Have studied INTRAmolecular forces—
the forces holding atoms together to
form molecules.
Now turn to forces between molecules —
INTERmolecular forces.
Forces between molecules, between ions,
or between molecules and ions.
© 2006 Brooks/Cole - Thomson
Intermolecular
Forces
Ion-Ion Forces
for comparison of magnitude
Na+—Cl- in salt
These are the strongest
forces.
Lead to solids with high
melting temperatures.
NaCl, mp = 800 oC
MgO, mp = 2800 oC
© 2006 Brooks/Cole - Thomson
Covalent Bonding Forces
for comparison of magnitude
C=C, 610 kJ/mol
C–C, 346 kJ/mol
C–H, 413 kJ/mol
© 2006 Brooks/Cole - Thomson
CN, 887 kJ/mol
Attraction Between Ions
and Permanent Dipoles
••
••
water
-
dipole
O
H
H +
Water is highly polar and can
interact with positive ions
to give hydrated ions in
water.
© 2006 Brooks/Cole - Thomson
Attraction Between Ions
and Permanent Dipoles
••
••
water
-
dipole
O
H
H +
Water is highly polar and can
interact with positive ions
to give hydrated ions in
water.
© 2006 Brooks/Cole - Thomson
Attraction Between Ions
and Permanent Dipoles
Many metal ions are hydrated. This is the reason
metal salts dissolve in water.
© 2006 Brooks/Cole - Thomson
Attraction Between Ions
and Permanent Dipoles
Attraction between ions and dipole depends on ion charge
and ion-dipole distance.
Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+
- H
O
H
+
•••
Mg2+
-1922 kJ/mol
© 2006 Brooks/Cole - Thomson
- H
O
H
+
- H
O
H
+
•••
Na +
-405 kJ/mol
•••
Cs+
-263 kJ/mol
Dipole-Dipole Forces
Such forces bind molecules having permanent dipoles
to one another.
© 2006 Brooks/Cole - Thomson
Dipole-Dipole Forces
Influence of dipole-dipole forces is seen in the boiling
points of simple molecules.
Compd
Mol. Wt.
Boil Point
N2
28
-196 oC
CO
28
-192 oC
Br2
160
59 oC
ICl
162
97 oC
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding
A special form of dipole-dipole attraction, which
enhances dipole-dipole attractions.
H-bonding is strongest when X and Y are N, O, or F
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding in H2O
H-bonding is especially strong
in water because
• the O—H bond is very polar
• there are 2 lone pairs on the
O atom
Accounts for many of water’s
unique properties.
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding in H2O
Ice has open
lattice-like
structure.
Ice density is
< liquid.
And so solid floats
on water.
Snow flake: www.snowcrystals.com
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding in H2O
Ice has open lattice-like structure.
Ice density is < liquid and so solid floats on water.
One of the VERY few
substances where solid is
LESS DENSE than the
liquid.
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding in H2O
H bonds ---> abnormally high specific heat capacity of water (4.184 J/g•K)
This is the reason water is used to put out fires, it is the reason lakes/oceans
control climate, and is the reason thunderstorms release huge energy.
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding
H bonds leads to abnormally high boiling point of water.
© 2006 Brooks/Cole - Thomson
Boiling Points of Simple
Hydrogen-Containing
Compounds
Active Figure 13.8
© 2006 Brooks/Cole - Thomson
Hydrogen Bonding in Biology
H-bonding is especially strong in biological systems —
such as DNA.
DNA — helical chains of phosphate groups and sugar
molecules. Chains are helical because of tetrahedral
geometry of P, C, and O.
Chains bind to one another by specific hydrogen bonding
between pairs of Lewis bases.
—adenine with thymine
—guanine with cytosine
© 2006 Brooks/Cole - Thomson
Double helix
of DNA
Portion of a
DNA chain
© 2006 Brooks/Cole - Thomson
Base-Pairing through H-Bonds
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Base-Pairing through H-Bonds
© 2006 Brooks/Cole - Thomson
FORCES INVOLVING
INDUCED DIPOLES
How can non-polar molecules such as O2 and I2 dissolve in
water?
