chemical kinetics of chemical reaction

1. 1.

2. Chemical kinetics is a branch of chemistry that studies the rate of reaction and
the mechanism of the reaction. The main contents of chemical kinetics include
the following:
1) Determine the rate of the chemical reaction and the
influence of external factors such as temperature, pressure,
catalyst, solvent and light on the reaction rate;
2) Study the chemical reaction mechanism and reveal the
nature of the chemical reaction rate;
3) Explore the relationship and laws between material
structure and reaction ability.

3. What is reaction rate
The time taken for the disappearance of the reactant or the
appearance of the product . Rate is a ratio as the amount of reactant
disappeared divided by the time.
Average rate: The change in the concentration divided by the total
time elapsed.
Rate = amount reacted or produced/ time interval
units: g/s, mol/s, or %/s

4. Factors That Affect Reaction Rates
• The Nature of the Reactants
• Chemical compounds vary considerably in their chemical
reactivities.
• Concentration of Reactants
• As the concentration of reactants increases, so does the
likelihood that reactant molecules will collide.
• Temperature
• At higher temperatures, reactant molecules have more kinetic
energy, move faster, and collide more often and with greater
energy.
• Catalysts
• Change the rate of a reaction
by changing the mechanism.
4.

5. Reaction Rates
The rate of a chemical reaction can be determined by
monitoring the change in concentration of either
reactants or the products as a function of time.
[A] vs t
5.

6. Example 1: Reaction Rates
In this reaction, the
concentration of butyl
chloride, C4H9Cl, was
measured at various times,
t.
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
[C4H9Cl] M
Rate =
[C4H9Cl]
t
6.

7. 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.
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
7.

8. Reaction Rates
• The reaction slows down with
time because the concentration
of the reactants decreases.
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
8.

9. Reaction Rates and Stoichiometry
• 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(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
Rate =
-[C4H9Cl]
t
=
[C4H9OH]
t
9.

10. Reaction Rates & Stoichiometry
Suppose that the mole ratio is not 1:1?
10.
 H2 (g) + I2 (g)  2 HI (g)
 Disappearance of reactant
 Appearance of product
 2 moles of HI are produced for each mole of H2 used.
 The concentration of HI increases at 2x the rate of that I2
or H2 decreases
 For the overall rate to have the same value when defined
with respect to any reactant the change in HI must be
multiplied by 1/2

Rate = 
[H2]
t
or 
[I2]
t

Rate = 1/2
[HI]
t

11. • For 2N2O (g)  2 N2 (g) + O2 (g)
1.Express the rate in terms of change in concentration of
reactants and products.
2.In the first 15 sec., 0.015 mol of O2 is produced in a 0.500 L
reaction vessel. Determine the average rate of reaction.
3.What is [N2O] during this same time interval?
11
The Rate of a Chemical Reaction

12. Solution
• For 2N2O (g)  2 N2 (g) + O2 (g)
1.
2.
3.
12
The Rate of a Chemical Reaction

Rate = 
1
2
[N2O]
t
= 
1
2
[N2]
t
= 
[O2]
t

0.015 mol/0.500 L = 0.030 M Rate =
0.030 M
15 sec
 0.0020 M/s
Rate = 
1
2
[N2O]
t
 
[O2 ]
t
so
[N2O]
t
=  2
[O2 ]
t
=  0.0040 M/s
Δ [N2O] = - 0.004 ×15 sec. = - 0.06 M

13. Concentration and Rate
* Each reaction has its own equation that gives its rate as a function of reactant
concentrations.
- This is called its Rate Law
The general form of the rate law is
Rate = k[A]x[B]y
* K: is the rate constant
* [A] and [B] are the concentrations of the reactants.
* X and y are exponents known as rate orders that must be determined
experimentally
To determine the rate law we measure the rate at different starting concentrations.
13.

14. Concentration and Rate
Compare Experiments 1 and 2:
when [NH4
+] doubles, the initial rate doubles.
NH4
+ (aq) + NO2
- (aq)  N2 (g) + 2H2O (l)
14.

15. Concentration and Rate
Likewise, compare Experiments 5 and 6:
when [NO2
-] doubles, the initial rate doubles.
NH4
+ (aq) + NO2
- (aq)  N2 (g) + 2H2O (l)
15.

16. Concentration and Rate
This equation is called the rate law, and k is the
rate constant.
NH4
+ (aq) + NO2
- (aq)  N2 (g) + 2H2O (l)
16.

17. Measuring Reaction Rates
• Spectroscopy
• Light is passed thru a sample
• Absorption of light is proportional to concentration.
• More light absorbed = higher concentration.
17
The Rate of a Chemical Reaction

18. Measuring Reaction Rates
• Pressure Measurement
• Concentrations of gaseous reactants can be monitored by
measuring changes in pressure
18
The Rate of a Chemical Reaction
For gas-phase reactants use PA instead of [A].

19. • Rates of reaction often depend on the concentration of one or more
reactants
• Consider a simple reaction A  products
• We can express this relationship in a Rate Law
• Rate = k [A]n
• Where k is a proportionality constant called a rate constant and n is the
reaction order
• The value of n determines how the rate depends on the concentration of the
reactant
19
The Rate Law
The Effect of Concentration on Reaction Rate

20. • Rate = k [A]n
• The value of n determines how the rate depends on the
concentration of the reactant
20
The Rate Law
The Effect of Concentration on Reaction Rate
Value of n Reaction order rate
0 zero Independent of the [A]
1 first Directly proportional to [A]
2 second Proportional to [A]2

21. The Rate Law
• Exponents tell the order of the reaction with respect to each reactant.
• This reaction is
First-order in [NH4
+]
First-order in [NO2
−]
• The overall reaction order can be found by adding the exponents on the reactants
in the rate law.
• This reaction is second-order overall.
Rate = K [NH4
+ ]1[NO2
- ]1
21.

