Physical Chemistry 1 CHEM 3310. Chem 3310 Credit hrs: 4 (3+1) Prerequisite: chem 2020 Level: 5 th...
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Transcript of Physical Chemistry 1 CHEM 3310. Chem 3310 Credit hrs: 4 (3+1) Prerequisite: chem 2020 Level: 5 th...
Physical Chemistry 1CHEM 3310
Chem 3310Credit hrs: 4 (3+1)Prerequisite: chem 2020 Level: 5th
Class Time: Thursday Lecture 10:00 AM-12:30 PM Monday Lab. 12:00 – 2:00 PMOffice Hours: Sunday 10:00- 12:00 PM, Monday 10:00 -12:00 PMText book: Atkin’s, Except First Lecture from General Chemistry book, Chang 3rd Edition.
الدرجات توزيع
100 points
Class Workاالعمال الفصلية
60 points 2 Exams 2 x 10= 20 points
Final Examالنهائي االختبار
40 points Oral Exams+ Quizzes
2 x 10 = 20 points
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
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
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.
The Four Postulates of the Kinetic Theory
KE = ½ mu2
Compressibility of Gases:• Gases are compressed hence molecules are
separated by large distances compressed easily.
Application of Kinetic Molecular Theory (KMT) to the Gas Laws
Experimentally P a 1/V
• 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 densityNumber density a 1/VP a 1/V
Boyle’s Law
Experimentally P a T
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 wallCollision rate a average kinetic energy of gas moleculesAverage kinetic energy a TP a T
Charles’ 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
Avogadro’s Law
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 anotherP exerted by one type of molecule is unaffected by the
presence of another gasPtotal = SPi
Dalton’s Law of Partial Pressures
Charles’ law: V a T (at constant n and P)
Avogadro’s law: V a n (at constant P and T)
Boyle’s law: V a (at constant n and T)1P
V a nT
P
V = constant x = RnT
P
nT
PR is the gas constant
PV = nRT
Ideal Gas Law
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 N2 @ three different temps.
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 = 3RTM
M: molar mass Heavier gas is more slowly than lighter gas. R = (8.314 J/mol K), T is the temperature in kelvin
Distribution of Molecular Speeds
Example
Calculate the root mean- square speeds of helium atoms and nitrogen N2 molecules in m/s at 25oC.
Solution1. R= 8.314 J/K. mol, OR2. R= 0.082 L atm K−1 mol−1
3. Molar mass should be Kg/mol
Solution
R= 8.314 J/K. mol, R= 0.082 L atm K−1 mol−1, Molar mass = 4 g/molMolar mass He = 4 x 10-3 Kg/mol
smxu
smxu
skgkgmxu
SKgmJ
KgJxx
xxu
/1036.1
/1086.1
./1086.1
/11
/1086.1104
298314.83
3
226
226
22
63
M
RTurms
3
For Nitrogen
smu
smxu
molkgx
KmolKJu
/515
/1065.2
/10802.2
)298)(./314.8(3
225
2
Molar mass = 28.02g/ mol Molar mass N2= 2.802 x10-2 Kg/mol
Deviation of Real Gas Behavior
For 1 mole of ideal gas at any pressure and temperature.
PV/RT = 1
18
•As pressure increases and temperature decreases, real gases deviate from ideal gas behavior, • The value ≠ (1)
•For real gases. PV/RT ≠ 1
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
Behavior of Ideal Gas
Volume of gas molecules can’t be ignored.
Volume of gas molecules can be
ignored
Low pressure
High Pressure
V ideal = volume of distances separate moleculesV measured = volume of distances separate molecules + volume of particles
V ideal = V measured - nb
Van Der Waals’ Correction of Volume
Where:n: number of molesb: constant related to molecular volume
21
Behavior of Real Gas
2- Pressure on the walls of a container is caused by molecular collisions with walls of container.
No attractive forces between molecules
Why????
22
Kinetic Theory of Gases postulates:
It is applicable at high temperatureIt in inapplicable low temperature
23
•At High Temperature: , molecules have high kinetic energy an move very fast, so potential energy due to attractive forces between molecules can be ignored.
Behavior of Real Gas
•At low temperature, molecules move very slow because they have small kinetic energy, therefore collisions with container wall decrease
Van Der Waals’ Equation for Real Gases
24
Van Der Waals’ Correction of Pressure
• To make real gases have similar behavior to the ideal gases. They should be under
High Temperature and Low Pressure.
Behavior of Real Gas
25
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= 0.0371 L/mol
Real Gases: Deviations from Ideality
Solution
1) Ideal Gas Equation
PV = nRT
P = nRT/V V= 5.2 L n = 3.5 mol
T = 47 + 273= 320 K
P = (3.5 mol)(0.08206 L atm mol-1 K-1)(320K)
(5.2 L)
P = 17.7 atm
R= 0.0821 L. atm/K.mol
P + n a
VV nb nRT
2
2
atmP
KmolKatmLmolLLatmP
LmolLmolnb
atmL
molmolatmL
V
an
2.16
)320)(./.0821.0)(5.3()130.02.5)(89.1(
130.0)/0371.0)(5.3(
89.1)2.5(
)5.3)(/17.4(2
222
2
2
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.
2) Van der Waals Equation