The document discusses bonding and intermolecular forces in covalent molecular substances. It can be summarized as follows: 1. Bonding in covalent molecular substances involves strong intramolecular forces within molecules and weaker intermolecular forces between molecules. 2. The main types of intermolecular forces are dipole-dipole attractions, hydrogen bonding, and London dispersion forces. Hydrogen bonding is the strongest and London dispersion forces are the weakest. 3. The kinetic molecular theory describes gases as small particles in constant random motion, with repulsive and attractive intermolecular forces between particles. The theory relates properties like pressure, volume, temperature, and amount of gas.

1. GAS AND
CONDENSED
MATTER

2. Bonding model for covalent
molecular substances

1.
2.
Bonding for covalent molecular substances falls into
two categories
The strong forces of attraction which holds atoms
together within molecules
The weak forces of attraction between molecules

3. Forces between molecules
(intermolecular forces)
will learn about the forces between molecules
or compounds are called intermolecular forces
 Inter means between or among
 Internet, interstate, international
 What would Interstellar travel be?
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4. Intermolecular Forces

Forces that occur between
molecules.

5. Intramolecular forces
 What
would intramolecular forces be?
 Forces within molecules e.g covalent, metallic or
ionic
 intra means within
 Intrastate, intranet, intracellular
 Most of the intermolecular forces we look at occur
between covalently bonded molecules or covalent
molecular substances
 Intramolecular bonds are stronger than
intermolecular forces.

6. Overview
 All
matter is held together by force.
 The force between atoms within a
molecule is a chemical or intramolecular
force.
 The force between molecules are a
physical or intermolecular force.
 These physical forces are what we
overcome when a chemical changes its
state (e.g. gas liquid).

7. What causes intermolecular forces?
 Molecules
are made up of charged particles: positive
nuclei and negative electrons.
 When one molecule approaches another there is a
multitude of forces between the particles in the two
molecules.
 Each electron in one molecule is attracted to the nuclei in
the other molecule but also repelled by the electrons
in the other molecule.
 The same applies for nuclei

8. Types of Intermolecular forces
 The
three main types of intermolecular
forces are:
1. Dipole-dipole attraction
occur only
btw polar molecules
2. H bonding – only with Hydrogen and
Oxygen, Fluorine and Nitrogen)
3. Dispersion forces (London Dispersion
Forces)

9. Intermolecular Forces

Forces that occur between
molecules.

Dipole–dipole forces


Hydrogen bonding
London dispersion forces

10. Dipole–Dipole Attraction

11. Dipole-Dipole Forces
 Dipole
moment – molecules with polar
bonds often behave in an electric field as if
they had a center of positive charge and a
center of negative charge.
 Molecules with dipole moments can attract
each other electrostatically. They line up so
that the positive and negative ends are
close to each other.
 Only about 1% as strong as covalent or ionic
bonds.

12. Hydrogen Bonding
Strong
dipole-dipole forces.
Hydrogen is bound to a highly
electronegative atom –
nitrogen, oxygen, or fluorine.

13. Hydrogen Bonding in Water

Blue dotted
lines are the
intermolecula
r forces
between the
water
molecules.

14. Hydrogen Bonding
Affects

physical properties
Boiling point

15. London Dispersion Forces



Instantaneous dipole that occurs
accidentally in a given atom induces a
similar dipole in a neighboring atom.
Significant in large atoms/molecules.
Occurs in all molecules, including
nonpolar ones.

16. London Dispersion Forces
Nonpolar Molecules

17. London Dispersion Forces

Become
stronger as the
sizes of atoms
or molecules
increase.

18. Melting and Boiling Points

In general, the stronger the
intermolecular forces, the higher
the melting and boiling points.

19. Strength of Intermolecular
Interactions
Hydrogen Bonding
↑
Dipole – Dipole
↑
London Dispersion Forces

20. Kinetic Molecular Theory
The kinetic theory of matter is based on the following
postulates:
1. Matter is composed of small particles called
molecules
2. The particles are in constant random motion
They possess kinetic energy due to their motion
3. There are repulsive and attractive forces between
particles.
They posses potential energy due to these
forces
4. Average particle speed increases with temperature
5. No energy is lost when the particles collide, called
elastic collision

21. Kinetic Molecular Theory
The kinetic energy of a particle is given by the equation:
1
KE = mv 2
Where:
2
m = particle mass in kg
v = particle velocity in m/s
KE = kg-m2/s2 = j (joule)
According to postulate 4 of our kinetic theory particle velocity
increases with temperature. This means as temperature
increases then kinetic energy increases.

22. Potential Energy
Potential energy is the sum of the attractive and
repulsive forces between particles.
Examples of these types of forces are the
gravitational attractive forces between objects and
the repulsive forces between the same poles of
magnets.
Alternatively we can say forces between particles may
be either cohesive or disruptive.

23. Interparticle Forces
Cohesive forces include dipole-dipole interactions,
dispersion forces, attraction between oppositely
charged ions.
Cohesive forces are largely temperature
independent.
e.g. magnets and gravity function the same way at
different temperature.

