1. Physical Chemistry II
CHEM 3320

2. List of Topics
No. of
Weeks
Contact Hours
Kinetics of elementary reactions 3 11
Composite reaction mechanisms
Solution of electrolytes &
Debye-Huckel theory
Electrochemical cells 2 8
Surface chemistry
Transport properties 1
A comprehensive review

3. Proportion of Total Assessment By the end of the semester
Assessment task
Week Due
Proportion of Total Assessment
Mid term I
6th week
10%
Mid term II
11th week
Participation and attendance 5%
Quizzes
presentation
Assignment
From 2nd to 14th week
Final practical exam
15th week
20%
Final written Exam.
By the end of the semester
40%

4. List Required Textbooks
"Physical Chemistry", 4th edition, by kieth Laidler, Meisser and Sanctury.
"Physical Chemistry", 8th Edition, By Peter Atkins and Julio de Paula.

5. Chapter I Chemical reaction Mechanism 1
1. Rate law and order of reaction
2. Differential rate laws
3. integrated rate laws
4.Molecularity
5.Reaction mechanisms and elementary processes
6. Collision Theory, activation energy
7. Arrhenius Equation
8. Activated Complex Theory

6. Factors Affecting Reaction Rate
UNIT 3
Chapter 6: Rates of Reaction
Section 6.2
Factors Affecting Reaction Rate
1. Nature of reactants
Ions faster than molecules
2. Concentration
Higher concentration a greater number of effective collisions
3. Temperature
Higher temperature, higher sufficient energy needed for a reaction (energy is ≥ Ea)

7. Factors Affecting Reaction Rate
UNIT 3
Chapter 6: Rates of Reaction
Section 6.2
Factors Affecting Reaction Rate
4. Pressure for gases
increased pressure , increased the number of collisions .
5. Surface area
a greater surface area of solid reactant  a greater chance of effective collisions
6. Presence of a catalyst
a catalyst is a substance that increases a reaction rate without being consumed by the reaction

8. The Rate Law rate = k[A]m[B]n UNIT 3
Chapter 6: Rates of Reaction
Section 6.3
The Rate Law
The rate law shows the relationship between reaction rates k and concentration of reactants [] for the overall reaction.
rate = k[A]m[B]n
m: order of the reaction for reactant A
n: order of the reaction for reactant B
k: rate constant
overall order of the reaction : m + n

9. First-order Reactions
UNIT 3
Chapter 6: Rates of Reaction
Section 6.3
First-order Reactions
A  Product
the rate law equation:
rate = k [A]
k: rate constant, [A]: concentration of reactant A

10. Second-order Reactions
UNIT 3
Chapter 6: Rates of Reaction
Section 6.3
Second-order Reactions
A + B  Product
the rate law equation:
rate = k [A][B] OR
rate = k [A]2
Rate of reaction is first order with respect to reactant A
Rate of reaction is first order with respect to reactant B
Overall order of chemical reaction is 2nd order reaction

11. Reaction Mechanisms
Reaction Mechanism: explains how the overall reaction proceeds.
Reaction Mechanisms: is the sequence of several steps that describes the actual process by which reactants become products.
Each of these step is known as an elementary reaction or elementary process.

12. N2O2 is detected during the reaction!
Reaction Mechanisms
Elementary step: any process that occurs in a single step
For example, oxygen and nitrogen are not formed directly from the decomposition of nitrogen dioxide:
2NO (g) + O2 (g) NO2 (g)
N2O2 is detected during the reaction!
Elementary step 1:
NO + NO N2O2
Overall reaction:
2NO + O NO2 +
Elementary step 2:
N2O2 + O NO2

13. Reaction Mechanisms For Example: A  B
Now we will examine what path the reactants took in order to become the products.
The reaction mechanism gives the path of the reaction.
Mechanisms provide a very detailed picture of which bonds are broken and formed during the course of a reaction.
A  B

14. Elementary Steps & Molecularity
Molecularity: number of molecules present in an elementary step.
Unimolecular: one molecule in the elementary step,
Bimolecular: two molecules in the elementary step, and
Termolecular: three molecules in the elementary step.
(It is uncommon to see termolecular Processes…statistically improbable for an
effective collision to occur.)

15. Rate Laws and Molecularity
rate = k [A]
Unimolecular reaction
A  products
Bimolecular reaction
A + B  products
rate = k [A][B]
Bimolecular reaction
A + A  products
rate = k [A]2

16. Rate Laws of Elementary Steps
Since this process occurs in one single step, the stoichiometry can be used to determine the rate law!
The rate law for an elementary step is written
directly from that step

17. Collision Theory
Molecules of reactants must collide each other . Not all collisions are effective (i.e. leads to chemical reaction). Conditions of occurring chemical reaction according to collision theory:

18. Effective Collision Criteria
1.The Correct Orientation of Reactants
For a chemical reaction to occur, reactant molecules must collide with the correct orientation relative to each other, which called (collision geometry).
Five of many possible ways that NO(g) can collide with NO3(g) are shown. Only one has the correct collision geometry for reaction to occur.

