Transcript Intermolecular Forces, Solids, and Liquids

Intermolecular Forces,
Gases, and Liquids
Ch.13
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Gases
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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
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Intermolecular Forces
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Various electrostatic forces that attract
molecules in solids/liqs
Much weaker than ionic forces
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!
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Intermolecular Forces
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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)
2) solubility of gases, liquids, and solids
3) determining structures of biochemicals
(DNA, proteins)
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Remember dipole moments?
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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)
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1) stronger attraction if ion/dipole closer
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2) higher ion charge, stronger attraction
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Li+ vs. Cs+ in water
Be2+ vs. Li+ in water
3) greater dipole, stronger attraction
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Dissolved salt has stronger attraction to water than
methanol
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Solvation energy
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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
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Easier to dissociate
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Permanent dipoles
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Positive end of one molecule attracted to negative end
of other
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For ex: HCl
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Dipole-dipole attractions
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Cmpds that exhibit greater d-d attractions have higher
b.p., and Hvap
Polar cmpds exhibit greater d-d attractions than non-polar cmpds
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NH3 vs. CH4
 equivalent molar masses (g/mol): 17 vs. 16, respectively
Boiling points: -33°C vs. -162°C, respectively
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Hydrogen Bonding
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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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Inordinately high heat capacity of water
High surface tension
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Opposite of almost every other substance
Insects walk on water
Concave meniscus
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Hydrogen Bonding
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Boiling pts. of H2O, HF, and NH3 much higher
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Surface Tension
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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
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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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Mercury
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Forms a convex meniscus
Doesn’t “climb” a glass tube
Due to cohesive forces > adhesive forces
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Viscosity
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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”?
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Dipole/Induced Dipole Forces
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Polar entities induce dipole in nonpolar species like O2
O2 can now dissolve in water
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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
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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
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
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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)
Keep in mind  a compound can have more than one
of the above!
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Problem
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Rank the following in order of increasing
boiling point and explain why:
NH3, CH4, and CO2
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Properties of Liquids
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(l)  (g)
Vaporization =
endothermic
Condensation =
exothermic
Boiling
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Why do we have
bubbles?
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Vapor Pressure
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Leave a bottle of water
open….
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Keep the lid on….
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Will evaporate
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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What does this graph tell us?
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Volatility
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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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Clausius-Clapeyron Equation
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Calculates ∆Hvap
What is this an equation
for?
What are the variables?
C = constant unique to
cmpd
R = ideal gas constant
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Ln Pvap  -
H vap
R
1
 C
T
8.314472 J/molK
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Clausius-Clapeyron Equation
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Or, if given two pts.:
H vap
P2
1 1
ln( )  
(  )
P1
R
T2 T1
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Clausius-Clapeyron Problem
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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?
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
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Bp  temp. at which vapor pressure = external
(atmospheric pressure)
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
 Tc/Tp
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H2O:
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Gives supercritical fluid
 Density of a liq
 Viscosity of gas
Tc = 374 °C
Tp = 217.7 atm!
Normal earth pressure  1 atm
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Supercritical fluid
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CO2 used in
decaffeinating coffee
Read about it on page
614
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Phase diagram
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Gives info on phase states of a substance at
varying pressures and temperatures
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Deciphering a phase diagram
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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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Water’s phase diagram
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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
 It’s the opposite of most
species
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Why?
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Sublimation
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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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CO2’s Phase Diagram
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Explains
sublimation
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How?
Why is it called
“dry ice”?
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Iodine’s Phase Diagram: But does it
really sublimate?
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Problem
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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.
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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