24. Interparticle Forces
Disruptive forces are those forces that make
particles move away from each other.
These forces result predominately from the particle
motion.
Disruptive forces increase with temperature in
agreement with postulate 4.
We can conclude that as we increase the
temperature particles will become further apart from
each other.

25. Tutorial 1
 State
and describe briefly three (3) main types of
intermolecular forces.
 State five (5) assumption in the kinetic molecular
theory

26. A Gas
 Has
neither a definite volume nor
shape.
 Uniformly
fills any container.
 Mixes
completely with any other gas
 Exerts
pressure on its surroundings.

27. Earth-like Atmosphere
Composition of Earth’s Atmosphere
Compound
%(Volume)
Mole Fractiona
Nitrogen
Oxygen
Argon
Carbon dioxide
78.08
20.95
0.934
0.033
0.7808
0.2095
0.00934
0.00033
Methane
Hydrogen
2 x 10-4
5 x 10-5
2 x 10-6
5 x 10-7
a. mole fraction = mol component/total mol in mixture.

28. A mercury barometer
The column height is proportional to the
atmospheric pressure.
Atmospheric pressure results
from the mass of the
atmosphere and gravitational
forces.
The pressure is the force per
unit area.
P = F/A
1 atm = 760 mmHg
1 atm = 1.01325 E5 Pa
1 mmHg = 1 torr

29. Units for Expressing Pressure
Unit
Atmosphere
Pascal (Pa)
Kilopascal (kPa)
mmHg
Torr
Bar
mbar
psi
Value
1 atm
1 atm = 1.01325 x 105 Pa
1 atm = 101.325 kPa
1 atm = 760 mmHg
1 atm = 760 torr
1 atm = 1.01325 bar
1 atm = 1013.25 mbar
1 atm = 14.7 psi

30. Pressure
is equal to force/unit area
 SI units = Newton/meter2 = 1 Pascal (Pa)
 1 standard atmosphere = 101,325 Pa
(100,000 Pa = 1 bar)
 1 standard atmosphere = 1 atm =

760 mm Hg = 760 torr = 1013.25 hPa
= 14.695 psi
Meteorologists often report pressure in
millibar; 1 mbar =0.001bar =0.1 kPa =
1hPa

31. Variables Affecting Gases
 Pressure
(P)
 Volume
(V)
 Temperature
 Number
(T)
of Moles (n)

32. Elevation and Atmospheric
Pressure

33. Manometer
Manometers are
used to measure
gas pressure in
closed systems.
For instance in a
reaction vessel.

34. The Gas Laws and Absolute Temperature
The relationship between the volume, pressure, temperature, and mass of a gas is
called an equation of state.
We will deal here with gases that are not too dense.
Boyle’s Law: the volume of a given amount of gas is
inversely proportional to the pressure as long as the
temperature is constant.

35. Boyle Law…pressure is inversely
proportional to volume (at constant
T and moles, n).

36. The Gas Laws and Absolute Temperature
The volume is linearly proportional to the temperature, as long as the
temperature is somewhat above the condensation point and the
pressure is constant:
Extrapolating, the volume becomes zero at −273.15°C; this temperature
is called absolute zero.

37. Avogadro’s Law

For a gas at
constant
temperature and
pressure, the volume
is directly
proportional to the
number of moles of
gas (at low
pressures).
V α n
V1 = V2
n1
n2

38. The Gas Laws and Absolute Temperature
The concept of absolute zero allows us to define a third temperature scale –
the absolute, or Kelvin, scale.
This scale starts with 0 K at absolute zero, but otherwise is the same as the
Celsius scale.
Therefore, the freezing point of water is 273.15 K, and the boiling point is
373.15 K.
Finally, when the volume is constant, the pressure is directly proportional to
the temperature:

39. Combined Gas Law
 Combining
the gas laws the relationship
P α T(n/V) can be obtained.
 If
n (number of moles) is held constant,
then PV/T = constant.
P1
V
P2
V
1
2
=
T
T
1
2

40. Ideal Gas Law
PV = nRT
R = universal gas constant
= 0.08206 L atm K-1 mol-1
P = pressure in atm
V = volume in liters
n = moles
T = temperature in Kelvin

41. Standard Temperature
and Pressure (for gases)
“STP”




P = 1 atmosphere
T = 0°C
The molar volume of an ideal gas is 22.42 liters at STP
(put 1 mole, 1 atm, R, and 273 K in the ideal gas law
and calculate V)
Note STP is different for other phases, e.g. solutions or
enthalpies of formation.

42. The Ideal Gas Law
A mole (mol) is defined as the number of grams of a substance that is numerically
equal to the molecular mass of the substance:
1 mol H2 has a mass of 2 g
1 mol Ne has a mass of 20 g
1 mol CO2 has a mass of 44 g
The number of moles in a certain mass of material:

43. The Ideal Gas
Equation
• Charles’s Law:
1
V ∝ (constant n, T )
P
V ∝ T (constant n, P )
• Avogadro’s Law:
V ∝ n (constant P, T )
• Boyle’s Law:
• We can combine these into a general gas law:
nT
V∝
P

44. The Ideal Gas
Equation
• R = gas constant, then
 nT 
V = R

 P 
• The ideal gas equation is:
PV = nRT
• R = 0.08206 L·atm/mol·K = 8.3145 J/mol·K
• J = kPa·L = kPa·dm3 = Pa·m3
• Real Gases behave ideally at low P and high T.

