Chemical kinetics power point, energy changes, IGCSE chemistry

1. Chemical Kinetics

2. Chemical Kinetics
• Chemical Kinetics – the area of chemistry
concerned with the speeds, or rates, at which
chemical reactions occur
Study of speed with which a chemical reaction
occurs and the factors affecting that speed.
Provides information about the feasibility of a
chemical reaction.
Provides information about the time it takes for a
chemical reaction to occur.
Provides information about the series of
elementary steps which lead to the formation of
product.

3. Factors that
Affect Reaction
Rates

4. Rate Factors: Physical State
• 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.

5. Rate Factors: Concentration
• Concentration of Reactants
As the concentration of reactants increases, so
does the likelihood that reactant molecules will
collide.

6. Rate Factors: Temperature
• Temperature
At higher
temperatures,
reactant molecules
have more kinetic
energy, move faster,
and collide more
often and with
greater energy.

7. Rate Factors: Catalyst
• 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.

8. Rate Factors: Pressure &
Surface Area
• Pressure and Surface Area
Pressure – an increase in the
pressure on a gaseous
reactant causes an increase in
concentration and increases
the speed of the reaction
Surface Area – more surface
area (smaller particle size)
increases the reaction surface
and the speed of the reaction,
assuming equal quantity of the
reactant

9. Reaction Rates

10. Reaction Rate
• Reaction Rate – the decrease in
concentration of a reactant or the increase in
concentration of a product with time
Rates of reactions can be determined by
monitoring the change in concentration of either
reactants or products as a function of time.

11. Defining Reaction Rate
•

12. Reaction Rates
C4
H9
Cl(aq) + H2
O(l) ⎯⎯→ C4
H9
OH(aq) + HCl(aq)
• In this reaction, the concentration
of butyl chloride, C4
H9
Cl, was
measured at various times.
• The average rate of the reaction
over each interval is the change
in concentration divided by the
change in time:
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.
Average rate =
Δ[C4
H9
Cl]
Δt

13. Reaction Rates
• 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.
C4
H9
Cl(aq) + H2
O(l) ⎯⎯→ C4
H9
OH(aq) + HCl(aq)

14. Instantaneous Rates

15. Reaction Rates
• All reactions slow
down over time.
• Therefore, the best
indicator of the rate
of a reaction is the
instantaneous rate
near the beginning.
C4
H9
Cl(aq) + H2
O(l) ⎯⎯→ C4
H9
OH(aq) + HCl(aq)

16. Reaction Rates and
Stoichiometry
• In this reaction, the ratio
of C4
H9
Cl to C4
H9
OH is
1:1.
• Thus, the rate of
disappearance of
C4
H9
Cl is the same as
the rate of appearance
of C4
H9
OH
C4
H9
Cl(aq) + H2
O(l) ⎯⎯→ C4
H9
OH(aq) + HCl(aq)
Rate =
-Δ[C4
H9
Cl]
Δt
=
Δ[C4
H9
OH]
Δt

17. Reaction Rates and
Stoichiometry
• What if the ratio is not 1:1?
• Therefore…
• To generalize, then, for any reaction
2 HI(g) ⎯⎯→ H2
(g) + I2
(g)
Rate = − 1
2
Δ[HI]
Δt
= Δ[I2
]
Δt
aA + bB cC + dD
Rate = −
1
a
Δ[A]
Δt
= −
1
b
Δ[B]
Δt
=
1
c
Δ[C]
Δt
1
d
Δ[D]
Δt
=

18. Rate Laws

19. Concentration and Rate
• One can gain information
about the rate of a reaction
by seeing how the rate
changes with changes in
concentration.
Comparing Experiments 1
and 2, when [NH4
+
] doubles,
the initial rate doubles.
• Likewise, comparing
Experiments 5 and 6, when
[NO2
−
] doubles, the initial
rate doubles.
NH4
+
(aq) + NO2
−
(aq) N2
(g) + 2 H2
O(l)

20. Concentration and Rate
• This means
Rate ∝ [NH4
+
]
Rate ∝ [NO2
−
]
Rate ∝ [NH4
+
] [NO2
−
]
• or
Rate = k [NH4
+
] [NO2
−
]
• This equation is called the rate law, and k is
the rate constant.

