Intermolecular Forces, Gases, and Liquids

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Intermolecular Forces, Gases, and Liquids

• Ch.13

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Gases

• Kinetic-Molecular Theory says molecules/atoms
  separated
• Little, if any, interactions
• Not so in solids and liquids
• Examples
• Big difference in volume between liquids solids
  and gases
• Gases compressible, liqs solids not

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Intermolecular Forces

• Various electrostatic forces that attract
  molecules in solids/liqs
• Much weaker than ionic forces
• About 15 (or less) that of bond energies
• Remember, ionic bonds extremely powerful
• Boiling pt of NaCl 1465 C!

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Intermolecular Forces

• Reason behind importance of knowing about IMF
• 1) b.p. m.p. and heats of vaporization (l?g)
  and fusion (s?l)
• 2) solubility of gases, liquids, and solids
• 3) determining structures of biochemicals (DNA,
  proteins)

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Remember dipole moments?

• Dipole moment product of magnitude of partial
  charges (?/?-) their distance of separation
• (1 Debye 3.34 x 10-30 C x m)
• Important in IMF

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Ion-dipole Ionization in aqueous medium (water)

• 1) stronger attraction if ion/dipole closer
• Li vs. Cs in water
• 2) higher ion charge, stronger attraction
• Be2 vs. Li in water
• 3) greater dipole, stronger attraction
• Dissolved salt has stronger attraction to water
  than methanol

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Solvation energy

• Or, enthalpy of hydration (if water) energy of
  ionization in aq. media
• Water molecules surround both ions
• Example
• Take hydration energies of G I metal ions
• Exothermicity decreases as you go down the column
• Cations become larger
• Easier to dissociate

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Permanent dipoles

• Positive end of one molecule attracted to
  negative end of other
• For ex HCl
• Dipole-dipole attractions
• Cmpds that exhibit greater d-d attractions have
  higher b.p., and Hvap
• Polar cmpds exhibit greater d-d attractions than
  non-polar cmpds
• NH3 vs. CH4
• ? equivalent molar masses (g/mol) 17 vs. 16,
  respectively
• Boiling points -33C vs. -162C, respectively

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Hydrogen Bonding

• A type of super dipole-dipole interaction
• Interaction between e--rich atom connected to H
  entity another H attached to erich atom
• e--rich atom O, F, N
• Density water gt than ice
• Opposite of almost every other substance
• Inordinately high heat capacity of water
• High surface tension
• Insects walk on water
• Concave meniscus

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Hydrogen Bonding

• Boiling pts. of H2O, HF, and NH3 much higher

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Surface Tension

• Outer molecules interact with surface, while
  inner interact with other molecules
• It has a skin
• Skin toughness surface tension
• Energy required to break through surface
• Smaller surface area reason that water drops
  spherical

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Capillary Action

• When water goes up a small glass tube
• Due to polarity of Si-O bonding with water
• Adhesive forces gt cohesive forces of water
• Creates a chain or bridge
• Pulls water up tube
• Limited by balancing gravity with
  adhesive/cohesive forces
• Thus, water has a concave meniscus

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Mercury

• Forms a convex meniscus
• Doesnt climb a glass tube
• Due to cohesive forces gt adhesive forces

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Viscosity

• Hydrogen-bonding increases viscosity
• But large non-polar liquids like oil have
• 1) large unwieldy molecules w/greater
  intermolecular forces
• 2) greater ability to be entangled w/one another
• Did you ever hear the expression, Youre as slow
  as molasses in January?

