1. Chemistry 40S
Unit 3 - Kinetics

2. About Kinetics
Chemistry focuses largely on chemical
reactions. Some reactions, such as rusting,
occur very slowly, while others, like
explosions, occur quite fast. Kinetics is the
branch of chemistry that studies the speed
or rate with which chemical reactions occur.
Some reactions do not occur in one simple
step. Some occur in several more complex
steps. Part of Kinetics is studying the steps,
or mechanism, of a chemical reaction.

3. Defining Rate
• Recall that a chemical reaction is one in
which a new substance is formed. They
are usually expressed in the shorthand
form, known as a chemical equation.
REACTANTS →
PRODUCTS
• The rate of a reaction refers to the speed
with which a chemical reaction takes
place. That is, how fast reactants are used
up or products are formed.

4. Measuring Rate
• Rate can be measured using several different
methods.
• The reaction below shows that A and B
combine to form C and heat is given off:
A(s) + B(l) → C(g) + heat
• In this reaction, rate can be measured in
terms of several different properties:
– i) Temperature change over time (°C/min)
– ii) Pressure change over time (kPa/s or mmHg/s)
– iii) Mass change over time (g of C/min)
– iv) Other possibilities:

5. Measuring Rate
Therefore generally,
Rate is usually described in terms of change in concentration
of reactant or product over time.
Or, in the case of our example: A(s) + B(l) → C(g) + heat

6. Average Rate

7. Average Rate
• The numerical value for the rate of a reaction can be
determined by examining the change in the amount
of a substance at a particular time, or over a period of
time. The graphical representation of this may look
like the figure below:

8. Average Rate

9. Average Rate
• The average rate for a reaction is given by the
following equation:
• The instantaneous rate of a reaction is the rate
at a specific time. This is determined by
calculating the slope of the line tangent to the
point on the concentration vs time curve.

10. Average Rate
• Example 1
– According to the reaction A → B, the following data was
collected:
a) What is the average rate over the entire 50
seconds?
b) What is the average rate for the interval 20
s to 40 s?

11. Average Rate
• Example 2
– The decomposition of nitrogen dioxide produces nitrogen
monoxide and oxygen according to the reaction:
2 NO2(g) → 2 NO(g) + O2(g)
– the following data has been collected:
Calculate:
a) The average rate of decomposition of NO2 over 400 s.
b) The average rate of production of NO over 400 s.
c) The average rate of production of O over 400 s.

12. Rate and Stoichiometry

13. Graphical Determination of
Average Rate

14. Rate and Stoichiometry
• Rate of creation of products or disappearance of
reactants can be predicted from reaction stoichiometry.
• From the example on the previous page, for the reaction:
2 NO2(g) → 2 NO(g) + O2(g)

15. Rate and Stoichiometry
• From the graph and the calculations, it can be
seen:
– the rate of decomposition of NO2 is equal to the
production of NO. The ratio of the molar coefficients is
1:1. This means as one molecule of NO2 is
decomposed, one molecule of NO is created. Their
rates should be equal.
– the rate of production of oxygen is half that of the NO.
The ratio of the molar coefficients is 2:1, as well. The
rate of production of NO should be double that of the
oxygen.
• Therefore,

16. Rate and Stoichiometry
• Example 1
– For the decomposition of nitrogen dioxide,
if the rate of decomposition is determined
to be 0.50 mol/Ls at a certain
temperature, predict the rate of creation
of both products.
2 NO2(g) → 2 NO(g) + O2(g)

17. Rate and Stoichiometry
• Example 2
– For the reaction 2 A + B → 3 C, what is
the rate of production of C and the rate of
disappearance of B if A is used up at a
rate of 0.60 mol/Ls?

18. Collision Theory
A Thought Experiment
• A group of students are taken into the school hall and
blindfolded. They are asked to move around the hall. If
two students crash into each other and both students
fall over, then they stay lying on the ground - they have
"reacted".
• How can we increase the rate at which students fall
down? There are two obvious choices:
– 1) Put more students into the hall. This will lead to more
collisions between students.
– 2) Ask the students in the hall to run more quickly. This will also
lead to more collisions and to a greater chance of a collision
leading to the students falling over.
• We will use this "thought experiment" to try and explain
rates of reaction.

19. Collision Theory
• The collision theory attempts to explain known
information concerning the rates of chemical
reactions. It uses kinetic molecular theory, and
other factors, such as molecular shape and
structure, in its explanations.
• The basic assumption of the collision theory is
that for a reaction to occur the reacting species
must collide. The collision theory applies to
three factors effecting reaction rate (effect of
concentration, effect of temperature, effect of
surface area).

