1. Physical Chemistry 1 CHEM 3310

2. Chem 3310 Credit hrs: 4 (3+1) Prerequisite: chem Level: 5th Class Time: Thursday Lecture 10:00 AM-12:30 PM Monday Lab. 12:00 – 2:00 PM Office Hours: Sunday 10:00- 12:00 PM, Monday 10:00 -12:00 PM Text book: Atkin’s, Except First Lecture from General Chemistry book, Chang 3rd Edition.

3. توزيع الدرجات 100 points Class Work الاعمال الفصلية 60 points 2 Exams
2 x 10= 20 points
Final Exam
الاختبار النهائي
40 points
Oral Exams+ Quizzes
2 x 10 = 20 points

4. CHEM 3310 Course Plan Week # Topic 1,2
Gases & Kinetic Molecular Theory (KMT)
3,4
First Law of Thermodynamics & Thermochemistry
5,6
Second & Third Laws of Thermodynamic 7
Chemical Equilibrium 8
Phases & Solutions
9,10
Phase Equlibria 11
Solutions of Electrolytes 12
Electrochemical Cells

5. The Four Postulates of the Kinetic Theory
A pure gas consists of a large number of identical molecules separated by distances that are great compared with their size
( sizes of molecules are negligible comparing to distances).
The gas molecules are in constant motion and they collide to one another. And collide to walls of container. Collisions among molecules are elastic.
Pressure =
Force
Area

6. The Four Postulates of the Kinetic Theory
The molecules exert neither attractive nor repulsive on one another. No energy lost.
The average kinetic energy of molecules is proportional to temperature of the gas in kelvin. Any two gases at the same temperature will have the same average kinetic energy. Where m is the mass, u is molecular speed.
KE = ½ mu2

7. Application of Kinetic Molecular Theory (KMT) to the Gas Laws
Compressibility of Gases:
Gases are compressed hence molecules are separated by large distances  compressed easily.

8. Boyle’s Law Experimentally P a 1/V KMT explained Boyle’s Law as:
Collisions of gas molecules with walls of container causes Pressure.
Collision rate is (the number of molecular collisions with walls per second).
Volume of container decreases the number density (number of molecules per unit volume of the gas) increases- collision rate increases. pressure of a gas is inversely proportional to volume that gas occupies.
KMT explained Boyle’s Law as:
P a collision rate with wall of container.
Collision rate a number density
Number density a 1/V
P a 1/V

9. Charles’ Law Experimentally P a T KMT explained Charles’ Law as:
Average kinetic energy is proportional to absolute temperature K. Consequently number of collisions of gas molecules increases , and thus the pressure increases.
KMT explained Charles’ Law as:
P a collision rate with wall
Collision rate a average kinetic energy of gas molecules
Average kinetic energy a T
P a T

10. Avogadro’s Law P a n ( number of moles)
KMT explained Avogadro’s Law as:
P a collision rate with wall
Collision rate a number density
Number density a n
P a n

11. Experimentally Ptotal = SPi
Dalton’s Law of Partial Pressures
Experimentally Ptotal = SPi
Consider a case in which two gases, A and B, are in a container of volume
Ptotal = P1 + P2
Ptotal = Spi
KMT explained Dalton’s Law of Partial Pressure as:
Molecules do not attract or repel one another
P exerted by one type of molecule is unaffected by the presence of another gas
Ptotal = SPi

12. Ideal Gas Law 1 Boyle’s law: V a (at constant n and T) P
Charles’ law: V a T (at constant n and P)
Avogadro’s law: V a n (at constant P and T)
V a nT P
V = constant x = R nT P
R is the gas constant
PV = nRT

13. Distribution of Molecular Speeds
Maxwell analyzed the behavior of gas molecules at different temperature.
Distribution curves tells us:
Peak of each curve represent the most probable speed( speed of largest number of molecules)
Most probable speed increases as temp. increases (i.e. shifts toward right)
Curve flatten out with increasing temp. indicates larger number of molecules are moving at greater speed.
Maxwell speed distribution curves for three different temps.

14. Distribution of Molecular Speeds
Lighter gas molecules move faster
than heavier gases.
He is faster than N2,
N2 is faster than Cl2
The distribution of speeds of three different gases at the same temperature
urms =
3RT M 
M: molar mass Heavier gas is more slowly than lighter gas.
R = (8.314 J/mol K),
T is the temperature in kelvin

15. Example
Calculate the root mean- square speeds of helium atoms and nitrogen N2 molecules in m/s at 25oC.
Solution
R= J/K. mol, OR
R= L atm K−1 mol−1
3. Molar mass should be Kg/mol

16. Solution
R= J/K. mol, R= L atm K−1 mol−1, Molar mass = 4 g/mol
Molar mass He = 4 x 10-3 Kg/mol

17. For Nitrogen
Molar mass = 28.02g/ mol
Molar mass N2= x10-2 Kg/mol

18. Deviation of Real Gas Behavior
For 1 mole of ideal gas at any pressure and temperature.
PV/RT = 1
For real gases. PV/RT ≠ 1
As pressure increases and temperature decreases, real gases deviate from ideal gas behavior,
The value ≠ (1)

19. Why does real gas behavior deviate from ideal gas behavior?
Kinetic Theory of Gases postulates:
1- Volume of gas molecules ignored relative to the total volume of gas occupies the whole container.
It is applicable at low pressure i.e. volume is big.
It in inapplicable at high pressure i.e. volume is small

20. Volume of gas molecules can be ignored
Behavior of Ideal Gas
Low pressure
High Pressure
Volume of
gas molecules can’t be ignored.
Volume of gas molecules can be ignored
V ideal = volume of distances separate molecules
V measured = volume of distances separate molecules + volume of particles

21. Van Der Waals’ Correction of Volume
V ideal = V measured - nb
Where:
n: number of moles
b: constant related to molecular volume

22. Behavior of Real Gas
Kinetic Theory of Gases postulates:
2- Pressure on the walls of a container is caused by molecular collisions with walls of container.
No attractive forces between molecules
It is applicable at high temperature
It in inapplicable low temperature
Why????

23. Behavior of Real Gas
At High Temperature: , molecules have high kinetic energy an move very fast, so potential energy due to attractive forces between molecules can be ignored.
At low temperature, molecules move very slow because they have small kinetic energy, therefore collisions with container wall decrease

24. Van Der Waals’ Correction of Pressure
Van Der Waals’ Equation for Real Gases

25. High Temperature and Low Pressure.
Behavior of Real Gas
To make real gases have similar behavior to the ideal gases. They should be under
High Temperature and Low Pressure.

26. Real Gases: Deviations from Ideality
PROBLEM:
Given that 3.5 moles of NH3 occupy 5.2 L at 47oC, Calculate the pressure of the gas in (atm) using
1. The ideal gas equation
2. The van der Waals equation, where a= 4.17 atm
L2/mol2, b= L/mol

27. P = (3.5 mol)(0.08206 L atm mol-1 K-1)(320K) (5.2 L) P = 17.7 atm
Solution
Ideal Gas Equation
PV = nRT
P = nRT/V V= 5.2 L n = 3.5 mol
T = = 320 K
P = (3.5 mol)( L atm mol-1 K-1)(320K)
(5.2 L)
P = 17.7 atm
R= L. atm/K.mol

28. 2) Van der Waals Equation
Lowering the actual gas indicates attraction forces between NH3 gas molecules results in a lower gas pressure than that of an ideal gas under the same condition.