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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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  • 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


  • Forms a convex meniscus

  • Doesn’t “climb” a glass tube

  • Due to cohesive forces > adhesive forces


  • 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!


  • 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.


  • 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?


  • 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”?


  • 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?