Kinetic molecular theory explains the states of matter based on the idea that matter is composed of tiny particles in constant motion. It describes the differences between solids, liquids, and gases in terms of intermolecular forces and molecular motion. In liquids, molecules are close together and can move past one another, allowing liquids to flow. Solids have molecules held rigidly in position with little motion. The properties of liquids, such as viscosity, surface tension, vapor pressure, and boiling point, depend on the balance between kinetic energy and intermolecular forces between molecules.

1. KINETIC MOLECULAR MODEL
OF SOLIDS AND LIQUIDS
What is Kinetic Molecular Theory?
- a theory that explains the states is matter
and is based on the idea that matter is
composed of tiny particles that are always in
motion.
- explain observable properties and
behaviors of solids, liquids, and gases.
- An application of the theory is that it helps
to explain why matter exists in different
phases (solid, liquid, and gas) and how
matter can change from one phase to
another.
➢ The state of a substance depends on the
balance between the kinetic energy of
individual particles (molecules or atoms)
and the intermolecular forces.
➢ The kinetic energy keeps the molecules
apart and moving around and is a function
of the temperature of the substance.
➢ The intermolecular forces are attractive
forces that try to draw the particles together.
Postulates of Kinetic Molecular
Theory
1. Matter is made of particles that are
constantly in motion. This energy of motion
is called kinetic energy.
2. The amount of kinetic energy in a
substance is related to its temperature. An
increase in temperature means greater
speed.
3. There is space between particles. The
amount of space in between particles is
related to the substance’s state of matter.
4. Phase changes happen when the
temperature of the substance changes
sufficiently.
5. intermolecular forces are attractive
forces in between particles. The strength of
these forces increases as particles get
closer together.
KMT of Liquids and Solids
➢The principal difference between the
condensed states (liquids and solids) and
the gaseous state is the distance between
molecules.
LIQUID
❖In a liquid, the molecules are so close
together that there is little space. Thus,
liquids are much more difficult to compress
than gases, and they are also much denser
under normal conditions.
❖Molecules in a liquid are held together by
one or more types of attractive forces.
❖A liquid also has a definite volume,
because molecules in a liquid do not break
away from the attractive forces.
❖The molecules can, however, move past
one another freely, so a liquid can flow, be
poured, and assume the shape of its
container.
SOLID
❖In a solid, molecules are held rigidly in
position with virtually no freedom of motion,
so they only vibrate about fixed positions.
❖There is even less space in a solid than in
a liquid because their particles are tightly
packed.
❖Thus, solids are almost incompressible
and possess definite shape and volume.
This is due to a more potent intermolecular
force of attraction compared to liquids.

2. INTERMOLECULAR FORCES
Intramolecular Forces
❖Intramolecular (within molecules) forces
hold atoms together in a molecule.
❖Intramolecular forces stabilize individual
molecules.
❖Generally, these forces are simply the
chemical bonds such as ionic and covalent
bonding.
Intermolecular Forces
❖Intermolecular forces are attractive forces
between molecules.
❖Intermolecular forces are responsible for
the non-ideal behavior of gases, but they
exert more influence in the condensed
phases of matter which are liquids and
solids.
❖Intermolecular forces are collectively
known as the Van der Waals force named
after Dutch chemist Johannes van der
Waal.
❖Van der Waals forces are electrical; that
is, they result in the attraction between
centers of opposite charge in two molecules
close to each other.
INTRA molecular forces - the forces
holding atoms together to form molecules.
INTER molecular forces - Forces between
molecules, between ions, or between
molecules and ions.
Types of Intermolecular Forces
1. London Dispersion Forces
(dispersion forces)
➢Intermolecular forces of attraction that
exist between all atoms and molecules
➢The only kind of intermolecular forces
present among symmetrical nonpolar
substances such as O2 and CO2 and
monoatomic species such as noble gases
Figure 1. The boiling point of noble gases
increases as one goes from top to bottom of
the periodic table. Dispersion forces
increase with increasing atomic size.

