2. As was the case for gaseous substances, the kinetic
molecular theory may be used to explain the behavior
of
solids and liquids. In the following description, the term
particle will be used to refer to an atom, molecule, or
ion. Note that we will use the popular phrase
“intermolecular attraction” to refer to attractive forces
between
the particles of a substance, regardless of whether these
particles are molecules, atoms, or ions.
3. Kinetic Molecular Theory
The KMT explains the properties of solids and liquids
in terms of intermolecular forces of attraction and the
kinetic energy of the individual particles.
4. Recalling the concept:
Solid Liquid Gas
Arrangement of
particles
Closely and
orderly packed
Less closely
packed
Very far apart
Kinetic energy Vibrate and rotate
about a fixed
position
Particle slide over
each other
Particles move
at great speed
Particle motion Very low Low High
Attractive forces Very strong Strong Low
5.
6. Plasma
• Like gas, plasma does not have a definite shape or
volume.
• Produce magnetic fields and electric current and
respond to strong electromagnetic forces.
7.
8. The 4th state of matter (Plasma)
• Formed by providing heat to the gas
• It is an ionized gas consisting of
positive ions and free electrons,
typically at low pressures
(fluorescent lamp) or at very high
temperatures (Sun)
• It is where all the states of matter
arises from.
9. Thermal vs. non-thermal plasma
Thermal Plasma
• Have electrons and
heavy particles at the
same temperature; they
are in thermal
equilibrium with each
other.
Non-thermal Plasma
• Have ions and neutrals
at a much lower
temperature, whereas
electrons are much
hotter.
10.
11. Types of plasma
Natural plasma
• Stars, solar wind,
interstellar nebulae
Artificial plasma
• TV screens, fluorescent
lamps, Laser-produced
plasmas, Plasma used in
semiconductor device
fabrication.
Natural plasma
• Lightning, Polar
aurora, Blue jets
12.
13.
14.
15.
16.
17. Compressibility of gas
Gaseous butane is
compressed within the
storage compartment of
a disposable lighter,
resulting in
its condensation to the
liquid state.
18. Intermolecular Forces of Attraction
Intermolecular Forces are
attractive forces between
molecules or particles in
the solid or liquid states
19. Intermolecular Forces of Attraction
Intermolecular Forces (IMF) are relatively
weaker than the forces within the molecules
forming bonds (intramolecular forces)
Intramolecular Forces holds atoms together
in a molecule
20. Intermolecular Forces of Attraction
The intermolecular forces of attraction in a
pure substance are collectively known as
van der Waals forces
1. Dipole-dipole
2. Hydrogen bonding
3. Ion-dipole
4. London dispersion
5. Dipole-induced dipole forces
21. van der Waals forces:
❑The term for all known
Intermolecular forces.
❑Named after a Dutch
scientist: Johannes van
der Waals (1837 –
1932)
22. Types of van der Waals forces:
Dipole – dipole
Exist between polar molecules. One end of a
dipole attracts the oppositely charged end of
the other dipole.
Weaker force compared to ion-dipole
(depending on size)
Example:
a. Dichloromethane
b. Hydrochloric Acid
23. Types of van der Waals forces:
Hydrogen Bonds
⦿ Plays an important role in life processes
⦿ It can easily be broken and reformed
⦿ Occurs in water, DNA molecules, and protein
⦿ It is an attractive interaction between a hydrogen atom
bonded to an electronegative Fluorine, Oxygen and
Nitrogen atom and an unshared electron pair of another
nearby electronegative atom.
24. Types of van der Waals forces:
Hydrogen Bonds
Example:
a. Water (H2O)
b. Ammonia
c. Ammonia and Water (NH3)
d. Hydrofluoric Acid (HF)
Ammonia
25. Types of van der Waals forces:
Ion – dipole
⦿ Results when an ion and the partial charge
found at the end of the polar molecule attract
each other.
⦿ Positive ions are attracted to the negative
end of a dipole and vice versa. This
explains the solubility of ionic
compounds in water which is polar
molecule.
Example:
a. Salt (NaCl) Dissolved in Water (H2O)
b. Potassim (K+) Dissolved in Hydrochloric
Acid (HCl)
26. Types of van der Waals forces:
Ion – dipole
Example:
a. Salt (NaCl) Dissolved
in Water (H2O)
b. Potassim (K+)
Dissolved in
Hydrochloric
Acid (HCl)
27. Types of van der Waals forces:
London Dispersion Forces
⦿ It is the weakest type of intermolecular forces
⦿ When two non-polar molecules approach each other,
an instantaneous dipole moment forms.
