CHEM 1411 CHAPTER 11
LIQUIDS AND SOLIDS
Intermolecular forces
The attractive forces between the molecules in compound are the intermolecular forces.
Intra molecular forces
These are the attractive forces within a molecule that hold the atoms together.
Both these forces constitute the secondary forces.
Types of Secondary forces
1. Dipole – Dipole Forces
These are the attractive forces between the dipoles in a polar molecule.
Ex. Water, CHCl3
2. Ion – Dipole Forces
Attractive forces between an ion (cation or anion) and oppositely charged end of a polar molecule are
the ion – dipole forces
These forces depend on the charge and size of the ion and on the magnitude of the dipole moment and
size of the polar molecule. Ex. Hydration
3. Dispersion Forces or London Forces
Weak attractive forces that arise as a result of temporary dipoles induced in atoms or molecules are
dispersion forces. Non-polar molecules exhibit this kind of forces. The dispersion forces appear in
molecules by the following ways.
 Ion – Induced dipole attraction: - Unequal distribution of electrons takes place in some nonpolar molecules under the influence of an ion resulting in the formation of temporary dipoles.
 Dipole – Induced dipole attraction: - A polar molecule bearing a dipole can induce temporary
dipoles in a non – polar molecule.
 By the erratic motion of molecules of the same substance resulting in the unequal distribution
of electrons resulting in the formation of temporary dipoles. This dipole can polarize other
molecules. For example He and N2 condenses at low temperatures due to dispersion forces.
Dispersion Forces usually increase with molecular mass because compounds with higher molecular mass
contain more number of electrons. For example, CCl4, which is non-polar, boils at 76.50C but polar
CH3F boils at a lower temperature of –78.40C.
4. The Hydrogen Bond
The unique and strong attractive force between molecules where a hydrogen atom is covalently
bonded to an electronegative atom like Oxygen, Nitrogen or Fluorine is the hydrogen bond.
Effects of Hydrogen Bonding
 Increases the boiling point of liquids and melting point of solids.
 Increases solubility of substances in polar solvents (water)
 Decides the physical state of some substances
Properties of Liquids
Surface Tension
This is the energy required to increase the surface area of a liquid by one unit (1 cm2). The effects of
surface tension are,
 Capillary action of liquids
 Drops of liquid assume spherical shape
Viscosity
Viscosity is a measure of fluid’s resistance to flow. The greater the viscosity, the more slowly the
liquid flows. Viscosity decreases with increase in temperature. Liquids having strong intermolecular
forces have higher viscosities.
Cohesion: - Intermolecular attractive forces between like molecules (attractive forces between water
molecules)
Adhesion: - Intermolecular attractive forces between unlike molecules (between water and glass)
The level of water in glass tube is concave because adhesion > cohesion
On the other hand level of mercury in a glass tube is convex because cohesion > adhesion
Structure and Properties of Water
In water the molecules are held together by strong intermolecular hydrogen bonding. Therefore,
water has a high boiling point and specific heat.
Ice floats on water because it is less dense than liquid water. In the structure of ice, 4 water molecules
surround each water molecule leaving a lot of empty space inside. This arrangement reduces the
density. When ice melts, water molecules get closely packed due to the formation of more H bonds
resulting in an increase in density. This continues till the temperature reaches 40C. Thus water shows
a maximum density at 40C and afterwards density decreases as usual.
Crystalline solids
In crystalline solids the constituents have a perfectly ordered 3 - dimensional arrangement. They have
a sharp melting point. They are characterized by a crystal lattice.
Amorphous Solids
In amorphous solids the constituents do not have a perfectly ordered arrangement. They do not have
a sharp melting point.
Ex. Glass, Rubber etc.
Co ordination Number
The number of constituents surrounding an atom / ion in the crystal lattice is the co ordination
number.