The water dipole INDUCES a dipole in
the O2 electric cloud.
Dipole-induced
dipole
© 2006 Brooks/Cole - Thomson
FORCES INVOLVING
INDUCED DIPOLES
Solubility increases with mass the gas
© 2006 Brooks/Cole - Thomson
FORCES INVOLVING
INDUCED DIPOLES
• Process of inducing a
dipole is polarization
• Degree to which electron
cloud of an atom or
molecule can be
distorted in its
polarizability.
© 2006 Brooks/Cole - Thomson
IM FORCES — INDUCED DIPOLES
Consider I2
dissolving
in ethanol,
CH3CH2OH.
-
I-I
- O
R
H
+
© 2006 Brooks/Cole - Thomson
I-I
The alcohol
temporarily
creates or
INDUCES a
dipole in I2.
+
- O
R
H
+
FORCES INVOLVING
INDUCED DIPOLES
Formation of a dipole in two nonpolar I2 molecules.
Induced dipole-induced dipole
The induced forces between I2 molecules are very weak, so solid I2 sublimes
(goes from a solid to gaseous molecules).
© 2006 Brooks/Cole - Thomson
FORCES INVOLVING
INDUCED DIPOLES
The magnitude of the induced dipole depends on the
tendency to be distorted.
Higher molec. weight ---> larger induced dipoles.
Molecule
Boiling Point (oC)
CH4 (methane)
- 161.5
C2H6 (ethane)
- 88.6
C3H8 (propane) - 42.1
C4H10 (butane)
- 0.5
© 2006 Brooks/Cole - Thomson
Boiling Points of Hydrocarbons
C4H10
C3H8
C2H6
CH4
Note linear relation between bp and molar mass.
© 2006 Brooks/Cole - Thomson
Intermolecular Forces Summary
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Measuring Equilibrium
Vapor Pressure
Liquid in flask evaporates and exerts pressure on manometer.
Active Fig. 13.17
© 2006 Brooks/Cole - Thomson
Equilibrium Vapor Pressure
Active Figure 13.18
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Liquids Equilibrium Vapor Pressure
FIGURE 13.18: VP as a function of T.
1. The curves show all conditions of P and T where LIQ and VAP are in
EQUILIBRIUM
2. The VP rises with T.
3. When VP = external P, the liquid boils.
This means that BP’s of liquids change with altitude.
4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING
POINT
5. VP of a given molecule at a given T depends on IM forces. Here the VP’s
are in the order
ether
O
C 2H 5
H 5C 2
dipoledipole
alcohol
O
H 5C 2
H
H-bonds
increasing strength of IM interactions
© 2006 Brooks/Cole - Thomson
water
O
H
H
extensive
H-bonds
Liquids
HEAT OF VAPORIZATION is the heat req’d (at
constant P) to vaporize the liquid.
LIQ + heat ---> VAP
Compd.
∆Hvap (kJ/mol)
H2O
40.7 (100 oC) H-bonds
SO2
26.8 (-47 oC) dipole
Xe
12.6 (-107 oC) induced dipole
© 2006 Brooks/Cole - Thomson
IM Force
Phase Diagrams
© 2006 Brooks/Cole - Thomson
TRANSITIONS
BETWEEN PHASES
Section 13.10
Lines connect all conditions of T and P where EQUILIBRIUM exists
between the phases on either side of the line.
(At equilibrium particles move from liquid to gas as fast as they move from
gas to liquid, for example.)
© 2006 Brooks/Cole - Thomson
Phase Diagram for Water
Liquid phase
Solid phase
Gas phase
© 2006 Brooks/Cole - Thomson
Phase Equilibria — Water
Solid-liquid
Gas-Liquid
Gas-Solid
© 2006 Brooks/Cole - Thomson
Triple Point —
Water
At the TRIPLE POINT all
three phases are in equilibrium.