21. Rate Law
• Rate Law – an
equation that relates
the reaction rate to
the concentrations of
reactants
The exponents tell
the order of the
reaction with respect
to each reactant.

22. Rate Constant
• Rate Constant – a
constant of
proportionality
between the reaction
rate and the
concentrations of
reactants that
appear in the rate
law

23. Reaction Order
• Reaction Order – the power to which the
concentration of a reactant is raised in a rate law
• Overall Reaction Order – the sum of the reaction
orders of all the reactants appearing in the rate
expression when the rate can be expressed as:
rate = k[A]a
[B]b
…..

24. Initial Rate Law Method

25. Determining Rate Laws
• Rate laws are ALWAYS determined experimentally.
• Reaction order is ALWAYS defined in terms of
reactant concentrations.
• The order of a reactant is NOT related to the
stoichiometric coefficient in the balanced chemical
equation

26. Integrated Rate
Laws

27. Integrated Rate Laws
• Using calculus to integrate the rate law for a
first-order process gives us…
• Where
[A]0
is the initial concentration of A.
[A]t
is the concentration of A at some time, t, during the
course of the reaction.
• Manipulating this equation produces…
• …which is in the form y = mx + b
ln
[A]t
[A]0
= −kt
ln [A]t
= − kt + ln [A]0

28. Zero-Order Processes

29. First Order Process
• First Order Reaction
– a reaction in which
the reaction rate is
proportional to the
concentration of a
single product, raised
to the first power
When ln[ ] is plotted
as a function of time, a
straight line results.

30. Second-Order Processes
• Second Order
Reaction – a
reaction in which the
overall reaction order
(the sum of the
concentrations-term
exponents) in the
rate law is 2
• When 1/[ ] is plotted
as a function of time,
a straight line results.

31. Half - Life
• Half-Life – the time
required for the
concentration of a reactant
substance to decrease to
half its initial value; the time
required for half of a
sample of a particular
radioisotope to decay
• For a first-order process,
this becomes
= t1/2
0.693
k
NOTE: For a first-order
process, the half-life does
not depend on [A]0
.

32. Half-Life
• For a second-order process,

33. Rate Law Summary Chart

34. Temperature and
Rate

35. Temperature and Rate
• Generally, as
temperature
increases, so does
the reaction rate.
• This is because k is
temperature
dependent.

36. Collision Model
• Collision Model – a model of
reaction rates based on the idea that
molecules must collide to react; it
explains the factors influencing
reaction rates in terms of the
frequency of collisions, the number of
collisions with energies exceeding
the activation energy, and the
probability that the collisions occur
with suitable orientations
In a chemical reaction, bonds are
broken and new bonds are formed.
Molecules can only react if they
collide with each other.
Furthermore, molecules must collide
with the correct orientation and with
enough energy to cause bond
breakage and formation.

37. Activation Energy
• Activation Energy – (Ea
) the minimum
energy needed for reaction; the height of the
energy barrier to formation of products
In other words, there is a minimum amount of
energy required for reaction: the activation
energy, Ea
.

38. Reaction Coordinate Diagrams
• It is helpful to visualize
energy changes throughout
a process on a reaction
coordinate diagram
It shows the energy of the
reactants and products (and,
therefore, ΔE).
The high point on the diagram
is the transition state.
• The species present at the
transition state is called the
activated complex.
• The energy gap between the
reactants and the activated
complex is the activation
energy barrier.

39. Maxwell–Boltzmann
Distributions
• 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.
As the temperature increases,
the curve flattens and
broadens.
• Thus at higher temperatures,
a larger population of
molecules has higher energy.

40. Maxwell–Boltzmann
Distributions
• 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.
• This fraction of molecules can be
found through the expression
where R is the gas constant and T is
the Kelvin temperature.
f = e−Ea/RT

41. Relationship Between k and Ea

42. Arrhenius Equaion
• Arrhenius Equation –
an equation that relates
the rate constant for a
reaction to the
frequency factor, A, the
activation energy, Ea
,
and the temperature. In
its logarithmic form it is
written (ln k = -Ea
/RT
+ lnA)

43. Reaction
Mechanisms

44. Reaction Mechanism
• Reaction Mechanism
– a detailed picture, or
model, of how the
reaction occurs; that is,
the order in which
bonds are broken and
formed and the
changes in relative
positions of the atoms
as the reaction
proceeds

45. Elementary Reactions
• Elementary Reactions – a process in a chemical
reaction that occurs in a single event or step. An
overall chemical reaction consists of one or more
elementary reactions or steps
Reactions may occur all at once or through several
discrete steps.
Each of these processes is known as an elementary
reaction or elementary process.