The basic repeating unit in a crystal lattice is called the unit cell. The unit cells are simple cubic; body
centered cubic and face centered cubic. In a simple cubic unit cell, the constituents are present at the
corners of the cube, a body centered cube contains constituents at the corners and at the center of the
body and in face centered cube, the constituents are present at the corners and at the center of each of
the six faces.
nxM
Density of the unit cell, D =
N0 x l 3
(Where ‘n’ = number of particles in the unit cell, ‘M’ = Molar mass, ‘N0’ = Avogadro’s number and
‘l’ = edge length of the cube)
Types of Crystals
1. Ionic crystals
In ionic crystals the constituents are ions. The forces of attraction between the constituents are
strong electrostatic forces. The anions are usually bigger than the cations. Therefore, the anions
form the close packed structure and the cations occupy the interstitial voids.
Ex. NaCl, CsCl, ZnS, CaCl2
2. Covalent Crystals
In covalent crystals the constituents are atoms, which are held in the 3 dimensional networks by
covalent bonds.
Ex. Graphite, Diamond, Quartz
In the structure of graphite, each C atom is bonded to 3 other C atoms to form hexagonal rings
resulting in the formation of a sheet like structure. 3 electrons from each C atom are used for this
bonding and the fourth electron is free. These free electrons make graphite a good conductor of
electricity.
In diamond each C atom is bonded to 4 other C atoms in a tetrahedral manner. No electron is free
in this arrangement and this makes diamond a bad conductor of electricity. The hardness and
high melting point of diamond is due to the presence of strong covalent bonds.
3. Molecular Crystals
The constituent particles are molecules held by weak van der Waals forces.
Ex. Solid SO2, Iodine, Ice, P4
4. Metallic Crystals
The constituents are metallic atoms held by strong metallic forces. Ex. Metals.
Phase changes
 Liquid – Vapor Equilibrium
Evaporation or vaporization is the process in which a liquid changes to its gaseous state. It is
an endothermic process. The rate of evaporation increases with rise in temperature because
the kinetic energy of molecules increases with temperature. The reverse of evaporation is
condensation, which is an exothermic process.
Evaporation
Liquid
Vapor
Condensation
Molar heat of vaporization is the amount of heat required to change one mole of a liquid into its vapor at
its boiling point.
Lowering the temperature and applying high pressure can liquefy a gas.
Equilibrium Vapor Pressure
The pressure exerted by the vapors of a liquid in a closed system when there is equilibrium between
vaporization and condensation is called the Equilibrium Vapor Pressure.
Larger the attractive forces between the molecules, the smaller the vapor pressure.
Boiling point of Liquids.
At this temperature the vapor pressure of a liquid becomes equal to the atmospheric pressure.
The boiling point of water at Mt. Everest is 700C, at normal level is 1000C and inside a
pressure cooker is above 1000C.
Liquids with stronger secondary forces have lower vapor pressure and higher boiling point.
Example: Water has higher boiling point than Carbon tetra chloride.
Critical temperature
This is the temperature at and above which, a gas cannot be liquefied no matter how great the applied
pressure is.
Greater the intermolecular forces, higher the critical temperature.
Critical Pressure
The minimum pressure that must be applied to a gas to bring about liquefaction at its critical
temperature is the critical pressure.

Liquid – Solid Equilibrium
The change of a liquid into its solid form is called freezing and the reverse process is melting
(Fusion).
Freezing
Liquid
Solid
Melting
Melting point or Freezing Point
The temperature at which the solid and liquid phases of a substance exist in equilibrium is the
Melting point or Freezing point of the substance.
Molar heat of fusion is the amount of heat required to change one mole of a solid into its liquid.
 Solid – Vapor Equilibrium
The change of a solid directly in to its vapor state is Sublimation. The reverse process is
Deposition.
Sublimation
Solid
Vapor
Deposition
Molar heat of sublimation is the amount of heat required to change one mole of a solid into its vapor.
H Sublimation =
H Fusion +
H Vaporization
Phase Diagram
A phase diagram summarizes the conditions of temperature and pressure at which a substance exists
as a solid, liquid, or a gas.
The point at which all the three curves for the solid, liquid and gaseous states of a substance meet in a
phase diagram is the Triple point. This is the only condition at which all the three phases of a
substance can be in equilibrium.
A phase diagram helps to predict changes in the melting and boiling points of a substance as a result
of changes in external pressure.