Ch.13
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Kinetic-Molecular Theory says molecules/atoms separated
Little, if any, interactions
Not so in solids and liquids
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Examples:
Big difference in volume between liquids & solids and gases
Gases compressible, liqs & solids not
2

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Various electrostatic forces that attract molecules in solids/liqs
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Much weaker than ionic forces
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About 15% (or less) that of bond energies
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Remember, ionic bonds extremely powerful
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Boiling pt of NaCl = 1465 ° C!
3

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Reason behind importance of knowing about
IMF:
1) b.p. & m.p. and heats of vaporization (l  g) and fusion (s  l)
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2) solubility of gases, liquids, and solids
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3) determining structures of biochemicals
(DNA, proteins)
4

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Dipole moment = product of magnitude of partial charges (+  /  -) & their distance of separation
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= (1 Debye = 3.34 x 10 -30 C x m)
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Important in IMF
5

Ion-dipole: Ionization in aqueous medium (water)
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1) stronger attraction if ion/dipole closer
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Li + vs. Cs + in water
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2) higher ion charge, stronger attraction
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Be 2+ vs. Li + in water
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3) greater dipole, stronger attraction
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Dissolved salt has stronger attraction to water than methanol
6

7

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Or, enthalpy of hydration (if water) = energy of ionization in aq. media
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Water molecules surround both ions
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Example:
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Take hydration energies of G I metal ions
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Exothermicity decreases as you go down the column
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Cations become larger
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Easier to dissociate
8

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Positive end of one molecule attracted to negative end of other
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For ex: HCl
Dipole-dipole attractions
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Cmpds that exhibit greater d-d attractions have higher b.p., and H vap
Polar cmpds exhibit greater d-d attractions than non-polar cmpds
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NH
3 vs. CH
4
 equivalent molar masses (g/mol): 17 vs. 16, respectively
Boiling points: -33 ° C vs. -162 ° C, respectively
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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
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Opposite of almost every other substance
Inordinately high heat capacity of water
High surface tension
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Insects walk on water
Concave meniscus
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Boiling pts. of H
2
O, HF, and NH
3
much higher
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Outer molecules interact with surface, while inner interact with other molecules
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It has a “skin”
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Skin toughness = surface tension
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Energy required to break through surface
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Smaller surface area reason that water drops spherical
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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
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Forms a convex meniscus
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Doesn’t “climb” a glass tube
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Due to cohesive forces > adhesive forces
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Hydrogen-bonding increases viscosity
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But large non-polar liquids like oil have:
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1) large unwieldy molecules w/greater intermolecular forces
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2) greater ability to be entangled w/one another
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Did you ever hear the expression, “You’re as slow as molasses in January”?
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Polar entities induce dipole in nonpolar species like O
2
O
2
 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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Induced dipole/induced dipole forces
Non-polar entities can cause temporary dipoles between other non-polar entities
 causing intermolecular attractions
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Pentane, hexane, etc.
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The higher the molar mass, the greater the intermolecular attractions
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N-pentane has greater interactions than neo-pentane
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Latter’s smaller area for interactions
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I
2 has a higher  H vap
& 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
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Strength
Strongest: Ion-dipole
Strong: Dipole-dipole (incl. H-bonding)
Less strong: dipole/induced-dipole
Least strong: induced-dipole/induced-dipole (London dispersion forces)
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Keep in mind  a compound can have more than one of the above!
19

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Rank the following in order of increasing boiling point and explain why:
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NH
3
, CH
4
, and CO
2
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(l)  (g)
Vaporization = endothermic
Condensation = exothermic
Boiling
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Why do we have bubbles?
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Leave a bottle of water open….
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Will evaporate
Keep the lid on….
 can have equilibrium between liquid and gas
Equilibrium vapor
pressure/vapor pressure
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Measure of tendency of molecules to vaporize at given temp.
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Ability of liquid to evaporate
Higher the vapor pressure, greater the volatility
Are polar cmpds or nonpolar cmpds of equal molecular mass more volatile?
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Calculates ∆H vap
What is this an equation for?
What are the variables?
C = constant unique to cmpd
R = ideal gas constant
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8.314472 J/mol  K
Ln P vap

-

H vap
R

1
T

C
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Or, if given two pts.: ln(
P
2 )
P
1
 

H vap
R

1
(
T
2

1
T
1
)
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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?
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Does the answer make sense?
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Would water have a higher heat of vaporization?
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Why?
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Heat of vaporization of water = 40.65 kJ/mol
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Bp  temp. at which vapor pressure = external
(atmospheric pressure)
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At higher elevations atmospheric pressure is lower
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Thus, water boils at less than 100 ° C
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Critical Temperature and Pressure
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As temp. rises so does vapor pressure, but not infinitely
At the critical point liq/gas interface disappears
Critical temp/pressure
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T c
/T p
Gives supercritical fluid
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Density of a liq
Viscosity of gas
H
2
O:
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T c
= 374 ° C
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T p
= 217.7 atm!
Normal earth pressure  1 atm
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CO
2 used in decaffeinating coffee
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Read about it on page
614
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Gives info on phase states of a substance at varying pressures and temperatures
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Triple point
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Where all 3 states coexist
Curves denote existence of two states
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Fusion (solid & liq)
Vaporization (liq & gas)
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Sublimation (solid & gas)
Off the lines
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Single state
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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
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It’s the opposite of most species
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Why?
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Going from solid to gas without going through the liquid state
Enthalpy of sublimation
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 H  sublimation
Iodine & dry ice (solid
CO
2
) sublimate
Opposite of sublimation
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Deposition (g  s)
Iodine demo
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2
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Explains sublimation
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How?
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Why is it called
“dry ice”?
35

Iodine’s Phase Diagram: But does it really sublimate?
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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.
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 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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