3. ❖From the figure, it can be observed that
the boiling point increases as you go from
top to bottom of the periodic table.
❖This means that the greater the atomic
size, the greater the boiling point.
❖It can also be observed that the greater
the molar mass, the higher the boiling point.
❖There are times when an atom loses its
symmetry, resulting in a temporary
separation of charges or temporary dipole.
❖Temporary dipole results from a shift in
the position of the electrons where one end
becomes more negative causing the other
end to be more positive.
❖At this point, the centers of positive and
negative charges do not coincide.
Figure 2. London dispersion forces are
temporary attraction forces that result when
the electrons in two adjacent atoms occupy
positions that make the atoms form a
temporary dipole.
❖It is a dipole created by the presence of a
neighboring dipole.
❖As the molar mass or atomic size
increases, dispersion forces increase
because more dipoles can be induced in
larger substances.
❖The greater the number of dipoles, the
greater the dispersion forces.
2. Dipole-Dipole Forces
➢Nonpolar molecules result when the
electronegativity difference between two
atoms is less than or equal to 0.4
Examples: Cl2, H2, CCl4
➢Polar molecule is formed when there is
an uneven sharing of electrons between
atoms
Examples: HCl, H2O
Difference Bond Symmtery
Less than 0.5 nonpolar symmetrical
0.5-1.9 polar unsymmetrical
Dipole-Dipole Forces
➢Dipole-dipole forces are formed between
neighboring molecules with permanent
dipoles.
➢The dipole-dipole forces are strong
because of the attraction of opposite
charges that are permanent within the entire
substance.
➢The strength for dipole-dipole forces
increases as the magnitude of the dipole

4. increases and the distance between the
molecules decreases.
➢ The solubility of a solute in a solvent can
be estimated by considering the energy
required to break bonds and the energy
released when bonds form.
➢Solubility of polar substances in polar
liquids can be explained by considering the
energy required to break the solute-solute
"bonds" and the solvent-solvent "bonds" in
comparison to the energy released when
the solvent-solute "bonds" form.
➢If the latter is too small when compared to
the former, the substance is not soluble.
➢Since this energy balance is rarely
achieved between substances that are not
similar, an often-quoted axiom is
" like dissolves like". " Like dissolves like”
is a statement of fact NOT, it is an
explanation of the phenomenon.
➢The relative magnitude of these forces
can also be used to explain trends in
melting points and boiling points.
➢It must be remembered that both melting
point and boiling point tend to increase with
increasing molar mass, all other factors
being equal.
3. Hydrogen Bond
❖This is a result of a high partial positive
charge on hydrogen and a large partial
negative charge for the more
electronegative atom (F, O, N).
❖Because of the very large dipole
produced between the hydrogen atom of
one molecule and the F, N, and O of
another molecule, a special name is given
to this kind of force-hydrogen bond.
❖Hydrogen bonding is a special case of
dipole-dipole forces and only exists between
hydrogen atoms bonded to F, N, or O, and
F, N, and O atoms bonded to hydrogen
atoms.
❖A special form of dipole attraction
enhances dipole dipole attractions.
❖H-bonding is strongest when X and Y are
N, O, or F
❖H-bonding is especially strong in
biological systems — such as DNA.
❖DNA — helical chains of phosphate
groups and sugar molecules. Chains are
helical because of the tetrahedral geometry
of P, C, and O.
❖Chains bind to one another by specific
hydrogen bonding between pairs of Lewis
bases.
- Adenine with Thymine, Guanine with
Cytosine
AMP = Adenosine monophosphate
4. Ion-Dipole Forces
➢It results from the interaction between an
ion and a polar molecule.
➢The ion-dipole attraction becomes
stronger when the charge on the ion
increases and when the magnitude of the
dipole becomes stronger.
Example: The ion-dipole forces in CaCl2 are
stronger than the ion-dipole forces in KCl
because the charge of the ion in calcium is
higher than in potassium

