1. Chapter II Chemical Kinetics

2. Lecture Outlines Introduction Factors affect reaction rate
Experimental determination of reaction rate
Computational determination of reaction rate (Mass Action Law)
Arrhenius Law
Orders of Reactions

3. Chemical 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).

4. Factors That Affect Reaction Rates
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. Factors That Affect Reaction Rates
Concentration of Reactants
As the concentration of reactants increases, so does the likelihood that reactant molecules will collide.

6. Factors That Affect Reaction Rates
Temperature
At higher temperatures, reactant molecules have more kinetic energy, move faster, and collide more often and with greater energy.

7. Factors That Affect Reaction Rates
Pressure
As the pressure of reactants increases, so does
the likelihood that reactant molecules will collide.

8. Factors That Affect Reaction Rates
Presence of a Catalyst
Catalysts speed up reactions by changing the mechanism of the reaction.

9. Classification of chemical reactions according to type of chemical substance
Gas phase reaction : Reactants of gaseous phases
Liquid phase reaction: Reactants of liquid phases
Solid phase reaction : Reactants of solid phases
Heterogeneous reaction : Reactants of different phases

10. Classification of chemical reactions according to speed of reaction
Explosive: dealing with determination of conditions under which chemical system undergoes very fast reaction and examination of the reaction mechanism
Non-explosive: steady reaction; like many pollutants are formed in reacting zones of rather reaction in various combustion system

11. Reaction Rates
Rate of reactions can be expressed in terms of the concentration ( rate of decrease of concentrations of reactants or the rate of increase of concentrations of products). Reaction rate units are moles/cm3.sec
Note : Concentration units are moles/cm3 and it is denoted by [ ] Concentration of A = [ A ]

12. Reaction Rates
Rates of reactions can be determined by monitoring the change in concentration of either reactants or products as a function of time.

13. Reaction Rates
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
mole/cm3
In this reaction, the concentration of butyl chloride, C4H9Cl, was measured at various times.

14. Reaction Rates
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
mole/cm3
mole/cm3.sec
The average rate of C4H9Cl over each interval is the change in concentration divided by the change in time:
Average rate =
[C4H9Cl] t

15. Reaction Rates C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
mole/cm3
mole/cm3.sec
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. 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.

17. Reaction Rates C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
All reactions slow down over time.
Therefore, the best indicator of the rate of a reaction is the instantaneous rate near the beginning.

18. 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]

19. Mass Action Law
A stoichiometric relation describing a one-step chemical reaction
Example

20. Mass Action Law
The law of mass action, which is confirmed experimentally, states that the rate of disappearance of a chemical species i , defined as RRi , is proportional to the product of the concentrations of the reacting chemical species, where each concentration is raised to a power equal to the corresponding stoichiometric coefficient; that is,
k is the proportionality constant called the specific reaction rate coefficient.

21. Mass Action Law

22. Rate Equation

23. Arrhenius Equation k = A e−Ea/RT where
It is a well-known fact that raising the temperature increases the reaction rate. Quantitatively this relationship between the rate a reaction proceeds and its temperature is determined by the Arrhenius Equation:
k = A e−Ea/RT
where
A : is the frequency factor, a number that represents the likelihood that collisions would occur with the proper orientation for reaction.
Ea: is the activation energy, the minimum energy required to start a chemical reaction (J/mole)
R: universal gas constant (8.314 J/mole.K)
T: Temperature (K)

24. Arrhenius Equation
The Arrhenius equation is often written in the logarithmic form:
y = ln k
m = - Ea / R
b = ln A
The activation energy Ea can be determined from the slope m of this line: Ea = - m · R . The value of the y-intercept corresponds to ln A.
Hint
An accurate determination of the activation energy requires at least three runs completed at different reaction temperatures. The temperature intervals should be at least 5 °C.

25. Arrhenius Equation Eq.(1) Eq.(2) Subtracting two equations
The activation energy can also be found algebraically by substituting two rate constants (k1, k2) and the two corresponding \temperatures (T1, T2) into the Arrhenius Equation
Eq.(1)
Eq.(2)
Subtracting two equations
Rearrangement

26. Zero Order of reaction rate [A]
A reaction is of zero order when the rate of reaction is independent of the concentration of materials. The rate of reaction is a constant.
k = rate, 0th order
[A]
rate
Rearranging
Integrating

27. First Order of reaction
For a general unimolecular reaction
where A is a reactant and P is a product, the reaction rate expression for a first order reaction is
Separation of the variables
Integrating

28. First Order of reaction
[A] = [A]o e – k t t½
This means that the concentration of A decreases exponentially as a function of time.
[A]=1/2[A]o
The rate constant k can also be determined from the half-life t1/2. Half-life time is the time it takes for the concentration to fall from [ A ]o to [ A ]o / 2 ==> [ A ] = 1 / 2 [ A ]o.

29. Second Order of reaction
For the general case of a reaction between A and B, such that
the rate of reaction will be given by :
Case 1 Initial concentrations of the two reactants are equal
Separating the variables and integrating
2nd order
1st order

30. Second Order of reaction
Case 2 Starting concentrations of the two reactants are different
If [ A ]o and [ B ]o are different the variable, x is used which represent the reacted part of A or B.
[ A ] = [ A ]o - x
[ B ] = [ B ]o - x
To integrate, perform a partial fraction expansion
d [ A ] = -dx