20. Collision Theory
• The effect of concentration
• The effect of temperature
• The effect of particle size

21. Collision Theory
• The rate of chemical reaction, specifically
depends upon the following two factors:
– Number of collisions that occur per unit
of time (i.e. frequency of collisions).
– The fraction of the collisions that is
effective in causing a “reaction”.

22. Collision Theory

23. Collision Theory
• Fast Reactions:
F reactants that are in a liquid (aq) or
gaseous (g) state have greater mobility of
the molecules, thus a good chance of
contacting each other
co when reactants are ionic, the positive-
negative attraction works rapidly

24. Collision Theory
• Slow Reactions:
Sl when reactants are both solids, the
particles have very little mobility
pa often slow reactants are neutral
molecules, this means that electron
transfer and bond arrangement takes
some time to occur
so complex structures take longer to
rearrange than do small ones (LEGO
analogy)

25. Collision Theory
• The collision theory suggest an additional
reason as to why it may take “a while” for
certain species to react:
• Collisions will only successful in producing a
reaction if the molecules have a particular
orientation, so that particular chemical bonds
may be either broken or formed.

26. Activation Energy
• Particles must collide with enough energy in
order to cause a reaction.
• Chemical reactions involve the making and
breaking of bonds, which requires energy. If
colliding particles do not have sufficient velocity
(Kinetic Energy) they will not produce a
reaction.
• The minimum amount of energy required for
colliding particles to produce a chemical reaction
is called the Activation Energy (EA) of that
reaction.

27. Collision Theory
& Activation Energy

28. Activation Energy
• The greater the activation energy, the slower the
reaction rate, the longer the reaction takes.
• The rate of a reaction is determined by the
height of the barrier that the particles must cross
in order that they are converted into products.
• The activation energy is the barrier colliding
particles must overcome.

29. Activation Energy
• As the particles collide, they form an unstable
intermediate particle called the activated
complex or transition state. The energy required
to produce the activated complex is the
activation energy.
Enthalpy (H) is the heat content or total energy
possessed by the particles in a system.

30. Activation Energy
• The energy released or absorbed by a reaction is
called the change in enthalpy, ∆H, or heat of reaction.
∆H = HPRODUCTS - HREACTANTS
• where ∆H is the change in enthalpy or energy that
occurs during a reaction in Joules (J).
• If ∆H is negative, heat flows out of the system. This
type of reaction gives off heat and the reaction vessel
feels warmer. This is called an exothermic reaction.
• If ∆H is positive, heat is absorbed or flows into the
system. The reaction vessel feels cooler as energy is
absorbed from the surroundings. This type of reaction
is called an endothermic reaction.

31. Reaction Coordinate Diagrams

32. Reaction Coordinate Diagrams

33. Reaction Coordinate Diagrams
•The activated complex in this reaction is CH3CH2(OH)Br¯.
•The activation energy is 88.9 kJ per mole of CH3CH2Br.
•The enthalpy change is -77.2 kJ, indicating an exothermic
reaction.
•Consequently, this reaction does not take place unless 88.9 kJ
per mole of CH CH Br is added to the system.

34. Factors Affecting
Reaction Rate
• The rate of a reaction can be affected by a
number of different variables. Listed below
some factors that may affect the rate of
reaction.
– Nature of Reactants
– Concentration of Reactants
– Temperature
– Catalysts
– Surface Area
– Pressure (Volume)

35. Factors Affecting
Reaction Rate
• Nature of Reactants
Some chemical reactions involve the rearrangement of
atoms as a result of bond breaking and new bond formation.
Other reactions are a result of electron transfer. The nature
of the reactants involved in the reaction will affect the rate of
the reaction.
i) The fewer the number of bonds broken, the faster the rate.
The equations give us an indication as to why:
2 NO(g) + O2(g) 2 NO2(g)
2 C8H18(l) + 25 O2(g) 16 CO2(g) + 18 H2O(g)

36. Factors Affecting
Reaction Rate
• Nature of Reactants
ii) In general, the weaker the bonds in the reactants,
the faster the reaction.
Reactions with molecular reactants (covalent
bonds) are usually slower than ionic reactions.
iii)The state or phase of the reactants will also affect
the rate of a reaction.
Reactions with gases as reactants are usually faster
than those with liquids, which are both faster than
those with solids.

