States of matter in chemistry related to intermediate students Easily understandable

1. CHEMISTRY
UNIT# 3
• STATES OF MATTER
–GASES
–LIQUIDS
–SOLIDS
(REVIEW LECTURE)

2. TODAY’S TOPICS
Properties of gases
Gas laws
Boyle’s law
Charles’s law
General gas equation
Kinetic molecular theory of gases
Ideal gases equation
Distinguish between Real and ideal gases (Van
der Waals equation)

3. Gas Law
Boyles Law Charles Law Avogadro’s
Law
Statement V ∝ 1/P V ∝ T (T in
Kelvin)
V ∝ n
Constant n, T n, P T, P
Variables V and P V and T (T in
Kelvin)
V and n
Expressions PV = K
P1V1 = P2V2
V/T = K V/n = K
2
2
1
1
T
V
T
V

2
2
1
1
n
V
n
V


4. Graphs of Boyles Law
Which of the following is Isotherm?
Which of the following is isothermal curve?
Which of the following is used to predict ideality of
gases?

5. General Gas Equation
V∝ T ; V ∝ 1/P ; V ∝ n
V ∝ nT/P
V = nRT/P
PV= nRT
Expression of Density:
RT
PM
d 
Units Of General Gas Constant
R and its value in S.I unit
When P is in atm and V in dm3
R = 0.0821 dm3 atm K-1 mol-1
When P is in mm Hg or torr
and V in dm3 or cm3.
R = 62.4 dm3 mm Hg K-1 mol-1
R = 62400 cm3 torr K-1 mol-1
When P is in Nm-2 and V in m3
(SI units)
R = 8.314NmK-1 mol-1
R = 8.314 JK-1 mol-1
R = 1.987 calK-1 mol-1

6. Postulates of KMT
Fundamental postulates of kinetic theory of gases:
1. Every gas consists of a large number of very small
particles called molecules. Gases like He, Ne, Ar, Kr,
Xe, Rn have monoatomic molecules.
2. The molecules of a gas move randomly, colliding
among themselves and with the walls of the container
and change their directions.
3. The pressure exerted by a gas is due to the collision of
its molecules with the walls of a container. The
collisions among the molecules are perfectly elastic.
4. The molecules of a gas are widely separated from one
another and there are sufficient empty spaces among
the molecules.

7. Postulates of KMT
5. The molecules of a gas have no forces of attraction
6. The actual volume of gas molecule is negligible as compared
to the volume of the Vessel.
7. The motion imparted to the molecules by gravity is
negligible as compared to the effect of continued collisions
between them.
8. The Kinetic energy of the gas molecules varies directly with
the absolute temperature of the gas. K.E  T
PV =1/3 mNc_2
M
RT
s
m
Cr
3
.
. 

8. Ideal and Non- Ideal Behavior of Gases
Gases are ideal at Low
Pressure and High
temperature
Compressibility factor
nRT
PV
Z 
Z = 1 Ideal Gas
Z > 1 Positive Deviation (Non Ideal)
Z < 1 Negative Deviation (Non Ideal)

9. Vander Waal’s Equation
• S.I units of a = Nm+4mol-2
• S.I unit of b = m3 mol-1
Mr
T
P
b
a
A
O
F
ideality
1
1
1
1
.
.
1






nRT
nb
V
V
an
P 

 )
)(
( 2
2

10. MCQ-1
Root mean square velocity of a gas is directly
related to
A) Pressure
B) Temperature
C) Molar mass
D) Volume

11. MCQ-2
Which of the following gas is least ideal
A) NH3
B) SO2
C) Cl2
D) He

12. MCQ-3
If temperature and pressure increase to
double then volume of gas
A. Increase
B. Remain constant
C. Decrease
D. Depends upon gas

13. TODAY’S TOPICS
Properties of liquids
Intermolecular forces
Hydrogen bonding
Evaporation
Vapor pressure
Boiling point and external pressure
Anomalous behavior of water

14. LIQUIDS
INTERMOLECULAR INTRAMOLECULAR
Forces b/w two different molecules Forces b/w atoms of same molecule
Weaker Stronger
Physical Properties Chemical Properties
Represented by ………. Represented by _______
H.Bond, dipole forces , debye
forces, London forces
Ionic Bond, Covalent bond
, Metallic bond