5. ➢Ion - Ion-dipole forces exist between ions
and polar molecules.
➢The magnitude of these forces increases
as: –the distance between the ion and
the polar molecule decreases
–the magnitude of the charge on the ion
increases
–the magnitude of the dipole of the polar
molecule increases.
➢Hydration energies for cations and anions
are an excellent example of this concept.
➢When these hydration bonds form, energy
is released, exothermic.
➢This energy is then used to break the
ion-ion forces in the ionic solid.
➢When the hydration energy is large
enough, the ionic solid is soluble in water.
➢Solubility trends for ionic solids can be
explained by using this combination of
forces.
➢Explain the trend in hydration energies for
Fe+2, Ca+2, and Fe+3. The calcium ion has
the largest radius and the Fe+3 is the
smallest radius.
PROPERTIES OF LIQUIDS
LIQUIDS
• Molecules are in constant motion
• There are appreciable intermolecular
forces
• Molecules close together
• Liquids are almost incompressible
Properties of Liquids
1. Surface Tension and
Capillary Action
• Surface tension is the result of the
intermolecular force acting at the surface of
a liquid.
• Capillary action, i.e. rising of a fluid in a
very small diameter tube, results from the
combination of adhesive forces, between a
solid (like glass) and the liquid and the
cohesive forces, between the molecules of
the liquid.
• If the cohesive forces are stronger, the
liquid forms an upward-rounded meniscus.
• A downward-rounded meniscus forms if
the adhesive forces are stronger.
Surface tension
- is the force that causes the surface of a
liquid to contract.
- Phenomena such as insects walking on
the surface of water, droplets of liquid being
spherical in shape, and a needle remaining
suspended on the surface of water can all
be explained in terms of surface tension.

6. ➢Molecules at the surface behave
differently than those in the interior.
Molecules at the surface experience the net
INWARD force of attraction.
➢This leads to the SURFACE TENSION -
the energy needed to break the surface.
also leads to spherical liquid droplets.
➢The stronger the intermolecular force of
attraction, the greater the surface tension.
➢An increase in temperature decreases
surface tension.
➢Water has a high surface tension
because of its ability to form a hydrogen
bond.
➢IM forces also lead to CAPILLARY action
and the existence of a concave meniscus
for a water column.
2. Viscosity
❖Viscosity is the resistance of fluids to flow.
❖Viscosity is a measure of the substance’s
intermolecular force of attraction (IMFA).
❖The greater the IMFA, the higher the
viscosity, and the less readily the liquid
flows.
❖The greater the resistance in flowing, the
more viscous the liquid is.
❖Temperature also affects viscosity. The
higher the temperature, the lower the
liquid’s viscosity.
❖An increase in temperature causes the
kinetic energy to increase.
❖Heat breaks the intermolecular forces
causing the liquid molecules to move faster.
This makes the molecules flow more readily.
3. Vapor Pressure
❖Vaporization is a change of state from
liquid to gas.
❖When liquid molecules break free from
their neighbors and escape into the gas
phase, the process is called evaporation.
❖Vaporization is a broader term that
includes evaporation and boiling.
❖Vapor is used to refer to the gaseous
state of a substance which is normally a
liquid or solid at room temperature.
❖Substances that evaporate readily are
volatile. They have weak intermolecular
forces.
❖Examples of volatile liquids are alcohol,
gasoline, paint thinner, and dry-cleaning
solvents.
❖Volatile substances burn more readily
since they easily combine with oxygen.
❖At higher temperatures, more of the
molecules have sufficient energy to escape.
❖Since vaporization is an endothermic
process, condensation is an exothermic
process.

