Liquids, Solids, and Phase Changes

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Liquids, Solids, and Phase Changes

• Chapter 10

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Polar Molecules

• Dipole - A molecule such as HF which has a
  positive and a negative end. This dipolar
  character is often represented by an arrow
  pointing towards the negative charge.
• Dipole moments the measure of the net
  molecular polarity
• Measured in units of Debyes (D) Qr (charge x
  separation)

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• Non polar molecules have no net dipole moment
• Polar molecules have a dipole moment

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Intermolecular Forces

• Forces holding one molecule to another in a
  substance.
• van der Waals forces
• Several types ion-dipole, dipole-dipole,
  dipole-induced dipole and London dispersion
  forces
• Hydrogen bonding

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Molecular interactions

• Ion - Dipole forces
• Dipole Dipole forces or Polar - Polar
  interactions.
• Dipole Induced Dipole or Polar - Nonpolar
  interactions
• London forces Induced dipole induced dipole or
  Nonpolar - Nonpolar interactions

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Strength of Attractions

• Polar polar
• gt
• Polar nonpolar
• gt
• Nonpolar nonpolar

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How does polarity affect molecular properties?
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Melting point and boiling points are higher for
polar molecules than for nonpolar molecules of
similar molecular weight.
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Solubility

• Like dissolves Like

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Solubility

• Polar solvents dissolve polar molecules
• Nonpolar solvents dissolve nonpolar molecules
• Molecules with polar and nonpolar ends are
  frequently soluble in both polar and nonpolar
  solvents.
• Polar solvents are good for solubilizing salts.

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Hydrogen Bonds

• A special type of polar interaction between a
  hydrogen atom bonded to an electronegative
  element and another electronegative element.

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Liquids
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Viscosity

• resistance to flow
• If a liquid has strong intermolecular
  interactions then particles will not flow past
  each other easily and viscosity will be high.

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Surface Tension

• tendency to minimize surface area
• Adhesion Forces that bind a substance to a
  surface
• Cohesion Forces that bind a substance to itself

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Phase Transitions

• Vaporization Liquid to gas transition
• Melting Solid to liquid transition
• Condensation gas to liquid transition
• Freezing liquid to solid transition
• Sublimation solid to gas transition
• Deposition gas to solid transition

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• Dry ice sublimes at 78oC and has a ?Hsub of 25.2
  kJ/mol. Calculate the vapor pressure of CO2 at
  100oC.

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• Calculate the boiling point of water at the
  summit of Pikes Peak in Colorado where the
  atmospheric pressure is 447.

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• Crystalline solid atoms, ions, or molecules lie
  in an orderly array
• typically have flat well defined surfaces called
  faces.
• Amorphous solid atoms or molecules lie in
  random jumble.

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Phase Diagrams

• Graphically show conditions under which all
  phases are stable.

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Phase diagram for CO2
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Triple point

• A three-way intersection representing the unique
  temp, pressure where all three phases exist
  simultaneously.

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Critical point

• The temperature (critical temperature) where a
  gas cannot be liquefied no matter what the
  pressure.
• Notice that the triple point for carbon dioxide
  is at 5.11 atm pressure. The liquid form doesnt
  ever exist below this pressure and this is the
  reason that dry ice never melts but always
  sublimes.

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Supercritical fluid

• Neither a liquid nor a gas. The liquid and gas
  forms become indistinguishable at this point.

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Phase diagram for H2O
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Water

• colorless, odorless, tasteless, liquid at
  ordinary temperatures
• only inorganic compound occurring naturally as a
  liquid
• composes ?65 of mass of living organisms
• excellent solvent for many things
• abnormally high boiling and melting point
• ice is less dense than water (it floats)

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Water purification

• Hard water -- Contains Ca2, Mg2, Fe3 and other
  minerals.
• Soft water -- Doesnt contain Ca2, Mg2, Fe3
  ions.
• Softened water -- metal cations in hard water are
  replaced by Na.
• Deionized water -- cations are replaced by H and
  anions are replaced by OH-