45. Ideal Gas Law in Terms of Molecules:
Avogadro’s Number
Since the gas constant is universal, the number of molecules in one mole is the
same for all gases. That number is called Avogadro’s number:
The number of molecules in a gas is the number of moles times Avogadro’s
number:

46. Ideal Gas Law in Terms of Molecules:
Avogadro’s Number
Therefore we can write:
where k is called Boltzmann’s constant.
(13-4)

47. The Ideal Gas
Equation
 Calculate
the pressure exerted by 84.0 g of ammonia,
NH3, in a 5.00 L container at 200. oC using the ideal gas
law.
PV = nRT
P = nRT/V n = 84.0g * 1mol/17 g T = 200 + 273
P = (4.94mol)(0.08206 L atm mol-1 K-1)(473K)
(5 L)
P = 38.3 atm

48. Tutorial 2.
 The
pressure on a sample of an ideal gas was
increased from 715 torr to 3.55 atm at constant
temperature. If the initial volume of the gas was 485.
mL, what would be the final volume?
A
7.9 L sample of gas was cooled from 79°C to a
temperature at which the volume of the gas was
4.3 L. Assuming the pressure remains constant,
calculate the final temperature.
 Calculate the pressure in atmospheres and pascals
of a 1.2 mol sample of methane gas in a 3.3 L
container at 25°C.

49. Real Gases:
Deviations from Ideality



Real gases behave ideally at ordinary
temperatures and pressures.
At low temperatures and high pressures real
gases do not behave ideally.
The reasons for the deviations from ideality
are:
1.
The molecules are very close to one
another, thus their volume is important.
2.
The molecular interactions also become
important.
J. van der Waals, 1837-1923,
Professor of Physics,
Amsterdam. Nobel Prize 1910.

50. Real Gases:
Deviations from Ideality
 van
der Waals’ equation accounts for the
behavior of real gases at low temperatures
and high pressures.

n 2a 
V − nb) = nRT
P +
2 (
V 

•
The van der Waals constants a and b take into account two things:
1. a accounts for intermolecular attraction
a.
b.
For nonpolar gases the attractive forces are London Forces
For polar gases the attractive forces are dipole-dipole attractions
or hydrogen bonds.
2. b accounts for volume of gas molecules
At large volumes a and b are relatively small and van der Waal’s
equation reduces to ideal gas law at high temperatures and low
pressures.

51. Real Gases: Deviations from Ideal
Behavior
The van der Waals Equation
nRT
n 2a
P=
− 2
V − nb V
Corrects for
Corrects for
molecular
molecular
volume
attraction
• General form of the van der Waals equation:

n 2a 
P +
( V − nb ) = nRT

V2 



52. Example

53. Condensed matter : the three
states of matter.

54. Some Characteristics of Gases, Liquids and Solids and the Microscopic
Explanation for the Behavior
gas
liquid
solid
assumes the shape and
volume of its container
particles can move past
one another
assumes the shape of
the part of the container
which it occupies
particles can move/slide
past one another
retains a fixed volume
and shape
rigid - particles locked
into place
compressible
lots of free space
between particles
not easily compressible
little free space between
particles
not easily compressible
little free space between
particles
flows easily
particles can move past
one another
flows easily
particles can move/slide
past one another
does not flow easily
rigid - particles cannot
move/slide past one
another

55. Clearly, a theory used to describe
the condensed states of matter
must include an attraction between
the particles in the substance
 Condensed
of Matter:
.
 Liquids
 Solids
States

56. Kinetic Theory Description of the
Liquid State.
 Like
gases, the
condensed states of
matter can consist of
atoms, ions, or
molecules.
 What separates the
three states of matter is
the proximity of the
particles in the
substance.
 For the condensed
states of matter the
particles are close
enough to interact.

57. Phase Changes

58. Triple Point Diagram of Water
 Regions:
Each region corresponds
to one phase which is stable for
any combination of P and T within
its region
 Lines Between Region: Lines
separating the regions
representing phase-transition
curves
 Triple Point: The triple point
represents the P and T at which all
3 phases coexist in equilibrium
 Critical Point: At the critical point
the vapor pressure cannot be
condensed to liquid no matter
what pressure is applied.

59. Tutorial 3
 Van
der Waals, realized that two of the assumptions
mentioned above were questionable. He then
developed the Van der Waals equation of state
which predicted the formation of liquid phase. Write
the equation and state two corrections that he
made.
 (a) Calculate the pressure exerted by 1.00 mol of
CO2 in a 1.00 L vessel at 300 K, assuming that the
gas behaves ideally. (b) Repeat the calculation by
using the van der Waals equation.
 Sketch and label the liquid region, gas region solid
region and triple point in water phase diagram