37. Factors Affecting
Reaction Rate
Temperature
• According to the collision theory, reaction rate is
determined by the frequency of collisions
between molecules and increases as the
frequency of collisions increases. According to
the Kinetic Molecular Theory, the speed (Kinetic
Energy) of particles increases as the temperature
increases and the more likely the reactant
particles will have enough energy to overcome
the energy barrier to form products. A general
rule is that for every 10 ºC rise in temperature the
reaction rate doubles.

38. Factors Affecting
Reaction Rate
Temperature

39. Factors Affecting
Reaction Rate
Concentration of Reactants
• Recall that concentration refers to the
amount of reactant per unit volume. The units
for concentration are mol/L.
• Increasing the concentration of reactants
increases the total number of particles in a
container. If the number of particles in a
container increases, so do the number of
particles with activation energy. More
particles with activation energy increases the
frequency of effective collisions.

40. Factors Affecting
Reaction Rate
Concentration of Reactants
• Increasing the number of particles in a container will
also increase the chances of a collision. The diagram
below illustrates the effect of doubling the
concentration of just one of the reactants:
– In (a), with two particles of each, there exists four possible
collisions which could produce a reaction.
– In (b), by doubling the number of red particles, the number
of possible collisions increases to eight! This will increase
collision frequency.

41. Factors Affecting
Reaction Rate
Catalysts
• A catalyst is a substance that speeds up or initiates a
reaction without itself being permanently changed.
• Catalysts reduce the activation energy needed for a
reaction to occur and so increase the number of
effective collisions. Catalysts speed up reaction rate but
do not enter into the reaction and are therefore not
changed or used up. This makes them very effective
molecules for creating change. Our body has many
enzymes (organic catalysts) which speed up reaction
rates allowing the body to control its metabolism.
• An inhibitor also affects reaction rates by stopping or
slowing down reactions. It is the opposite of a catalyst.

42. Factors Affecting
Reaction Rate
Catalysts

43. Factors Affecting
Reaction Rate
Pressure
• Changing the pressure on a system usually only affects
the reaction rates of gaseous reactions.
• Pressure is the force of particles upon the walls of their
container. If the number of particles in a container
increases without changing volume, the pressure
increases.
• There are 3 ways to change pressure:
– Add more product and/or reactant particles to the
container.
– Increase or decrease the volume of the container
– Add an inert or unreactive gas.

44. Factors Affecting
Reaction Rate
Pressure
• If pressure is increased by adding more reactant particles, the
concentration increases causing an increased rate. If the pressure is
reduced by removing reactant, the rate decreases due to decreased
concentration.
If you decrease the volume of a container without changing the
number of particles in the container, as in figure c in the diagram
below, the concentration of the reactants increases.
If you increase the volume of a container without changing the number
of particles, as in figure a in the diagram above, the concentration of
the reactants decreases.

45. Factors Affecting
Reaction Rate
Surface Area
• The size of the reactants, or surface area in contact, is
usually only a factor in heterogeneous reactions.
Increasing the surface area of the reactants by crushing,
grinding or other means, increases the number of particles of
reactants in contact. The rate of reactions increases when
surface area in contact also increases.
• Increasing surface area increases the frequency of collisions,
increasing reaction rate.

46. Collision Theory &
Reaction Mechanisms
• The rate of a reaction cannot always be
determined from the chemical equation of the
reaction. Consider the reaction:
2 NO (g) + O2 (g) → 2 NO2 (g)
• According to the collision theory, for this
reaction to occur in one step, 3 particles must
collide: two NO molecules and one O2. These
particles must collide at exactly the same time
with the correct orientation and enough
energy.

47. Collision Theory &
Reaction Mechanisms
• Three particle collisions are quite rare. We
can think of a three particle collision in terms
of a pool shot. Striking the cue ball to create a
collision which puts one ball in a pocket is
difficult. Making a "combination" shot, striking
two balls with enough energy and the correct
angles to sink both, is much more difficult,
and rarely attempted.

48. Collision Theory &
Reaction Mechanisms
• Chemical reactions tend to take place in
steps, with each step involving a collision
between two particles, or bimolecular. The
chances of a two particle collision being
successful are much better than a collision
between three or more particles.