⦿ Cl2 and CH4
28. London Dispersion Forces
Even distribution
of e-
Uneven
distribution of e-
due to constant
movement making
one side partially +
(Polarization)
Temporary
Partial + attracts
nonpolar and form
forces of attraction
Atom will be
distorted
29. Types of van der Waals forces:
London Dispersion Forces
⦿ Dipole can be induced more likely on
molecules having larger molecular masses.
(Polarizability)
-This also affects the melting and boiling
points of the molecules.
33. POPERTIES OF LIQUIDS
A. VISCOSITY
⦿ What is the difference between fluid
and viscous liquids?
⦿ VISCOSITY is the ability of a fluid to
resist flowing.
34. POPERTIES OF LIQUIDS
A. VISCOSITY
⦿ Viscosity of a
liquid depends on
intermolecular
forces that is
present.
⦿ The stronger the
IMF, the higher
the viscosity
35. POPERTIES OF LIQUIDS
A. VISCOSITY
⦿ Non-polar molecules have low
viscosities because of weak London
Force. Example: Benzene, pentane and
carbon tetrachloride.
⦿ Polar molecules
such as glycerol
and aqueous
sugar solution
have high
viscosities.
36. POPERTIES OF LIQUIDS
A. VISCOSITY
⦿ Long-chained substances
like oil have greater IMF
because there are more
atoms that can attract one
another, contributing to the
substance’s total attractive
forces.
37. POPERTIES OF LIQUIDS
A. VISCOSITY
What do you think is the effect of an
increasing temperature to the viscosity of
a liquid?
The viscosity
decreases as
the temperature
increases
40. POPERTIES OF LIQUIDS
B. SURFACE TENSION
⦿ The measure of the resistance of a liquid to
spread out.
⦿ The higher IMF means higher surface tension
⦿ The higher the temperature,
the less the strength of the
attractive force that holds
the molecule together
41. POPERTIES OF LIQUIDS
B. SURFACE TENSION
⦿ Allows needles and paper clips to
float in water if placed carefully on
the surface. It also explains why
drop of water are spherical in
shaped
42. POPERTIES OF LIQUIDS
B. SURFACE TENSION
⦿ Molecules within a liquid are
pulled in all directions by IMF
⦿ Molecules at the surface are
pulled downward and
sideways by other molecules,
not upward away from the
surface
43.
44.
45. POPERTIES OF LIQUIDS
C. CAPILLARITY
⦿ The tendency of a liquid to
rise in narrow tubes or to be
drawn into small openings
such as in plants transport
system.
⦿ Results from competition
between liquid’s
intermolecular force and the
walls of the tube.
48. POPERTIES OF LIQUIDS
C. CAPILLARITY
Two types of forces are involved in capillary action
⦿ Cohesion – intermolecular attraction between like liquid
molecules
⦿ Adhesion – attraction between unlike molecules(water
and the one that make up the glass tube)
⦿ These forces also define the shape of the surface of a
liquid in a cylindrical container (meniscus)
49. POPERTIES OF LIQUIDS
C. CAPILLARITY
⦿ When the cohesive force between the liquid molecules
is greater than the adhesive forces between the liquid
and the walls of the container, the surface of the liquid
is convex.
⦿ When the cohesive forces between the liquid molecules
are lesser than the adhesive forces between the liquid
and the walls of the container, the surface of the liquid
is concave.
50.
51. POPERTIES OF LIQUIDS
D. VAPOR PRESSURE
⦿ It is the pressure exerted by its
vapor when in equilibrium with
liquid or soilid.
⦿ Example: when liquid or solid
substances is made to evaporate in
a closed container, the gas exerts a
pressure above the liquid.
52. POPERTIES OF LIQUIDS
D. VAPOR PRESSURE
⦿ Substances with relatively strong IMF
will have low vapor pressure because
the particles will have difficulty
escaping as a gas.
Example:
⦿ Water(H bonding) has vapor pressure of
0.03 atm
⦿ Ethyl Ether Dipole-dipole and London
forces has vapor pressure of 0.68 atm
53.
54. POPERTIES OF LIQUIDS
E. BOILING POINT
⦿ The boiling point of a liquid is the
temperature at which its vapor pressure
is equal to the external or atmospheric
pressure
⦿ Increasing the temperature of a liquid
raises the KE of its molecules, until
such a point where the energy of the
particle movement exceeds the IMF that
holds them together.
55. POPERTIES OF LIQUIDS
E. BOILING POINT
⦿ The liquid molecules then transform to
gas and are seen as bubbles that rises
to the surface of the liquids and escape
to the atmosphere.
⦿ Temperature at which a liquid boils
under 1 atmospheric pressure is
referred to as its normal boiling point.
