The Condensed Phase

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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.

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• 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.

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• 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.

Potential energy (kJ/mol)
Separation (Å)
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• 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.
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3) Induced Dipole interaction Example a water
molecule approaching an O2 molecule, can induce a
temporary dipole in the O2
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• 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.

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• 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.

water
Hydrogen fluoride
ammonia
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Hydrogen bonds affect physical properties of a
molecule,
boiling point (oC) H2O 100 HF 20 NH3
-34 HCl -85 CH4 -161
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Hydrogen bonds in liquid water
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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
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liquid water
solid water
http//www.nyu.edu/pages/mathmol/modules/water/inf
o_water.html
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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
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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.
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Universal Solvent Water can dissolve ionic and
polar compounds
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Polar compounds in water
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H-bonding defines the shape of the molecule for
example, the overall shape of proteins, the
double-helix in DNA.
H-bonding in DNA
http//michele.usc.edu/105b/biochemistry/dna.html
http//www.umass.edu/microbio/chime/dna/fs_pairs.h
tm
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Liquid Crystals
http//invsee.asu.edu/nmodules/liquidmod/spatial.h
tml
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• 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.

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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.
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• 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.

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• 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

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• 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

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• Fusion or melting solid --gt liquid
• Vaporization liquid --gt gas
• Sublimation solid --gt 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.

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Phase Diagrams Plots of pressure vs temperature
showing changes in the phase of a substance is
called a PHASE DIAGRAM.
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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.
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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.
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• Phase diagram of water and CO2

Note for CO2 freezing point increases with
increasing pressure, but for H2O freezing point
decreases with increasing pressure.