CHEMISTRY
The Central Science
8th Edition
Chapter 11
Intermolecular Forces, Liquids
and Solids
P. Hatch
A Molecular Comparison
of Liquids and Solids
A Molecular Comparison
of Liquids and Solids
• Physical properties understood in terms of kinetic
molecular theory:
– Gas molecules are far apart and do not interact much
with each other.
– Liquid molecules are held closer together than gas
molecules, but not so rigidly that the molecules cannot
slide past each other.
– Solid molecules are packed closely together. The molecules
are so rigidly packed that they cannot easily slide past each
other.
A Molecular Comparison
of Liquids and Solids
A Molecular Comparison
of Liquids and Solids
A Molecular Comparison
of Liquids and Solids
• 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:
– heat or reduce pressure.
• The forces holding solids and liquids together are called
intermolecular forces (“inter-” means between).
Intermolecular Forces
• The covalent bond holding a molecule together is an
intramolecular force (“intra-” means within).
O
O
H
H
H
H
• The attraction between molecules is an intermolecular
force (“inter-” means between).
H
H
-
+
O
O
+
-
H
H
Intermolecular Forces
• Intermolecular forces are much weaker than
intramolecular forces (covalent bonds).
• When a substance melts or boils the intermolecular
forces are broken (not the covalent bonds, intramolecular
forces).
Intermolecular Forces
Intermolecular Forces
Four types of intermolecular forces:
Ion-dipole forces:
- Exists between an ion and a polar molecule
Dipole-dipole forces:
- Exists between neutral polar molecules
London dispersion forces:
- Exists between all atoms and molecules, polar, nonpolar
Hydrogen bonds:
- Exists between the H atom in a polar bond and an
electronegative ion or atom (F, O, N).
Ion-Dipole Forces
• Interaction between an ion and a dipole (ex. NaCl in water).
• Strongest of all intermolecular forces.
Dipole-Dipole Forces
• Dipole-dipole forces exist
between neutral polar
molecules. Ex. HCl
• Weaker than ion-dipole
forces.
• If two molecules have about
the same mass and size,
then dipole-dipole forces
increase with increasing
polarity.
Intermolecular Forces
Dipole-Dipole Forces
London Dispersion Forces
• Weakest of all intermolecular forces.
• The nucleus of one molecule (or atom) attracts the electrons
of the adjacent molecule (or atom).
• In that instant a dipole is formed (called an instantaneous
dipole).
London Dispersion Forces
London Dispersion (LD) Forces
• Polarizability is the ease with which an electron cloud can be
deformed; the larger the molecule the more polarizable.
• LD forces increase as molecular weight increases.
• LD forces depend on the shape of the molecule.
greater the surface area available for contact, the greater
the dispersion forces.
LD forces between spherical molecules are lower than
between sausage-like molecules. WHY?
Intermolecular Forces
London Dispersion
Forces
London Dispersion Forces
London Dispersion Forces
Hydrogen Bonding
Hydrogen Bonding
• H-bonding occurs when H bonded to F, O, or N.
creates a polar bond with d+ on H
• Special case of dipole-dipole forces.
• By experiments: boiling points of compounds with H-F,
H-O, and H-N bonds are abnormally high. Why?
• H-bonds are abnormally strong.
Hydrogen Bonding
Hydrogen Bonding
Hydrogen Bonding
Hydrogen Bonding
• Hydrogen bonds are responsible for:
– Ice Floating
• Solids are usually more closely packed than liquids;
• Therefore, solids are more dense than liquids.
• Ice is ordered with an open structure to optimize H-bonding.
• Therefore, ice is less dense than water.
• In water the H-O bond length is 1.0 Å.
• The O…H hydrogen bond length is 1.8 Å.
• Ice has waters arranged in an open, regular hexagon.
• Each d+ H points towards a lone pair on O.
Intermolecular Forces
Some Properties of Liquids
Viscosity
• Viscosity: the resistance of a liquid to flow.
• A liquid flows by sliding molecules over each other.
• The stronger the intermolecular forces, the higher the
viscosity.
Some Properties of Liquids
Surface Tension
• Surface molecules are only attracted inwards towards the
bulk molecules.
– Therefore, surface molecules are packed more closely than bulk
molecules.
• Surface tension: the amount of energy required to
increase the surface area of a liquid.
• Cohesive forces bind molecules to each other.
• Adhesive forces bind molecules to a surface (ex. glass).
Surface Tension
Some Properties of Liquids
Surface Tension
• Meniscus is the shape of the liquid surface.
– Adhesive forces > Cohesive forces, the liquid surface is
attracted to its container more than the bulk molecules and
meniscus is U-shaped (e.g. water in glass).
– Cohesive forces > Adhesive forces, meniscus is curved
downwards.
• Capillary Action: When a narrow glass tube is placed in
water, the meniscus pulls the water up the tube.
Phase Changes
•
•
•
•
•
•
Sublimation:
Vaporization:
Melting or fusion:
Deposition:
Condensation:
Freezing:
solid  gas.
liquid  gas.
solid  liquid.
gas  solid.
gas  liquid.
liquid  solid.