22. Zero, First and Second-Order Reactions
• We can examine reaction order by looking at:
• [reactant] vs time
• The slope of this line = rate  [R] / t
22
The Rate Law: The Effect of Concentration on Reaction Rate

23. 23
In this figure:Rate vs [reactant]. How does [R] affect
rate?

24. Zero-Order Reaction
• Rate = k[A]0 = k
• Independent of reactant concentration
• [Reactant] decreases linearly with time
• Constant reaction rate
• Reaction does not slow down as reaction progresses
• Rate is the same at any [reactant]
24
The Rate Law: The Effect of Concentration on Reaction Rate

25. 25

26. Zero-Order Reaction
• Example of a zero-order reaction is decomposition of NH3 in presence of
molybdenum or tungsten catalyst
[Mo]
• 2NH3 (g) N2 (g) + 3H2 (g)
• The surface of the catalyst is almost completely covered by NH3 molecules. The
adsorption of gas on the surface cannot change by increasing the pressure or
concentration of NH3. Thus, the concentration of gas phase remains constant
although the product is formed.
26
The Rate Law: The Effect of Concentration on Reaction Rate

27. First-Order Reaction
27
The Rate Law: The Effect of Concentration on Reaction Rate
• Rate = k[A]1 = k
• Directly proportional to [reactant]
• [Reactant] decreases with time
• Not linear – slope becomes less steep
• Reaction slows down as reaction progresses
• Rate decreases
• Rate is directly proportional to [reactant]
• linear

28. 28

29. Second-Order Reaction
29
The Rate Law: The Effect of Concentration on Reaction Rate
• Rate = k[A]2 Rate is proportional to square [reactant]
• Proportional to square of [reactant]
• [Reactant] decreases with time
• Not linear – slope becomes less steep
• Reaction slows down as reaction progresses and the rate decreases

30. Determining Reaction Order (method of initial rates)
• Reaction order can only be determined by experiment.
• It cannot be deduced from the balanced equation.
• Remember – reaction order is how the rate depends on
reactant concentration.
• So how do we determine this?
30

31. Determining Reaction Order (method of initial rates)
31
• This set of data shows no effect on rate with
increasing [A].
• Rate is independent of [A].
• Zero order
• Rate = k

32. Determining Reaction Order (method of initial rates)
32
• Measure the (initial) rate of reaction at several different [reactant].
• Look for the effect of [reactant] on rate.
• In this set of data when [A] doubles the rate doubles
• Rate is directly proportional to [A]
• First order Rate = k [A]
• In this set of data when [A] doubles the rate quadruples.
• Rate is directly roportional to [A]2
• Second order Rate = k [A]2

33. Determining Reaction Order (method of initial rates)
• The math on that last one was obvious but this might not always be the
case.
• May need to actually solve for the value of n to determine the rate law.
33
rate 2
rate 1

k [A]2
n
k[A]1
n =
0.240 M/s
0.060 M/s
=
k(0.40)n
k(0.20)n
4.0 =
0.40
0.20




n
= 2n
log 4.0 = log 2n
= n log 2
n =
log 4
log 2
= 2

34. Other Factors that influence rate of
reaction
Temperature: An increase in temperature is accompanied by an
increase in the reaction rate. Temperature is a measure of the
kinetic energy of a system, so higher temperature means higher
average kinetic energy of molecules and more collisions per unit
time.
For most chemical reactions the rate at which the reaction
proceeds will approximately double for each 10°C increase in
temperature. Once the temperature reaches a certain point, some
of the chemical species may be altered (e.g., denaturing of
proteins) and the chemical reaction will slow or stop.

36. • Medium: The rate of a chemical reaction depends on the medium in
which the reaction occurs. It sometimes could make a difference
whether a medium is aqueous or organic; polar or nonpolar; or liquid,
solid, or gaseous.
• Surface area: It is easier to dissolve sugar if it is crushed. Crushing the
sugar increases its surface tension. The larger surface area allows more
sugar molecules to contact the solution.

37. Catalyst: A catalyst is a substance that alters the rate of a chemical
reaction without being used up or permanently changed chemically.
A catalyst works by changing the energy pathway for a chemical
reaction. It provides an alternative route (mechanism) that lowers the
Activation Energy meaning more particles now have the required
energy needed to undergo a successful collision.

38. • What is activation energy?
• The least amount of energy needed to permit a particular chemical
reaction.

39. Types of catalysts
There are 2 types of catalysts:
Homogeneous catalyst: Homogeneous catalysts are in the same phase as the
reactants.
Ex: 2H2O2(aq)+ KI(aq)2H2O(l)+O2(g)
Heterogeneous catalyst: Heterogeneous catalysts are present in different
phases from the reactants (for example, a solid catalyst in a liquid reaction
mixture), whereas homogeneous catalysts are in the same phase (for
example, a dissolved catalyst in a liquid reaction mixture).
Ex: Decomposition of H2O2 in presence of MnO2
• Hydrogen peroxide is a solution while manganese dioxide is a solid and
can be easily separated.

40. Thank you for your
attention