19. Effective Collision 2.Sufficient Activation Energy:
For a chemical reaction, reactant molecules must also collide with sufficient energy.
Activation energy, Ea, is the minimum amount of collision energy required to initiate a chemical reaction.
Collision energy depends on the kinetic energy of the colliding particles.

20. Potential Energy Hill
Ea: is the minimum energy that reactants must have to form products.
the height of the potential barrier  (sometimes called the energy barrier). 
Activation Energy
Curve called :
1. Potential Energy Hill Or
2. Potential Energy Barrier

21. Activation energy, Ea
The shaded part of the Maxwell-Boltzmann distribution curve represents number of particles ( i.e. number of collisions) that have enough collision energy for a reaction (i.e. the energy is ≥ Ea).
Suppose: Number of collisions 100
Where, Ea= 70 j/mole
25 collisions have 20 j/mole
40 collisions have 45 j/mole
20 collisions have 50 j/mole
10 collisions have 60 j/mole
5 collisions have 70 j/mole
Number of collisions
Energy

22. 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.

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

24. Maxwell–Boltzmann Distributions
If the dotted line represents the activation energy, as the temperature increases, so does the number of molecules ( i.e. number of collisions) that can overcome the activation energy barrier.
As a result, the reaction rate increases.

25. Arrhenius Equation
A mathematical relationship between k : ( rate constant of chemical reaction) and Ea: activation energy.
where
A : “Frequency Factor”-- a constant indicating how many collisions have the correct orientation to form products.

26. Arrhenius Equation: Temperature Dependence of the Rate Constant
Ea = the activation energy (J/mol)
R = the gas constant (8.314 J/K•mol)
T = is the absolute temperature ( in Kelvin)
A = is the frequency factor

27. Arrhenius Equation
Taking the natural logarithm (ln) of both sides, the equation becomes,
Ln (natural logarithm): is inverse function of exponential function.
eln(x) = x
ln(ex) = x

28. ln(k) = - Ea/R(1/T) + ln(A)
Arrhenius Equation
ln(k) = - Ea/R(1/T) + ln(A)
Straight Line Equation
y = mx + b
When k is determined experimentally at several temperatures, Ea can be calculated from the slope of a plot , Slope= -Ea/R

29. Exothermic Reaction A → B + Heat
Potential Energy of reactant = Energy of chemical bond = Heat content = H
H product < H reactant
Enthalpy (∆H)
∆H = H product - H reactant < 0
∆H = negative value (-)
Activation Energy (Ea) = H transition state- H reactant
Exothermic reaction has Low Ea
Exothermic

30. Endothermic Reaction A + Heat → B Endothermic reaction has High Ea
Potential Energy of reactant = Energy of chemical bond = Heat content = H
H product > H reactant
Enthalpy (∆H)
∆H = H product - H reactant > 0
∆H = positive value (+)
Activation Energy (Ea) = H transition state- H reactant
Endothermic reaction has High Ea
Endothermic

31. Activation Energy and Enthalpy
The Ea for a reaction cannot be predicted from ∆H.
∆H is determined only by the difference in potential
energy between reactants and products.
△H has no effect on the rate of reaction.
The rate depends on the size of the activation energy Ea
Reactions with low Ea occur quickly. Reactions with
high Ea occur slowly.
Potential energy diagram for the combustion of octane.

32. Activation Energy for Reversible Reactions
Potential energy diagrams  both forward and reverse reactions.
follow left to right for the forward reaction
follow right to left for the reverse reaction

33. Activated Complex (Transition State)
Activated complex is unstable compound and can break to form product.
Activated complex: The arrangement of atoms found at the top of potential energy hill or barrier.

34. Activated Complex (Transition State)
1. The collision must provide at least the minimum energy necessary to produce the activated complex. 2. It takes energy to initiate the reaction by converting the reactants into the activated complex. 3.If the collision does not provide this energy, products cannot form.

35. Analyzing Reactions Using Potential Energy Diagrams
Forward Reaction is Exothermic Reaction
Reversible Reaction is Endothermic Reaction
BrCH3 molecule and OH- must collide with the correct orientation and sufficient energy and an activated complex forms.
2. When chemical bonds reform,
potential energy decreases and
kinetic energy increases as the
particles move apart.
Ea(rev) is greater than Ea(fwd)