© 2006 Brooks/Cole - Thomson
Phases Diagrams—
Important Points for Water
T(˚C)
P(mmHg)
Normal boil point
100
760
Normal freeze point
0
760
Triple point
0.0098
© 2006 Brooks/Cole - Thomson
4.58
Critical T and P
As P and T increase, you finally reach the CRITICAL T and P
Above critical T
no liquid exists
no matter how
high the pressure.
At P < 4.58 mmHg and T < 0.0098 ˚C solid H2O can go directly to
vapor. This process is called SUBLIMATION
This is how a frost-free refrigerator works.
© 2006 Brooks/Cole - Thomson
Critical T and P
COMPD
Tc(oC)
H2O
374
218
CO2
31
73
CH4
-82
46
Freon-12
(CCl2F2)
112
41
Pc(atm)
Notice that Tc and Pc depend on intermolecular
forces.
© 2006 Brooks/Cole - Thomson
Solid-Liquid Equilibria
In any system, if you increase P the DENSITY will go up.
Therefore — as P goes up, equilibrium favors phase with the larger density
(or SMALLER volume/gram).
Liquid H2O Solid H2O
1 g/cm3
Density
cm3/gram
© 2006 Brooks/Cole - Thomson
1
0.917 g/cm3
1.09
Solid-Liquid Equilibria
P
Solid
H 2O
Liquid
H 2O
Raising the pressure at constant T
causes water to melt.
Normal
freezing
point
760
mmHg
0 °C
The NEGATIVE SLOPE of the S/L
line is unique to H2O. Almost
everything else has positive slope.
T
The behavior of water under pressure is an example of
LE CHATELIER’S PRINCIPLE
At Solid/Liquid equilibrium, raising P squeezes the solid.
It responds by going to phase with greater density, i.e., the liquid phase.
© 2006 Brooks/Cole - Thomson
CO2 Phase Diagram
© 2006 Brooks/Cole - Thomson
CO2 Phases
Separate
phases
Figure 13.41
© 2006 Brooks/Cole - Thomson
Increasing
pressure
More
pressure
Supercritcal
CO2
Dissolving An Ionic Solid
Active Figure 14.9
© 2006 Brooks/Cole - Thomson
Energetics
of the Solution Process
Figure 14.8
© 2006 Brooks/Cole - Thomson
Energetics
of the Solution Process
If the enthalpy of formation
of the solution is more
negative that that of the
solvent and solute, the
enthalpy of solution is
negative.
The solution process is
exothermic!
© 2006 Brooks/Cole - Thomson
Colligative Properties
On adding a solute to a solvent, the props. of the solvent are
modified.
• Vapor pressure
decreases
• Melting point
decreases
• Boiling point
increases
• Osmosis is possible (osmotic pressure)
These changes are called COLLIGATIVE PROPERTIES.
They depend only on the NUMBER of solute particles
relative to solvent particles, not on the KIND of solute
particles.
© 2006 Brooks/Cole - Thomson
Concentration Units
MOLE FRACTION, X
For a mixture of A, B, and C
mol A
X A  mol fraction A =
mol A + mol B + mol C
MOLALITY, m
mol solute
m of solute =
kilograms solvent
WEIGHT % = grams solute per 100 g solution
© 2006 Brooks/Cole - Thomson
Calculating Concentrations
Dissolve 62.1 g (1.00 mol) of ethylene glycol in
250. g of H2O. Calculate mol fraction, molality,
and weight % of glycol.
© 2006 Brooks/Cole - Thomson
Understanding
Colligative Properties
To understand colligative properties, study the LIQUIDVAPOR EQUILIBRIUM for a solution.
© 2006 Brooks/Cole - Thomson
Understanding
Colligative Properties
VP of H2O over a solution depends on the number of H2O
molecules per solute molecule.
Psolvent proportional to Xsolvent
Psolvent = Xsolvent • Posolvent
VP of solvent over solution
= (Mol frac solvent)•(VP pure solvent)
RAOULT’S LAW
© 2006 Brooks/Cole - Thomson
Raoult’s Law
An ideal solution is one that obeys Raoult’s law.