49. Collision Theory &
Reaction Mechanisms
Reaction Intermediates
• Reactions which take place in one elementary step
are called, appropriately, simple reactions.
• Reactions which take place in more than one step
are called complex reactions.
• The complex reaction:
2 NO (g) + O2 (g) → 2 NO2 (g)
does not take place in one elementary step, but
actually takes place in two steps:

50. Collision Theory &
Reaction Mechanisms
Reaction Intermediates
• The complex reaction:
2 NO (g) + O2 (g) → 2 NO2 (g)
• does not take place in one elementary step,
but actually takes place in two steps:

51. Collision Theory &
Reaction Mechanisms
Reaction Intermediates
• The complex reaction:
2 NO (g) + O2 (g) → 2 NO2 (g)
• Compounds such as N2O2, products of one reaction
that immediately become reactants in another
reaction, are called reaction intermediates. All
complex reactions contain at least one reaction

52. Collision Theory &
Reaction Mechanisms
Net Equations
• The steps in which a reaction occurs is called
that reaction's mechanism. The sum of the
steps of a mechanism must equal the total or net
equation.
• For the reaction, NO2 + CO → NO + CO2
the mechanism is as follows:

53. Collision Theory &
Reaction Mechanisms
Net Equations
• This mechanism agrees with the initial equation
for this reaction, as the sum of the steps equals
the original reaction. The NO3 is the reaction
intermediate, so it does not appear in the net
reaction. Since NO2 is found twice on the left and
once on the right, we can cancel one of the
NO2's just as we would adding equations in
math.

54. Collision Theory &
Reaction Mechanisms
Rate Determining Step
• Not all steps in a mechanism have the same rate.
The step with the slowest rate is called the rate
determining step (RDS), since that step affects
the rate of the reaction more than the others.
• The slowest step in the reaction mechanism
determines the rate, and therefore only reactants
in the slowest step affect rate. Increasing the
concentrations of these reactants will speeding up
the slow step will increase the overall reaction rate.
Increasing the concentration of reactant in fast
steps will have little effect on rate. To understand
how this works consider the following analogy.

55. Collision Theory &
Reaction Mechanisms
Rate Determining Step
• A student has to get to school in the morning, after having a good
night's sleep in his own bed. The student might perform the following
tasks (steps) in order to get to school (overall reaction)?
• Sleeping in bed (reactant) --------> Arriving at school (product)
– wake up
– uses the washroom to clean up and have a shower.
– have breakfast
– walk, ride with family or friends or take the bus to school
• The slowest step in this sequence will determine how fast the student
can get to school (product). Waking up may be the slow step for some
while taking a shower may be the slow step for others.
• Speeding up the slow step will speed up the overall reaction. If for
example taking a shower was the slow step the student might shower
the night before and simply wash their face in the morning. If taking the
bus is the slow step they might get a ride with a friend that might be
quicker and therefore speed up the mechanism.

56. Collision Theory &
Reaction Mechanisms
Rate Determining Step
• Assignment :Analogies of Reaction Rates
• List 4 jobs that you do either at home or at work, that involve
more that one task or step to accomplish.(For example
cleaning your room, or cooking chicken at a local
restaurant)
• List what might be considered the starting point (reactant).
• List what might be considered the finished product.
• List the steps required and the amount of time needed to
complete each step.
• Describe which step is rate determinining.
• Describe how this step might be sped up.

57. Collision Theory &
Reaction Mechanisms
Rate Determining Step
• According to our last mechanism:
step 1 is the RDS since it is the slowest step.

58. Collision Theory &
Reaction Mechanisms
Energy Changes
• According to our last mechanism:
• if the reaction is exothermic, we can predict
its shape:

59. Collision Theory &
Reaction Mechanisms
Energy Changes
• Without knowing the actual energy
values, we can predict the shape of the
curve:
– Recall that the larger the activation
energy, the slower the reaction. Step 1
is the rate determining step, so it
should have the largest activation
energy.
– Step 2 is fast so it should have the
lowest activation energy. Notice the
activation energy for Step 2 begins
where Step 1 ends.
– The enthalpy of reactants should be
higher than the enthalpy of products

60. Collision Theory &
Reaction Mechanisms
Mechanisms Problem
• Since the RDS affects the rate of the entire reaction the most,
changes to the reactants in the other steps will have very little effect
on the rate of the reaction.
• Example 1
• Given the following mechanism:
P+Q → X+T (slow)
X+P → Y+R (fast)
Y+S → T (moderate)
a) What is the net reaction?
b) What are the reaction intermediates?
c) Which is the rate determining step?
d) What would be the effect of increasing the concentration of P?
e) What would be the effect of decreasing the concentration of Q?
f) What would be the effect of increasing the concentration of S?

61. Rate Law
• Rate Law is an expression which relates the rate of a
reaction to the concentration of the reactants. Rate
law is a tool which helps us calculate the rate of a
reaction with given concentrations of reactants.
• The rate law shows the quantitative effect of
concentration on reaction rate.
• For the reaction:
A + B → products
• The rate of consumption of A and B is directly
proportional to their concentration. That is, the greater
the concentration of A and B, the faster the rate.