46. Molecularity
• Molecularity – the number of molecules that
participate as reactants in an elementary
reaction
The molecularity of a process tells how many
molecules are involved in the process.

47. Molecularity
• Unimolecular Reaction – an elementary
reaction that involves a single molecule
• Bimolecular Reaction – an elementary
reaction that involves two molecules
• Termolecular Reaction – an elementary
reaction that involves three molecules.
Termolecular reactions are rare.

48. Multistep Mechanisms
• Intermediate – a substance formed in one
elementary step of a multistep mechanism
and consumed in another

49. Multistep Mechanisms
• Rate-Determining Step – the slowest elementary step in a
reaction mechanism
In a multistep process, one of the steps will be slower than all others.
The overall reaction cannot occur faster than this slowest,
rate-determining step.
The sum of the elementary steps MUST give the overall balanced
equation for the reaction

50. Rate-Determining Step

51. Slow Initial Step
• The rate law for this reaction is found experimentally to be…
Rate = k [NO2
]2
CO is necessary for this reaction to occur, but the rate of the reaction does not
depend on its concentration.
This suggests the reaction occurs in two steps.
• A proposed mechanism for this reaction is
Step 1: NO2
+ NO2
⎯⎯→ NO3
+ NO (slow)
Step 2: NO3
+ CO ⎯⎯→ NO2
+ CO2
(fast)
• The NO3
intermediate is consumed in the second step.
• As CO is not involved in the slow, rate-determining step, it does not appear in the rate law.
NO2
(g) + CO (g) ⎯⎯→ NO (g) + CO2
(g)

52. Slow Initial Step

53. Fast Initial Step
• The rate law for this reaction is found to be
Rate = k [NO]2 [Br2]
Because termolecular processes are rare, this
rate law suggests a two-step mechanism.
• A proposed mechanism is…
Step 1 includes the forward and reverse
reactions.
2 NO (g) + Br2
(g) ⎯⎯→ 2 NOBr (g)
Step 2: NOBr2
+ NO ⎯⎯→ 2 NOBr (slow)
Step 1: NO + Br2
NOBr2
(fast)

54. Fast Initial Step
• The rate of the overall reaction depends upon the
rate of the slow step.
• The rate law for that step would be…
Rate = k2
[NOBr2
] [NO]
• But how can we find [NOBr2
]?
NOBr2
can react two ways:
• With NO to form NOBr
• By decomposition to reform NO and Br2
The reactants and products of the first step are in
equilibrium with each other.
Therefore,
Ratef
= Rater

55. Fast Initial Step
• Because Ratef = Rater ,
• k1
[NO] [Br2
] = k−1
[NOBr2
]
• Solving for [NOBr2
] gives us…
• Substituting this expression for [NOBr2
] in the
rate law for the rate-determining step gives
k1
k−1
[NO] [Br2
] = [NOBr2
]
k2
k1
k−1
Rate = [NO] [Br2
] [NO]
= k [NO]2
[Br2
]

56. Catalysis

57. Catalyst
• Catalyst – a substance
that changes the speed of
a chemical reaction without
itself undergoing a
permanent chemical
change in the process
Catalysts increase the rate of
a reaction by decreasing the
activation energy of the
reaction.
Catalysts change the
mechanism by which the
process occurs.

58. Types of Catalyst
• Homogenous
Catalyst – a catalyst
that is in the same
phase as the
reactant substances
• Heterogeneous
Catalyst – a catalyst
that is in a different
phase from that of
the reactant
substances

59. How Catalysts Work
• One way a catalyst can
speed up a reaction is
by holding the reactants
together and helping
bonds to break.
• In doing so, the catalyst
provides an alternate
mechanism for the
reaction with a lower
activation energy.

60. Adsorption
• Adsorption – the
binding of molecules
to a surface

61. Enzyme
• Enzyme – a protein
molecule that acts to
catalyze specific
biochemical reactions
The substrate fits into the
active site of the enzyme
much like a key fits into a
lock.

62. End of Unit
Be Prepared for the Exam.