INTERMOLECULAR FORCES, LIQUIDS, & SOLIDS
AIM
Explore the relationship among structure, intermolecular forces, and
physical properties of matter.
CONTENTS
11.1- A molecular Comparison of Liquids and Solids
11.2- Intermolecular Forces
11.3- Some Properties of liquids
11.4- Phase Changes
11.5- Vapor Pressure
11.6- Phase Diagrams
11.7- Structures of Solids
11.8- Bonding in Solids
 REVIEW QUESTIONS AND PROBLEMS
11.1- A Molecular Comparison of Liquids & Solids

Liquid molecules are held closer together than gas molecules, but
not so rigidly that the molecules can slide past each other.
 Solid molecules are packed closely together. Molecules are rigidly
packed that they cannot easily slide past each other.
 Converting a gas into a liquid or solid requires the molecules to
get closer to each other: cool or compress.
 Converting a solid into a liquid or gas requires the molecules to
move further apart to each other: heat or decrease the pressure
11.1 INTERMOLECULAR FORCES
Intermolecular forces are the forces that exist between molecules;
they are the attractive forces that hold together the particles of a
liquid or a solid.
Intermolecular forces are typically much weaker than covalent or
ionic bonds (intramolecular forces);
So it takes much more energy to break the covalent bond between H
and Cl than it does to vaporize HCl
 When a substance undergoes a physical change, the intermolecular
forces are broken and not the intramolecular forces.
When a substance condenses the intermolecular forces are formed.
The strength of the intermolecular forces in a compound are
reflected in its melting point and boiling point. The stronger the
intermolecular forces, the more energy is required to overcome those
forces
Forces that hold together molecules and atoms (both electrically
neutral species) are referred to collectively as van der Waals forces.
 Attractive forces include dipole-dipole attractive forces, hydrogen
bonding, and London dispersion forces.
Ion-Dipole Forces
An ion-dipole force exists between an ion and the partial charge on
the end of a polar molecule.
The positive ends will orient themselves toward a negative ion, and
the negative ends will orient themselves toward a positive ion. As
with all electrostatic attraction, the magnitude of the attraction
increases with the size of the charge (both on the ion and on the
dipole).

Ion-Dipole Forces are the strongest of all intermolecular forces.
Dipole-Dipole Forces
Dipole-Dipole Forces: exist between neutral polar molecules. The
positive end of one polar molecule (dipole) is attracted to the negative
end of another. The greater the polarity of the molecules, the
stronger the attractions between them.
 For the attraction to be greater the dipoles need to be close
together.
 There is a mix of attractive and repulsive forces as the molecules
tumble. The overall effect is a net attraction
For molecules of approximately the same size, boiling point increases
with increasing dipole moment:
 Dipole Dipole forces are weaker than ion-dipole forces
London Dispersion Forces
There are many nonpolar substances exist as condensed phases.
The intermolecular forces holding nonpolar atoms or molecules
together are known as London dispersion forces.
 It is possible for two adjacent molecules or atoms to affect each
other. The nucleus of one molecule or atom attracts the electron
clouds of the adjacent molecule or atom. We say the molecule has
become polarized.
This creates an instantaneous dipole, which can induce an other
instantaneous dipole in the adjacent molecule. Those dipoles attract
each other. This attractive force is called London dispersion force.
 Polarizability is the ease with which an electron cloud can be
deformed.
 The larger the molecule, the greater the polarizability, the more
easily the electron cloud can be distorted to give a momentary dipole
which leads to a stronger London dispersion forces.
London dispesion forces increase as the molecular weights increase
boiling points increase

London dispesion forces exist in all molecules.
 London dispersion forces depends on the shape of the molecule
(VSEPR)
The greater the surface area available for contact, the greater the
dispersion forces.
Hydrogen Bonding
Hydrogen bonding occurs when hydrogen is bound to a small, very
electronegative atom. (The most significant hydrogen bonding occurs
when H is covalently bonded to F, O, N.
The electronegative atom effectively strips hydrogen of its only
electron, leaving a nearly unshielded proton exposed. This proton is
powerfully attracted to the atoms on neighboring molecules that have
unshared electron pairs. Hydrogen bonds significantly impact the
properties of some substances
 Experimentally the boiling points of compounds with H-F, H-O,
and H-N bonds are abnormally high
 The hydrogen bonds are 5-10% as strong as covalent bonds.
 Hydrogen bonds are generally stronger than dipole-dipole
interaction or dispersions forces.
 Solids are usually more closely packed than liquidsSolids are
more dense than liquids
Exception:
 Ice is ordered with an opened structure(regular hexagon) to
optimize H-bonding  Therefore, ice is less dense than water
Comparing Intermolecular Forces
 Despersion forces are found in all substances.
 Dipole-dipole forces and hydrogen bonds add to the effect of
dispersion forces.
 None of the intermolecular forces are stronger than ionic or
covalent bonds.
11.3 Some Properties of Liquids
 Viscosity is the resistance of a liquid to flow
 The stronger the intermolecular forces, the higher the viscosity.
 The viscosity decreases with increasing temperature.
Surface Tension
 Bulk molecules (those in the liquid) are equally attracted to their
neighbors.
 Surfaces molecules are attracted sidewises and inwards towards
the bulk molecules  Therefore, surface molecules are packed more
closely than the bulk molecules causing the liquid to behaves like a
skin.
 Surface tension is the amount of energy required to increase the
surface area of a liquid.
 Cohesive forces bind the molecules to each other
 Adhesive forces bind the molecules to surface of the container
 If the adhesive forces are large than cohesive forces, the liquid is
attracted to the surface of the container more than the bulk
molecules. Therefore, the meniscus is U-shaped (water in glass)
 If the cohesive forces are greater than adhesive forces, the liquid is
attracted to the bulk molecules more the surface of the container.
Therefore, the meniscus is downward (mercury in glass).
11.4 Phase Changes
Energy Changes Accompanying Phase Changes
 Energy change of the system for the above processes are:

Sublimation:
sub > 0 (endothermic).

Vaporization: Hvap > 0 (endothermic).

Melting or Fusion: fus > 0 (endothermic).

Deposition:dep < 0 (exothermic).

Condensation:Hcon < 0 (exothermic).

Freezing: Hfre < 0 (exothermic).
 Generally heat of fusion (enthalpy of fusion) is less than heat of
vaporization: it takes more energy to completely separate
molecules, than partially separate them.
 The sequence
solid + heat
liquid (fusion)
liquid + heat
gas (vaporization)
solid + heat gas (sublimation)
is endothermic.
 The sequence
gas (cool)
liquid + heat (condensation)
Liquid (cool)
solid +heat (freezing)
gas (cool)
solid + heat (deposition)
is exothermic.
Heating Curves
Plot of temperature changes versus heat added is a heating curve.
 During a phase change, adding heat causes no temperature
change.
These points are used to calculate Hfus and Hvap
 Supercooling: When a liquid is cooled below its melting point and
it still remains a liquid.
11.5 Vapor Pressure
Explaining Vapor Pressure on the Molecular Level
 Some of the molecules on the surface of a liquid have enough
energy to escape the attraction of the bulk liquid. These molecules
move into the gas phase. As the number of molecules in the gas phase
increases, some of the gas phase molecules strike the surface and
return to the liquid.
 Dynamic Equilibrium occurs when as many molecules escape the
surface as strike the surface.
 A liquid and its vapor are in dynamic equilibrium when
evaporation and condensation occur at equal rates.
Vapor pressure is the pressure exerted when the liquid and
vapor are in dynamic equilibrium.
Volatility, Vapor Pressure, and Temperature
 If equilibrium is never established then the liquid evaporates (e.g
water in an open container).
 Volatile substances evaporate rapidly (e.g gasoline).