Phase Changes
Phase Changes
•
•
•
•
•
•
Energy Changes Accompanying
Phase Changes
Sublimation: Hsub > 0 (endothermic).
Vaporization: Hvap > 0 (endothermic).
Melting or Fusion: Hfus > 0 (endothermic).
Deposition: Hdep < 0 (exothermic).
Condensation: Hcon < 0 (exothermic).
Freezing: Hfre < 0 (exothermic).
Phase Changes
Energy Changes Accompanying
Phase Changes
• In general Hfus < Hvap
• it takes more energy to completely separate molecules,
than partially separate them.
Phase Changes
Phase Changes
Energy Changes Accompanying
Phase Changes
• All phase changes are possible under the right conditions.
• The sequence
heat solid  melt  heat liquid  boil  heat gas
is endothermic.
• The sequence
cool gas  condense  cool liquid  freeze  cool solid
is exothermic.
Heating Curves
Heating Curves
• Plot of temperature change 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.
achieved by keeping the temperature low and
increasing kinetic energy to break intermolecular
forces.
Phase Changes
Critical Temperature and Pressure
• Gases liquefied by increasing pressure at some
temperature.
• Critical temperature: the maximum temperature for
liquefaction of a gas using pressure.
• Critical pressure: pressure required for liquefaction.
Phase Changes
Critical Temperature and Pressure
Exe. 11.33, p. 429
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.
After some time the pressure of the gas will be constant
at the vapor pressure.
Vapor Pressure
Explaining Vapor Pressure on the
Molecular Level
Vapor Pressure
•
•
•
•
Explaining Vapor Pressure on the
Molecular Level
Dynamic Equilibrium: the point when as many molecules
escape the surface as strike the surface.
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.
Volatile substances evaporate rapidly.
Vapor Pressure
Volatility, Vapor Pressure, and
Temperature
• The higher the temperature, the higher the average kinetic
energy, the faster the liquid evaporates.
Vapor Pressure
Volatility, Vapor Pressure, and Temperature
Vapor Pressure
Vapor Pressure and Boiling Point
• Liquids boil when the external pressure equals the vapor
pressure.
• Temperature of boiling point increases as pressure
increases.
Vapor Pressure
Vapor Pressure and Boiling Point
• Two ways to get a liquid to boil: increase temperature or
decrease pressure.
– Pressure cookers operate at high pressure. At high pressure the
boiling point of water is higher than at 1 atm. Therefore, there
is a higher temperature at which the food is cooked, reducing
the cooking time required.
• Normal boiling point is the boiling point at 760 mmHg (1
atm).
Phase Diagrams
• Phase diagram: plot of pressure vs. Temperature
summarizing all equilibria between phases.
• Given a temperature and pressure, phase diagrams tell us
which phase will exist.
• Any temperature and pressure combination not on a
curve represents a single phase.
Phase Diagrams
• 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 for the gas.
– 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.
Phase Diagrams
Phase Diagrams
The Phase Diagrams of H2O and CO2
Phase Diagrams
The Phase Diagrams of H2O and CO2
• Water:
– The melting point curve slopes to the left because ice is less
dense than water.
– Triple point occurs at 0.0098C and 4.58 mmHg.
– Normal melting (freezing) point is 0C.
– Normal boiling point is 100C.
– Critical point is 374C and 218 atm.
Phase Diagrams
The Phase Diagrams of H2O and CO2
• Carbon Dioxide:
– Triple point occurs at -56.4C and 5.11 atm.
– Normal sublimation point is -78.5C. (At 1 atm CO2 sublimes
it does not melt.)
– Critical point occurs at 31.1C and 73 atm.
Bonding in Solids
• There are four types of solid:
 Molecular (formed from molecules) - usually soft with
low melting points and poor conductivity.
Covalent network (formed from atoms) - very hard with
very high melting points and poor conductivity.
Ions (formed form ions) - hard, brittle, high melting points
and poor conductivity.
Metallic (formed from metal atoms) - soft or hard, high
melting points, good conductivity, malleable and ductile.
Bonding in Solids
Molecular Solids
• Examples: C12H11O22,
• 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 at low temperature.
Bonding in Solids
•
•
•
•
Covalent-Network Solids
Intermolecular forces: dipole-dipole, London dispersion
and H-bonds.
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 (3550 C).
Bonding in Solids
Covalent-Network Solids
Bonding in Solids
Covalent-Network Solids
• 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 Å vs.
1.395 Å in benzene);
– the distance between layers is large (3.41 Å);
– electrons move in delocalized orbitals (good conductor).
Bonding in Solids
Ionic Solids
• Ions (spherical) held together by electrostatic forces of
attraction.
• There are some simple classifications for ionic lattice
types.
Ionic Solids
Bonding in Solids
Ionic Solids
• 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 from the NaCl structure (Cs+ is larger than Na+).
• Cation to anion ratio is 1:1.
Bonding in Solids
Ionic Solids
• 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.
Bonding in Solids
Ionic Solids
• 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.
Bonding in Solids
•
•
•
•
•
Metallic Solids
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 electrons are delocalized and
are mobile.
End of Chapter 11
Intermolecular Forces, Liquids
and Solids