7. • The magnitude of ΔHvap is related to the
type and magnitude of the intermolecular
forces found in the liquid.
➢The two key properties we need to
describe are EVAPORATION and its
opposite, CONDENSATION
➢To evaporate, molecules must have
sufficient energy to break IM forces.
Breaking IM forces requires energy.
➢The process of evaporation is
Endothermic.
➢When molecules of liquid are in the vapor
state, they exert a VAPOR PRESSURE
➢EQUILIBRIUM VAPOR PRESSURE is
the pressure exerted by a vapor over a
liquid in a closed container when the rate of
evaporation = the rate of condensation.
➢The vapor pressure is the equilibrium
pressure of the vapor above the liquid (or
solid) at a given temperature or the
pressure of the vapor resulting from the
evaporation of a liquid above a sample of
the liquid in a closed container.
➢The equilibrium vapor pressure is
dependent on temperature.
➢If the temperature is increased,
evaporation will take place more readily and
will further increase the number of vapor
molecules.
➢If the number of vapor molecules
increases, vapor pressure also increases.
➢Temperature is directly proportional to
vapor pressure; an increase in temperature
increases vapor pressure.
➢If a liquid has a weak intermolecular force
of attraction, the escaping tendency of the
molecules is high. This causes a high vapor
pressure for a liquid.
➢If the escaping tendency of the molecules
is low, evaporation is slow, and this
produces a low vapor pressure of the liquid.
➢Compounds with higher vapor pressures
are more volatile than those with lower
vapor pressures.
➢The stronger the intermolecular forces,
the lower the vapor pressure.
➢When VP = external P, the liquid boils.
* This means that the BP’s of liquids change
with
altitude.
➢As the temperature increases, the vapor
pressure increases since there are more
higher energy molecules at the higher
temperature.
4. Boiling Point
❑The boiling point, Tb, is the temperature
when the equilibrium vapor pressure equals
the external pressure.
❑The normal boiling point, Tbo, is the
temperature when the equilibrium vapor
pressure equals one-atmosphere pressure
or 760 torr.
❑The boiling point is the temperature at
which the vapor pressure of the liquid is
equal to the atmospheric pressure.
❑Liquids that have high vapor pressure
have low boiling points.

8. ❑The higher the vapor pressure of a liquid,
the lower its boiling point.
❑The lower the atmospheric pressure, the
faster it is to equalize the vapor pressure of
the liquid and the atmospheric pressure,
and the lower the boiling point is.
❑A liquid that evaporates readily has weak
intermolecular forces and a high vapor
pressure. The normal boiling point is low.
❑If the liquid has strong intermolecular
forces, the escaping tendency is low, its
vapor pressure is low, and it has a high
boiling point.
➢A liquid boils when its vapor pressure
equals atmospheric pressure.
➢When pressure is lowered, the vapor
pressure can equal the external pressure at
a lower temperature.
Critical Temperature and Pressure
• The critical temperature, Tc, is the
temperature at which the liquid state no
longer exists since all molecules have
sufficient energy to be separated from each
other.
• The critical pressure, Pc, is the pressure
corresponding to the critical temperature,
where no further increase in pressure will
cause the gas phase to condense into the
liquid phase.
• This (Tc, Pc) point is called the critical
point on the vapor pressure graph.
5. Molar Heat of Vaporization
❖Heat of vaporization is the amount of heat
needed to vaporize a given amount of
substance at its boiling point.
❖Substances with high heat of vaporization
have strong intermolecular forces of
attraction (IMFA).
❖The stronger the IMFA, the higher the
heat of vaporization.
➢For water, the amount of heat needed to
vaporize 1 mole of water (18 grams) is 40.7
kilojoules or 540 calories/gram at the same
temperature and one-atmosphere pressure.
❖Even if heat is added to the liquid to
convert it to gas, the kinetic energy of the
molecules remains constant.