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Dipole/Induced Dipole Forces

• Polar entities induce dipole in nonpolar species
  like O2
• O2 can now dissolve in water
• If not, fishes in trouble!
• Process called polarization
• Generally, higher molar mass, greater
  polarizability of molecule
• Why?
• (larger the species, more likely e- held further
  away ? easier to polarize)

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Polarizability
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Induced dipole/induced dipole forces

• Non-polar entities can cause temporary dipoles
  between other non-polar entities
• ? causing intermolecular attractions
• Pentane, hexane, etc.
• The higher the molar mass, the greater the
  intermolecular attractions
• N-pentane has greater interactions than
  neo-pentane
• Latters smaller area for interactions
• I2 has a higher ?Hvap b.p. than other halogens
• cause nonpolar substances to condense to liquids
• and to freeze into solids
• (when the temperature is lowered sufficiently)
• Also called London Dispersion Forces

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Intermolecular Bonding Compared

• Strength
• Strongest Ion-dipole
• Strong Dipole-dipole (incl. H-bonding)
• Less strong dipole/induced-dipole
• Least strong induced-dipole/induced-dipole
  (London dispersion forces)
• Keep in mind ? a compound can have more than one
  of the above!

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Problem

• Rank the following in order of increasing boiling
  point and explain why
• NH3, CH4, and CO2

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Properties of Liquids

• (l) ? (g)
• Vaporization endothermic
• Condensation exothermic
• Boiling
• Why do we have bubbles?

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Vapor Pressure

• Leave a bottle of water open.
• Will evaporate
• Keep the lid on.
• can have equilibrium between liquid and gas
• Equilibrium vapor pressure/vapor pressure
• Measure of tendency of molecules to vaporize at
  given temp.

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What does this graph tell us?
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Volatility

• Ability of liquid to evaporate
• Higher the vapor pressure, greater the volatility
• Are polar cmpds or non-polar cmpds of equal
  molecular mass more volatile?

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Clausius-Clapeyron Equation

• Calculates ?Hvap
• What is this an equation for?
• What are the variables?
• C constant unique to cmpd
• R ideal gas constant
• 8.314472 J/mol?K

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Clausius-Clapeyron Equation

• Or, if given two pts.

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Clausius-Clapeyron Problem

• Methanol has a normal boiling point of 64.6C and
  a heat of vaporization of 35.2 kJ/mol. What is
  the vapor pressure of methanol at 12.0C?
• Does the answer make sense?
• Would water have a higher heat of vaporization?
• Why?
• Heat of vaporization of water 40.65 kJ/mol

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Boiling Point

• Bp ? temp. at which vapor pressure external
  (atmospheric pressure)
• At higher elevations atmospheric pressure is
  lower
• Thus, water boils at less than 100 C

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Critical Temperature and Pressure

• As temp. rises so does vapor pressure, but not
  infinitely
• At the critical point liq/gas interface
  disappears
• Critical temp/pressure
• Tc/Tp
• Gives supercritical fluid
• Density of a liq
• Viscosity of gas
• H2O
• Tc 374 C
• Tp 217.7 atm!
• Normal earth pressure ? 1 atm

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Supercritical fluid

• CO2 used in decaffeinating coffee
• Read about it on page 614

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Phase diagram

• Gives info on phase states of a substance at
  varying pressures and temperatures

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Deciphering a phase diagram

• Triple point
• Where all 3 states coexist
• Curves denote existence of two states
• Fusion (solid liq)
• Vaporization (liq gas)
• Sublimation (solid gas)
• Off the lines
• Single state

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Waters phase diagram

• Graph explains why water boils at lower temps at
  higher altitudes (next slide)
• If you apply increasing pressure (const. T of
  0C) to ice will it convert to water?
• Solid-liquid line has negative slope
• Its the opposite of most species
• Why?

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Sublimation

• Going from solid to gas without going through the
  liquid state
• Enthalpy of sublimation
• ?H?sublimation
• Iodine dry ice (solid CO2) sublimate
• Opposite of sublimation
• Deposition (g?s)
• Iodine demo

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CO2s Phase Diagram

• Explains sublimation
• How?
• Why is it called dry ice?

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Iodines Phase Diagram But does it really
sublimate?
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Problem

• The normal melting and boiling points of xenon
  are -112C and -107C, respectively.
• Its triple point is a -121C and 0.371 atm and
  its critical point is at 16.6C and 57.6 atm.
• a) Sketch the phase diagram for Xe, showing the
  axes, the four points given above, and indicating
  the area in which each phase is stable.
• b) If Xe gas is cooled under an external pressure
  of 0.131 atm, will it undergo condensation or
  deposition?