must say
“number of collisions per unit time”
or “frequency”

48
In order to increase reaction rates…
Factor
Optimal Condition
to Increase
Reaction Rate
Reason
Particle Size
Concentration
Pressure
Temperature
Catalyst
49
In order to increase reaction rates…
Factor
Increase
Rxn Rate
by…
Particle Size
Decrease
Increases the surface area  greater
frequency of collisions.
Concentration
Increase
Greater frequency of collisions
Pressure
Increase
Greater frequency of collisions
Temperature
Increase
Greater frequency of collisions &
the proportion of molecules with E ≥ Ea
increases
Catalyst
Add one

Changing the mechanism of the reaction
and lowering the Ea
Reason
50
6.2.5 Sketch and explain qualitatively the Maxwell-Boltzman
energy distribution curve for a fixed amount of gas at different
temperatures and its consequences for changes in reaction rate.

Temperature is
defined as a
measure of the
average kinetic
energy of the
molecules in a
sample.
• At any temperature there is a wide
distribution of kinetic energies.
51
51
6.2.5 Sketch and explain qualitatively the Maxwell-Boltzman
energy distribution curve for a fixed amount of gas at different
temperatures and its consequences for changes in reaction rate.


As the temperature
increases, the
curve flattens and
broadens.
Thus at higher
temperatures, a
larger population of
molecules has
higher energy.
52
52
6.2.5 Sketch and explain qualitatively the Maxwell-Boltzman
energy distribution curve for a fixed amount of gas at different
temperatures and its consequences for changes in reaction rate.

If the dotted line represents the activation
energy, as the temperature increases, so does
the fraction of molecules that can overcome
the activation energy barrier.
• As a result, the
reaction rate
increases.
53
53
6.2.5
Sketch and explain qualitatively the MaxwellBoltzman energy distribution curve for a fixed amount of
gas at different temperatures and its consequences for
changes in reaction rate.



Area under the graph
directly related to the
number of particles
present.
The average kinetic energy
close to the peak of the
graph.
Area to the right of the Ea
depicts the number of
particles having sufficient
energy to react.
54
54
6.2.5
Sketch and explain qualitatively the MaxwellBoltzman energy distribution curve for a fixed amount of
gas at different temperatures and its consequences for
changes in reaction rate.



At higher temps, the
distribution flattens out as
more molecules gain
kinetic energy.
Area under the curves
remain the same since the
number of particles remain
the same.
However more particles
have kinetic energies
greater than the Ea as
temperature increases.
55
55
6.2.6 Describe the effect of a catalyst on a chemical
reaction.



Catalysts increase the rate of reactions by providing a new
alternative pathway or mechanism for the reaction that has
a lower barrier height than the uncatalyzed reaction.
Catalysts increase the forward and reverse rates of
reversible reactions and increase the rate of irreversible
reactions.
Catalysts do NOT alter the position of equilibrium.
56
56
6.2.6 Describe the effect of a catalyst on a
chemical reaction.
Since catalysts lower the forward and backward activation
energy barrier heights (Ea), the rates of the forward and
backward reactions increase to the same degree.
57
57
6.2.6 Describe the effect of a catalyst on a
chemical reaction.
In a reaction with a catalyst, without increasing the
temperature, a larger number of particles has the value of
kinetic energy greater than the activation energy (E ≥ Ea)
and so will be able to undergo successful collisions.
58
58
6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without catalysts.

Catalysts lower the activation energy barrier by positioning
reactant particles favorably. More reactant particles possess
this lower activation energy so the rate increases.
59
59
6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without catalysts.
60
60
6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without catalysts.
61
61
6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without catalysts.
62
62
6.2.7 Sketch and explain Maxwell-Boltzmann
curves for reactions with and without catalysts.

Maxwell Boltzmann
http://www.youtube.com/watch?v=66AxPceP
5QY
63
63
Terms to Know







Rate of reaction
Collision theory
Activation energy
Temperature
Catalyst
Maxwell-Boltzmann distribution
Factors affecting reaction rate (no need to
define, just describe/explain the impact of
each factor)
64
Resources




Steve Colgan
http://www1.kent.k12.wa.us/staff/AdamSkagen/IBp
pts/6-Kinetics-09.pdf
http://dl.dropbox.com/u/7362447/GMHS_Chemistry
/Unit%20Documents/Kinetics/Notes/Notes_Reaction
Rates.doc
http://www.chemactive.com/ppt/ib/ib_kinetics_bw.p
pt#604,33,What are catalysts?
65