9. STRUCTURE AND
PROPERTIES OF WATER
Structure of Water and Properties of Water
❖Water’s unique properties from the strong
intermolecular forces of attraction
characterized by the hydrogen bond.
❖H-bonding is especially strong in water
because
▪the O-H bond is very polar
▪there are 2 lone pairs on the O atom
▪Accounts for many of water’s unique
properties.
Properties of Water
1. High boiling point
➢The high boiling point of water is a
consequence of its strong intermolecular
forces of attraction caused by the formation
of the H-bond.
➢It also explains why water is liquid at
room temperature.
▪H bonds is an abnormally high boiling
point of water.
2. High specific heat
- refers to the amount of heat needed to
change the temperature of 1 gram of a
substance by 1°C.
➢Water can absorb and release large
quantities of heat without a change in
temperature.
➢This is the reason why body temperature
remains at 37°C even when there’s a
change in the surroundings.
➢Water has high specific heat which
requires large amounts of heat before it
vaporizes.
➢This explains why the Earth has minimal
temperature variations that can affect the
climate.
3. High density in its liquid form
➢Water is the only substance that contracts
when cooled.
➢Solid form for most substances is denser
than their liquid form.
➢Ice has an open structure because the
hydrogen bonds could not get inside the
hexagonal ring structure.
➢H-bonding in H2O ----> open lattice-like
structure of ice.
➢Ice density is less than that of liquid, and
solid
floats on water.
➢ This more open structure of the solid
form of water causes the ice to have a
smaller number of molecules packed in a
given volume.
➢ This causes the mass to be lower, hence,
the density of ice is lesser than the liquid
water and as a result, ice floats on water.
4. High surface tension
➢The hydrogen bond formation among
water molecules causes water to have high
surface tension.
➢This high surface tension causes water to
move from the roots of a tree to the top of
very tall trees and explains why water
moves into the fibers of a towel. This
phenomenon is called capillary.
5. High heat of vaporization
➢Large amount of heat is needed to
vaporize a given amount of water.
➢This causes a significant drop in
temperature during evaporation.
➢This explains why perspiration lowers the
body temperature.

10. TYPES AND PROPERTIES OF
SOLIDS
➢Solids have very strong forces of
attraction and therefore, their particles are
not free to move from one another.
➢Particles of solids are packed closely
together and are arranged in an organized
pattern.
➢Solids have definite shape and volume
and have high densities.
Crystalline Solids
❖Solids whose particles are arranged in
regular geometric patterns are called
crystalline solids.
❖The definite patterns that repeat
themselves in solid crystals are called unit
cells.
Examples: salt, sugar, snow, and many
precious stones used in jewelry
Different unit cells (representative unit)
of the seven principal crystal patterns in
crystalline solids
Amorphous Solids
❖Solids that have fixed shape and volume,
but their particles are not arranged in a
regular geometric pattern are called
amorphous solids and referred to as
“supercooled liquids” because these solids
appear to have been cooled at very low
temperatures and their viscosities are very
high, preventing the flow of the liquid.
Examples: glass, rubber, and some plastics
Types of Crystals
1. Ionic Solids
➢The particles of ionic solids are positive
and negative ions.
➢The crystal arrangement in ionic solids
maximizes attraction and simultaneously
minimizes repulsion making the compound
possess a high degree of stability.
➢These solids are hard, they are brittle,
have high melting points, and have poor
electrical and thermal conductivity.
2. Covalent solids
➢This type of solids is made up of atoms
and is joined by covalent bonds.
➢Some solids form covalent bonds
resulting in the formation of molecules.
➢In some solids, however, molecules are
not formed. Rather, a covalent network is
formed extending throughout the solid
crystal.

11. ➢They are very hard, have very high
melting points, and often have poor thermal
and electrical conductivity.
3. Molecular solids
➢The particles in molecular solids can
either be atoms or molecules held together
by intermolecular forces like dispersion
forces, dipole-dipole forces, and hydrogen
bonds.
➢Solids with dispersion forces that are also
large are solids at room temperature. This is
also true for solids that are highly polar.
➢This type of solids is soft, has low to
moderately high melting points, and has
poor electrical and thermal conductivity.
➢They are poor conductors of heat and
electricity because there are no free-moving
electrons and no charged particles.
4. Metallic solids
➢These are joined by metallic bonds.
➢Metallic bonds are characterized by the
presence
of mobile electrons around the positive
metal ion.
➢The force that binds the atoms together is
the force of attraction between the mobile
valence electrons and the fixed positive
metal ion.
➢The strength of the force depends on the
metal and depending on the nature of the
metal, they can be soft to hard, and melting
points range from low to high.
➢These solids are good electrical
conductors (because of the mobile
electrons), good thermal conductors, and
are malleable and ductile.