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Intermolecular Forces, Gases, and Liquids. Ch.13. 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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gases
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
intermolecular forces
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!
intermolecular forces1
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)
remember dipole moments
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
ion dipole ionization in aqueous medium water
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
solvation energy
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
permanent dipoles
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: -33°C vs. -162°C, respectively
hydrogen bonding
Hydrogen Bonding
  • A type of “super” dipole-dipole interaction
  • Interaction between e--rich atom connected to H entity & another H attached to e—rich atom
  • e--rich atom = O, F, N
  • Density water > than ice
    • Opposite of almost every other substance
  • Inordinately high heat capacity of water
  • High surface tension
    • Insects walk on water
  • Concave meniscus
hydrogen bonding1
Hydrogen Bonding
  • Boiling pts. of H2O, HF, and NH3much higher
surface tension
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
capillary action
Capillary Action
  • When water goes up a small glass tube
  • Due to polarity of Si-O bonding with water
  • Adhesive forces > 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
mercury
Mercury
  • Forms a convex meniscus
  • Doesn’t “climb” a glass tube
  • Due to cohesive forces > adhesive forces
viscosity
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, “You’re as slow as molasses in January”?
dipole induced dipole forces
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)
induced dipole induced dipole forces
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
    • Latter’s 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
intermolecular bonding compared
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!
problem
Problem
  • Rank the following in order of increasing boiling point and explain why:
  • NH3, CH4, and CO2
properties of liquids
Properties of Liquids
  • (l)  (g)
  • Vaporization = endothermic
  • Condensation = exothermic
  • Boiling
    • Why do we have bubbles?
vapor pressure
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.
volatility
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?
clausius clapeyron equation
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
clausius clapeyron equation1
Clausius-Clapeyron Equation
  • Or, if given two pts.:
clausius clapeyron problem
Clausius-Clapeyron Problem
  • Methanol has a normal boiling point of 64.6°C and a heat of vaporization of 35.2 kJ/mol. What is the vapor pressure of methanol at 12.0°C?
  • Does the answer make sense?
  • Would water have a higher heat of vaporization?
    • Why?
    • Heat of vaporization of water = 40.65 kJ/mol
boiling point
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
critical temperature and pressure
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
supercritical fluid
Supercritical fluid
  • CO2 used in decaffeinating coffee
  • Read about it on page 614
phase diagram
Phase diagram
  • Gives info on phase states of a substance at varying pressures and temperatures
deciphering a phase diagram
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
water s phase diagram
Water’s phase diagram
  • Graph explains why water boils at lower temps at higher altitudes (next slide)
  • If you apply increasing pressure (const. T of 0°C) to ice will it convert to water?
  • Solid-liquid line has negative slope
    • It’s the opposite of most species
      • Why?
sublimation
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
co 2 s phase diagram
CO2’s Phase Diagram
  • Explains sublimation
    • How?
  • Why is it called “dry ice”?
problem1
Problem
  • The normal melting and boiling points of xenon are -112°C and -107°C, respectively.
  • Its triple point is a -121°C and 0.371 atm and its critical point is at 16.6°C 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?
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