PA = XA • PoA
Because mole fraction of solvent, XA, is always less than 1, then PA is always
less than PoA.
The vapor pressure of solvent over a solution is always LOWERED!
© 2006 Brooks/Cole - Thomson
Raoult’s Law
Assume the solution containing 62.1 g of glycol in 250. g of
water is ideal. What is the vapor pressure of water over the
solution at 30 oC? (The VP of pure H2O is 31.8 mm Hg; see App. E.)
© 2006 Brooks/Cole - Thomson
Changes in Freezing and Boiling Points of Solvent
VP Pure solvent
1 atm
VP solvent
after adding
solute
P
BP solution
BP pure
solvent
T
• For a 2-component system where A is the solvent and B is the solute
• ∆PA = VP lowering = XBPoA
• VP lowering is proportional to mol frac solute!
• For very dilute solutions, ∆PA = K•molalityB where K is a
proportionality constant.
• This helps explain changes in melting and boiling points.
© 2006 Brooks/Cole - Thomson
Vapor
Pressure
Lowering
Figure 14.14
© 2006 Brooks/Cole - Thomson
The boiling point of a solution
is higher than that of the pure
solvent.
© 2006 Brooks/Cole - Thomson
Elevation of Boiling Point
Elevation in BP = ∆TBP = KBP•m
(where KBP is characteristic of solvent)
VP Pure solvent
1 atm
VP solvent
after adding
solute
P
BP solution
BP pure
solvent
© 2006 Brooks/Cole - Thomson
T
Change in Boiling Point
Dissolve 62.1 g of glycol (1.00 mol) in 250. g of water. What
is the BP of the solution?
KBP = +0.512 oC/molal for water (see Table 14.3).
© 2006 Brooks/Cole - Thomson
Change in Freezing Point
Pure water
Ethylene glycol/water
solution
The freezing point of a solution is LOWER than that
of the pure solvent.
FP depression = ∆TFP = KFP•m
© 2006 Brooks/Cole - Thomson
Lowering the Freezing Point
Water with and without antifreeze
© 2006 Brooks/Cole - Thomson
When a solution freezes, the solid
phase is pure water. The solution
becomes more concentrated.
Freezing Point Depression
Calculate the FP of a 4.00 molal glycol/water solution.
KFP = -1.86 oC/molal (Table 14.4)
© 2006 Brooks/Cole - Thomson
Freezing Point Depression
How much NaCl must be dissolved in 4.00 kg of water to
lower FP to -10.00 oC?.
© 2006 Brooks/Cole - Thomson
Freezing Point Depression
How much NaCl must be dissolved in 4.00 kg of water to
lower FP to -10.00 oC?.
© 2006 Brooks/Cole - Thomson
Boiling Point Elevation and
Freezing Point Depression
∆T = K•m•i
A generally useful equation
i = van’t Hoff factor = number of particles produced per
formula unit.
Compound
Theoretical Value of i
glycol
1
NaCl
2
CaCl2
3
© 2006 Brooks/Cole - Thomson
Osmosis
The semipermeable membrane allows only the movement of
solvent molecules.
Solvent molecules move from pure solvent to
solution in an attempt to make both have the
same concentration of solute.
Driving force is entropy
© 2006 Brooks/Cole - Thomson
Osmosic Pressure, ∏
Osmotic
pressure
© 2006 Brooks/Cole - Thomson
Equilibrium is reached when
pressure — the OSMOTIC
PRESSURE, ∏ — produced by
extra solution counterbalances
pressure of solvent molecules
moving thru the membrane.
∏ = cRT
(c is conc. in mol/L)
Osmosis and Living Cells
© 2006 Brooks/Cole - Thomson
Reverse Osmosis
Water Desalination
Water desalination plant in Tampa
© 2006 Brooks/Cole - Thomson
Osmosis
Calculating a Molar Mass
Dissolve 35.0 g of hemoglobin in enough water to make 1.00
L of solution. ∏ measured to be 10.0 mm Hg at 25 ˚C.
Calculate molar mass of hemoglobin.
© 2006 Brooks/Cole - Thomson