The Condensed Phase
The kinetic theory of gases presents a microscopic model for
the behavior of gases.
As pressure increases or temperature decreases, gas
molecules begin to feel the presence of other gas molecules
and interactions between them cannot be ignored.
As pressure on a gas is increased, or the temperature of the
gas is decreased, the gas liquefies forming a more ordered
phase, the liquid phase. Further increase in P or drop in T
results in a still more ordered solid phase.
Non-bonding Interactions between molecules:
Intermolecular interactions
NON-BONDING interactions are weaker than bonding
interactions (covalent and ionic)
Yet, these interactions are critical in defining physical
properties of compounds.
Potential energy (kJ/mol)
The relative energy of two molecules interacting with each
other can be plotted as a function of the distance between
the two molecules - Potential energy curves.
Separation (Å)
Types of Nonbonding Interactions
1) Dipole-Dipole interactions
Interactions between polar molecules - e.g. between two H2O
molecules
2) Ion-Dipole interactions
Interactions between an ion and a polar molecule - e.g.
dissolution of NaCl in water.
3) Induced Dipole interaction
Example: a water molecule approaching an O2 molecule, can
induce a temporary dipole in the O2
4) Dispersion forces or van der Waals interactions
Example: interaction between two H2 molecules
When two non-polar molecules approach one another, each
can influence the electron distribution in the other to a small
extent.
A small fluctuation in the electron distribution around one of
the molecules, will, at close distance, affect the electron
distribution on the neighboring molecule.
Hydrogen-bonding: a special case of dipole-dipole interaction
Hydrogen bonds form when a H atom is covalently bonded to
a N, O, or F atom and interacts with the lone electron pair on
the N, O or F atom in an adjacent molecule.
H
H O H
H F
H N H
water
Hydrogen fluoride
ammonia
Hydrogen bonds affect physical properties of a molecule,
boiling point (oC)
H2O
100
HF
20
NH3
-34
HCl
-85
CH4
-161
Hydrogen bonds in liquid water
Consequences of hydrogen bonding in water
Ice floats because hydrogen bonds hold water molecules
further apart in a solid than in a liquid - density of ice is
less than density of water
Density of ice at 0oC - 0.9997 g/ml
Density of water at 0oC - 0.9170 g/ml
liquid water
solid water
http://www.nyu.edu/pages/mathmol/modules/water/info_water.html
Water has a high specific heat index.
It takes much more heat to raise the
temperature of a volume of water than the
same volume of air.
Some Consequences:
Water is used as a coolant
Effects global climates and rates of global
climate change
- changes in temperatures are gradual
Water has a high surface tension
surface tension (dynes/cm at 20oC)
Water
73
Methanol
22
Ethanol
22
Ether
17
The surface tension makes
air-water boundaries
distinctive microhabitats.
“Universal” Solvent
Water can dissolve ionic and polar compounds
Polar compounds in water
H-bonding defines the shape of the molecule
for example, the overall shape of proteins, the doublehelix in DNA.
H-bonding in DNA
http://michele.usc.edu/105b/biochemistry/dna.html
http://www.umass.edu/microbio/chime/dna/fs_pairs.htm
Liquid Crystals
http://invsee.asu.edu/nmodules/liquidmod/spatial.html
Vaporization and Condensation
Molecules in a liquid are in constant motion; some moving
faster, others slower.
Those molecules with enough kinetic energy escape from the
liquid surface, i.e. vaporize.
Molecules with higher energy are able to overcome
interactions between other molecules
If the container is kept open, vaporization continues until no
more liquid remains; molecules escape from the liquid and
heat flows in from the surroundings, replacing the energy
lost to vaporization and maintaining the rate of vaporization.
Condensation: When molecules in the gas phase collide with
the liquid surface, they loose energy and return to the liquid.
At some point the rate of vaporization and the rate of
condensation become equal and the system is at equilibrium
(a dynamic equilibrium)
The partial pressure of the vapor above the liquid established
at equilibrium is called the equilibrium vapor pressure or the
vapor pressure.
Boiling Point - the temperature at which the vapor pressure of
the liquid equals the atmospheric pressure.
Normal boiling point - temperature at which the vapor
pressure equals 1 atm.
If the external pressure is reduced, the boiling point
decreases (e.g. at high altitudes).
If the external pressure increases,the boiling point increases
(e.g. a pressure cooker).
Melting point - temperature at which a substance turns from
solid to liquid.
At the melting point, the solid and liquid are in equilibrium
and co-exist at this temperature
Phase Transitions
When a compound changes its state from a solid to a liquid or
a liquid to a gas, it is said to have undergone a phase change
or a phase transition.
Changes in temperature and pressure cause phase transitions
Fusion or melting
Vaporization
Sublimation
solid --> liquid
liquid --> gas
solid --> gas
At the melting point, boiling point or sublimation point of the
substance, temperature remains the same even if the sample
is heated
These points correspond to phase changes, and the energy
supplied is being used by the substance to undergo the
phase change.
Once the phase change is complete, and if heat is still
applied, then the temperature increases.
Phase Diagrams
Plots of pressure vs temperature showing changes in the
phase of a substance is called a PHASE DIAGRAM.
liquid
solid
gas
liquid
solid
gas
1) Any point along a curve represents an equilibrium
between two phases. Any point not on a curve
corresponds to a single phase.
2)The line from A to B is the vapor pressure curve of the
liquid. The vapor pressure ends at the critical point (B),
beyond which a gas cannot be compressed to a liquid - a
supercritical fluid exists.
liquid
solid
gas
3) Line from A to C represents variation in the vapor pressure
of the solid as it sublimes at different temperatures.
4) Line from A to D represents change in melting point of the
solid with increasing pressure
5) Point A, where the three curves intersect is called the
TRIPLE point, which corresponds to the pressure and
temperature at which all three phases coexist.
Phase diagram of water and CO2
Note: for CO2 freezing point increases with increasing
pressure, but for H2O freezing point decreases with
increasing pressure.