15. • Present in polar molecules
• DDF ∝ Electronegtivity
difference
• force in HCl or acetone
Dipole
forces
• Present in the mixture of polar
and non polar molecule
• DID ∝ size of non polar molecule
• DID ∝ polarity of polar molecule
Debye
forces

16. •Present in every compound
•Only force in non polar
compound
London
forces
Factors
affecting
•LDF ∝ size of molecules
•LDF ∝ Atomicity
•LDF ∝ 1/ Branches

17. •Special case of DDF
•H……FON
•F= Fluorine, O= oxygen, N=
nitrogen
Hydrogen
Bonding
Factors
affecting
•Strength of H.Bond ∝ ∆E.D
•Strength of H.Bonding ∝ no
of H-Bonds

18. STRUCTURE OF ICE:
0°C 4 °C 100°C
•Temperature
decreases
•Volume decreases
•Density increases
•Spaces decreases
•KE decreases
•Temp decreases
•Volume increases
•Density decreases
9%
•Spaces increases
9%
•KE decreases
•Water has maximum density
at 4 degree centigrade
•Structure of
ice= Hexagonal
•Structure of
H2O= tetrahedral

19. • spontaneous process
• Property of open system
• Carries at all T
• Endothermic
EVAPORATION
• Surface area
• Temperature
• Distance between molecules
Directly
affecting
factors
Inveresly
affecting
factors
• IMF
• Size of molecule

20. • Pressure when Rate of
evaporation becomes
equal to condensation
• Property of Closed system
VAPOR
PRESSURE
Depends
Upon
• V.P varies directly with
Temperature
• V.P ∝ Distance between molecules
• V.P ∝ 1/IMF

21. • When Vapor pressure
becomes equal to
external pressure
BOILING POINT
• B.Pt ∝ 1/ Distance Between
molecules
• B.P ∝ External Pressure
Depends
Upon
Does not
depend
• Surface Area
• Volume of container
• Amount of liquid
• Temperature

22. MCQ-4
If we provide a very large amount of
heat to liquid, its boiling point will
A. Increase
B. Remain constant
C. Decrease
D. There will no boiling

23. MCQ-5
All directly affect the boiling point
except
A. Inter–molecular forces
B. Surface area
C. Size of molecule
D. External Pressure

24. MCQ-6
Liquid hydrocarbon is
A. Methane
B. Propane
C. Ethane
D. Hexane

25. TODAY’S TOPICS
Introduction
Types of solids
Ionic solids
Molecular solids
Crystal lattice
Lattice energy

26. Comparison of Properties
Properties Ionic
solid
Covalent
solid
Molecular
solid
Metallic
solid
Smallest
particles
Cations ,
Anions
Atoms Molecules Cations,
free elec’s
Hardness Hard Hardest Softest Soft to V.
hard
Force
responsible
for being solid
Ionic
bond
Covalent
bond
Inter
molecular
forces
Metallic
bond

27. Melting
Point
High Highest Lowest High-low
Vapor
pressure
Low Lowest Highest High-low
Conductivity Molten
solution
X (Graphite) Absent Conductor
Solubility Polar Non-polar Both Acids or
bases
examples NaCl, MgO,
KCl
Diamond, BN
Graphite,
AlN, Silica
Iodine,
Sugar, dry
ice, S8, P4
Fe, Au, Ag

28. Example of ionic solid
Total sodium ion in unit cell of
NaCl= 4
Total formula unit in unit cell is = 4
Crystal system of NaCl is cubic
Coordination number = 6
Distance between similar
ions= 5.63Ao
Distance between
opposite ion=2.815Ao
Total Chloride ion in unit
cell of NaCl= 4

29. MCQ-7
Solid CO2 (Dry ice) is an example of _______
crystals
A. Covalent
B. Ionic
C. Metallic
D. Molecular

30. Which is incorrect about structure of iodine
A. Face-centered cubic
B. Metallic appearance
C. I – I (g) > I – I (s)
D. I – I (g) < I – I (s)
MCQ-8