56. POPERTIES OF LIQUIDS
E. BOILING POINT
⦿ At higher altitudes, the atm is lower,
hence, the boiling point will
subsequently decrease.
⦿ The greater IMF, the higher the energy
needed to increase the KE of the
molecules to break these forces.
57.
58. POPERTIES OF LIQUIDS
F. HEAT OF VAPORIZATION
⦿ Molar Heat of vaporization is the
amount of heat required to vaporize one
mole of substance at its boiling point.
⦿ The application of heat disrupts the IMF
of attraction of the liquid molecules and
allows them to vaporize
59. POPERTIES OF LIQUIDS
F. HEAT OF VAPORIZATION
⦿ Boiling point generally increases as
molar heat of vaporization increases.
⦿ The heat of vaporization is also
determined by the strength of IMF
between molecules.
60.
61. POPERTIES OF LIQUIDS
ANALYSIS
⦿ Normal boiling point happens when a
liquid reaches an internal temperature
of 100OC under
1 atm (atmospheric pressure)
Which level with respect to sea
level foods cooks faster and
slower?
62. POPERTIES OF LIQUIDS
At higher altitude,
atmospheric pressure is
lesser.
Thus water boils faster at a
lower temperature because
less pressure is exerted on
water molecules.
Inefficient delivery of heat to
cook the food and it takes time
for the food to be cooked.
63. PROPERTIES OF SOLIDS
How are the structures and
properties of
solids related?
https://www.slideshare.net/marvinnbustamante1/general-
chemistry-2-chapter-1-the-kinetic-molecular-model-and-
intermolecular-forces-of-attraction-in-matter
64. POPERTIES OF SOLIDS
⦿ A SOLID is formed when
the temperature of a
liquid is low and the
pressure is sufficiently
high causing the
particles to come very
close to one another.
⦿ They are rigid
⦿ Their particles hardly
diffuse
66. POPERTIES OF SOLIDS
A. CRYSTALLINE
⦿ Atoms, ions, or molecules are arranged
in well defined arrangement
⦿ Having flat surface and sharp edges
⦿ Example: gems, salts, sugar and ice.
68. POPERTIES OF SOLIDS
1. Ionic Crystalline Solids
⦿ Composed of (+) and (-) ions
⦿ Held by electrostatic attractions
⦿ They are hard, brittle and poor
electrical and thermal conduction
⦿ Example: NaCl
69. POPERTIES OF SOLIDS
2. Molecular Crystalline Solids
⦿ Composed of atoms and molecules
⦿ Held together by: H-Bond, dipole-
dipole, and London dispersion forces
⦿ Soft, low to moderate melting point and
poor thermal and electrical
conductivity
⦿ Examples: CH4, C12H22O11, CO2, H2O
and Br2
70. POPERTIES OF SOLIDS
3. Covalent Network Crystalline
Solids
⦿ Atoms connected in a network of
covalent molecules
⦿ Held together by covalent bonds
⦿ Very hard, very high melting point and
often poor thermal and electrical
conductivity.
⦿ Examples: Plastics, Allotropes of
carbon, silicon carbide
71. POPERTIES OF SOLIDS
4. Metallic Crystalline Solids
⦿ Composed of atoms and molecules
⦿ Held together by metallic bonds
⦿ Soft to hard, low to high melting point,
malleable, ductile and good thermal
and electrical conduction
⦿ All metallic elements: Cu, Na, Zn, Fe
and Al
72. POPERTIES OF SOLIDS
How are molecules being
arranged in microscopic level?
Unit Cell
Crystal
Lattice
The smallest portion of the
crystal which shows the
complete pattern of its
particles
The repetition of unit cells
in all directions
74. POPERTIES OF SOLIDS
B. AMORPHOUS SOLIDS
⦿ From the Greek word for “without
form”
⦿ Solid particles which do not have
orderly structures.
⦿ They have poorly defined shapes
⦿ are rigid, but they lack repeated
periodicity or long-range order in their
structure.
⦿ examples include thin film lubricants,
metallic glasses, polymers, and gels
77. POPERTIES OF SOLIDS
⦿ It can be noted that as temperature of
crystalline solid is increased, the
particles vibrate back and forth about
its lattice point.
⦿ The crystal becomes less ordered.
⦿ The heat added increases the kinetic
motion of the particles.
⦿ Until the crystalline structure is
completely destroyed by the vibrations
of the particles, melting is achieved.
78. POPERTIES OF LIQUIDS
ANALYSIS
⦿ What will happen if heating stops and
no heat is allowed to escape?
Both solid and liquid phases are
present in equilibrium.