 The higher the temperature, the higher the average kinetic
energy, the faster the liquid evaporates.
Vapor Pressure and Boiling Point
 Liquids boil when the external pressure acting on the surface of
the liquids equals the vapor pressure.
 Temperature of boiling point increases as pressure increases.
 Pressure cookers operate at high pressure. At high pressure the
boiling point of water is higher than 1 atm. Therefore, there is a
higher temperature at which the food would be cooked.
11.6 Phase Diagrams
 Phase diagram: plot of Pressure vs. Temperature
summarizing all equilibria between phases.
 Features of a phase diagram:
Triple point: temperature and pressure at which all three
phases are in equilibrium.
Vapor-pressure curve: generally as pressure increases,
temperature increases.
Critical point: Critical Temperature and Pressure
 Critical temperature: A temperature above which the gas cannot
be liqufefy regardless of the pressure.
 Critical pressure: pressure required for liquefaction.
Melting point curve: as pressure increases, the solid phase is
favored if the solid is more dense than the liquid.
Normal melting point: melting point at 1 atm.
 Any temperature and pressure combination not on a curve
represents a single phase.
Phase Diagrams of H2O and CO2
For H2O
 The melting point curve slopes to the left because ice is less dense
than water.
 Triple point occurs at 0.0098o C and 4.58 mmHg.
 Normal melting (freezing) point is 0o C.
 Normal boiling point is 100 oC.
 Critical point is 374o C and 218 atm.
For CO2
 Triple point occurs at -56.4o C and 5.11 atm.
 Normal sublimation point is -78.5o C (At 1 atm CO2 sublimes it
does not melt.)
 Critical point occurs at 31.1o C and 73 atm.
11.7 Structures of Solids
 Crystalline solid: well-ordered, definite arrangements of
molecules, atoms or ions.
Examples: quartz, salt, sugar.
Tend to melt at specific temperatures.
Therefore, have a narrow range of intermolecular forces.
 Amorphous solid: molecules, atoms or ions do not have an orderly
arrangement.
Examples: rubber, glass.
Tend to melt over a range of temperatures.
Therefore, amorphous solids have variable intermolecular
forces.
Unit Cells
 Crystals have an ordered, repeated structure.
 The smallest repeating unit in a crystal is a unit cell.
 Unit cell is the smallest unit with all the symmetry of the entire
crystal.
Three dimensional stacking of unit cells is the crystal lattice.
 Three types of unit cell .
1. Primitive cubic, atoms at the corners of a simple cube,
each atom shared by 8 unit cells;
2. Body-centered cubic (bcc), atoms at the corners of a
cube
plus one in the center of the body of the cube,
corner atoms shared by 8 unit cells, center atom
completely enclosed in one unit cell;
3. Face-centered cubic (fcc), atoms at the corners of a cube plus one
atom in the center of each face of the cube,corner atoms shared by 8
unit cells, face atoms shared by 2 unit cells.
 Only in the sample cubic that the atoms at the corners of the cube
touch each other.
 In the body centered cubic cell atom contact is a body diagonal.
 In the face centered cubic atoms touche a face diagonal.
Close Packing of Spheres
 Solids have maximum intermolecular forces.
 Molecules, atoms or ions can be modeled by spheres.
 Molecular crystals are formed by close packing of the molecules.
 We rationalize maximum intermolecular force in a crystal by the
close packing of spheres.
 When spheres are packed as closely as possible, there are small
spaces between adjacent spheres called interstitial holes.
 A crystal is built up by placing close packed layers of spheres on
top of each other.
 If unequally sized spheres are used, the smaller spheres are placed
 in the interstitial holes
X-Ray Diffraction
 When waves of X rays hit a crystal , the electron’s clouds of the
atoms of the crystal are diffracted giving a diffraction pattern
 Efficient diffraction occurs when the wavelength of light is close to
the size of the atoms of the crystal.
 X-rays are scattered by atoms  interference (constructive or
destructive).
 The spacing between layers in a crystal is 2 - 20 Ao , which is the
wavelength range for X-rays.
 X-ray diffraction (X-ray crystallography):
X-rays are passed through the crystal and are detected on a
photographic plate.
 The photographic plate has one bright spot at the center (incident
beam) as well as a diffraction pattern.
 Each close packing arrangement produces a different diffraction
pattern.
 Knowing the diffraction pattern, we can calculate the positions of
the atoms required to produce that pattern.
 We calculate the crystal structure based on a knowledge of the
diffraction pattern.
 Bragg equation:
2 d sin = n 
d: spacing between 2 successive layers of particles : angle at which X
rays enter & leave the layers
n: diffraction order (an integer)
: wavelength of the X rays
 The Bragg equation serves as the basis for the study of crystalline
structures by XRD.
 Method: measure  and calculate d then compare structure of the
solid
Intensities of diffracted rays  atomic positions
11.8 Bonding in Solids
There are four types of solids:
1. Molecular solids
Molecular (formed from molecules) - usually soft with low melting
points and poor conductivity.
 Intermolecular forces: dipole-dipole, London dispersion and
 H-bonds.
 Weak intermolecular forces give rise to low melting points.
 Room temperature gases and liquids usually form molecular
solids and low temperature.

2.Covalent-Network Solids
(formed from atoms) - very hard with very high melting points and
poor conductivity.
 Atoms held together in large networks.
 Examples: diamond, graphite, quartz (SiO2 ), silicon carbide
(SiC), and boron nitride (BN).
In diamond:
 each C atom has a coordination number of 4;
 each C atom is tetrahedral;
 there is a three-dimensional array of atoms.
 Diamond is hard, and has a high melting point (3550o C).
In Graphite
 each C atom is arranged in a planar hexagonal ring;
 layers of interconnected rings are placed on top of each other;
 the distance between C atoms is close to benzene (1.42 Ao vs. 1.395
Ao in benzene); the distance between layers is large (3.41 Ao );
 electrons move in delocalized orbitals (good conductor).

3.Ionic Solids
(formed form ions ) - hard, brittle, high melting points and poor
conductivity .
F = k Q1Q2
d2
 Ions (spherical) held together by electrostatic forces of attraction:
The higher the charge (Q) and smaller the distance (d) between ions,
the stronger the ionic bond.
 There are some simple classifications for ionic lattice types:
NaCl Structure
 Each ion has a coordination number of 6.
 Face-centered cubic lattice.
 Cation to anion ratio is 1:1.
 Examples: LiF, KCl, AgCl and CaO.
CsCl Structure
 Cs has a coordination number of 8.
 Different fron the NaCl structure (Cs+ is larger than Na+).
 Cation to anion ratio is 1:1.
Zinc Blende Structure
 Typical example ZnS.
 S2- ions adopt a fcc arrangement.
 Zn2+ ions have a coordination number of 4.
 The S2- ions are placed in a tetrahedron around the Zn2+ ions.
 Example: CuCl.
Fluorite Structure
 Typical example CaF2 Ca2+ ions in a fcc arrangement.
 There are twice as many F- per Ca2+ ions in each unit cell.
 Examples: BaCl2, PbF2
4. Metallic Solids
Meltallic Solids (formed from metal atoms) - soft or hard, high
melting points, good conductivity, malleable and ductile
 Metallic solids have metal atoms in hcp, fcc or bcc arrangements.
 Coordination number for each atom is either 8 or 12.
 Problem: the bonding is too strong for London dispersion and
there are not enough electrons for covalent bonds.
 Resolution: the metal nuclei float in a sea of electrons.
Metals conduct because the electons are delocalized and are mobile.
