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 2

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 (+δ/δ-) & 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





2) higher ion charge, stronger attraction





Li+ vs. Cs+ in water 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: -33°C vs. -162°C, respectively 9

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






Inordinately high heat capacity of water High surface tension





Opposite of almost every other substance

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 > 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 13

Mercury




Forms a convex meniscus Doesn’t “climb” a glass tube Due to cohesive forces > 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, “You’re as slow as molasses in January”? 15

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) 16

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









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 18

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! 19

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….





Keep the lid on….





Will evaporate can have equilibrium between liquid and gas

Equilibrium vapor pressure/vapor pressure


Measure of tendency of molecules to vaporize at given temp. 22

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 nonpolar cmpds of equal molecular mass more volatile? 24

Clausius-Clapeyron Equation






Calculates ∆Hvap What is this an equation for? What are the variables? C = constant unique to cmpd R = ideal gas constant


Ln Pvap = -

∆H vap R

1 × +C T

8.314472 J/mol•K

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Clausius-Clapeyron Equation


Or, if given two pts.:

∆H vap P2 1 1 ln( ) = − ×( − ) P1 R T2 T1

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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


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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





H2O:






Gives supercritical fluid
Density of a liq
Viscosity of gas Tc = 374 °C Tp = 217.7 atm!

Normal earth pressure ≈ 1 atm 29

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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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?

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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 34

CO2’s Phase Diagram


Explains sublimation





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









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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