62. Rate Law
The Molarity Relationship
• M = n,
V
• the molar concentration α number of moles of substance
– therefore R α [A] and R α [B]
– AND: R α [A]x [B]y AND: R = k [A]x [B]y
• R = rate of reaction (Moles/1/s or M/s)
• [A] = molar concentration of A
• [B] = molar concentration of B
• k = specific rate constant
• x, y = is the power, called the order of the reaction

63. Rate Law
The Molarity Relationship
• R = k [A]x [B]y
• The constant k, is known as the specific rate constant
for the reaction. The rate constant and the order can
only be determined experimentally.
• The rate constant is specific for each reaction at a
specific temperature, since it’s value depends upon the
size, speed and types of molecules in the reaction.
• Changing temperature would change the speed of the
reactant particles and hence change the rate constant.
• Temperature is the only factor, which affects the rate
constant.

64. Rate Law
The Law of Mass Action
• The rate of a reaction is proportional to the
product of the concentrations of the
reactants, each raised to a power equal to
the coefficient in the original equation.

65. Rate Law
Reaction Order
• The order of a reaction indicates how
concentration of reactants affects the rate of a
reaction.
• Example, A → products,
• If the order of a reaction is a first order reaction,
x = 1, this means the reaction rate was directly
proportional to changes in reactant concentration.
If the concentration of A were doubled, the rate
would double. If the concentration were tripled,
the rate would triple, etc.

66. Rate Law
Reaction Order
• If the reaction is a second order reaction, x = 2,
doubling the concentration would increase the rate by
a factor of 2x = 22 = 4. That is, the rate would
increase four times. Tripling the concentration of A
would cause the rate to increase nine times (3x=32=9).
• If the rate of the reaction did not depend on the
concentration of A, it would be a zero order reaction,
x = 0. This means a change in the concentration of A
does NOT change the rate of the reaction.

67. Rate Law
Reaction Order
• For a reaction with more than one reactant, such as
A + B → products
• The rate law would be: Rate = k[A]x[B]y
• The rate depends on both A and B concentrations.
Each reactant can affect the rate differently.
• The total order of the reaction is the sum of the
order with respect to A and the order with respect to
B, x + y.

68. Rate Law
Calculating Rate Law
• The rate law can only be determined experimentally.
Even though the rate of formation of the products
and the rate of consumption of reactants is related
to their reaction stoichiometry, the rate law cannot
usually be determined from the molar coefficients.
• One method of determining rate law is by measuring
the effect of changes in concentration of one
reactant on the initial reaction rate, while keeping
the other reactant concentrations constant.

69. Rate Law
Calculating Rate Law
• Example 1:
• What is the rate law for the following reaction,
given the experimental data below?
H2O2 + 2 HI → 2 H2O + I2

70. Rate Law
Calculating Rate Law
• Example 2:
• For the reaction
3 A(g) + B(g) + 2 C(g) → 2 D(g) + 3 E(g)
a. Write the rate law for this reaction.
b. Calculate the value of the rate constant.
c. Calculate the rate for Trial #5.
d. Calculate the concentration
of A in Trial #6
.

71. Rate Law
Rate Law & Stoichiometry
• For most reactions, the rate law, the specific rate constant,
k, and the mechanism of a reaction can only be determined
experimentally, not from the reaction stoichiometry.
• However, for reactions that occur in a single step
(elementary reactions) the order of each reactant in the
rate law is equal to the coefficient in the reaction's
balanced equation.
• For the elementary reaction:
aA + bB → cC + dD
• the rate law is
rate = k[A]a[B]b
• where a and b are the molar coefficients for the elementary
reaction.

72. Rate Law
Rate Law & Reaction Mechanism
• We have seen that rate law for reactions
which occur in more than one step do NOT
correspond to the coefficients in its
balanced net equation.
• The rate law for these equations
corresponds to the stoichiometry of the rate
determining step. Those reactants which
are zero order, do not appear in the RDS
and, hence do not affect reaction rate
significantly.

73. Rate Law
Rate Law & Reaction Mechanism
• Example
• The mechanism for the reaction
3M + N → P + 2Q
• is below:
2M → P + X (slow)
X+M→ Q + Y (moderate)
N+Y → Q (fast)
a) What is the rate law for this reaction?
b) What would be the effect of tripling the [M]?
c) What would be the effect of doubling the [N]?