Liquids, Solids, and Intermolecular Forces
(Ch 15)
Suggested HW: (Ch 15) 13-17
States of Matter Differ By Intermolecular Distance
• The state of a substance at a given temperature and pressure is
determined by two factors:
• Thermal energy of the molecules
• Intermolecular forces (called Van der walls forces) between molecules
Intermolecular Forces
• Intermolecular forces are the attractive forces that exist between
molecules. These are not the same as intramolecular forces (i.e. bonds)
which exist within a molecule.
• Intermolecular forces are much weaker than intramolecular forces.
– Consider water
– Water would have to be heated to thousands of degrees to break the
O-H covalent bonds.
– However, the intermolecular forces can be overcome simply by boiling
the water (100oC) which separates the molecules.
• Nonetheless, intermolecular forces are very important and dictate many
of the important physical characteristics of molecules
Intermolecular Forces: Coulombic Attractions
• As you recall, ionic compounds are solids
at room temperature. There are ion-ion
attractions in ionic compounds.
• The coulombic force that holds ions
together is very strong. Coulombic
attractions are the strongest of all
intermolecular forces.
• Therefore, all ionic compounds have very
high melting/boiling points, and are solid
at room temperature.
Dipole-Dipole Forces
• Polar molecules attract one another. This type of intermolecular force is
called dipole-dipole attraction.
+
δ
δ
Covalent bond:
Very Strong
-
+
δ
δ
-
Dipole-dipole interaction:
• Polar molecules will orient themselves in a way to maximize these
attractions. The strength of these attractions increases with increasing
polarity. Polar molecules have higher melting/boiling points than non
polar ones.
Intermolecular Forces: Dipole-Dipole Forces
• The magnitude of the melting/boiling temperatures of various substances
reflect how strongly the molecules attract one another.
• The more strongly the molecules attract, the harder it is to separate them.
Hence, the higher the melting/boiling temperatures.
• Recall polarity from chapter 8. Any molecule with a net dipole is polar.
+
δ
Partial positive
character
H
Cl
δ
Partial negative
character
Ion-Dipole Interactions
• When salts are dissolved
in water, the water dipoles
are oriented around the
ions.
• This strong interaction
explains why adding salt
to water raises its boiling
temperature, or why
salting snow causes it to
melt.
Solution
Boiling Temp. (oC)
Boiling Temp. (oF)
Pure Water
100.0
212.0
Satd. NaCl (aq)
108.7
227.7
Hydrogen Bonding
• There is a special, unusual, very strong type of dipole-dipole interaction
known as hydrogen bonding
• Hydrogen bonding is a dipole interaction that exists only between the
H atom in an H—F, H—O, or H—N bond on one molecule and an
adjacent lone pair on another F, O, or N atom in another molecule
• Because hydrogen atoms are so small, the large partial positive charge
induced on H when bonding with these highly electronegative elements
is highly concentrated. Therefore, it strongly attracts neighboring N, O,
and F atoms, and can approach them very closely.
hydrogen bonds
δ+
δ-
δ-
O
O
H
H
δ+
δ+
H
H
δ+
δ-
O
δ+
H
H
δ+
δ-
δ-
O
δ+
H
O
H
δ+
δ+
H
H
δ+
Hydrogen Bonding Causes Abnormalities in Boiling Point
Trend
Actual Boiling Temp
of Water (100oC)
Predicted Boiling Temp
of Water (-90oC)
Without hydrogen bonding,
all water would be gas
(steam) at normal room
temperature!!!
Hydrogen Bonding
• For example, consider ethanol and dimethyl ether (C2H6O, MW 46 g/mol)
• The –OH bond in ethanol is susceptible to
hydrogen bonding, whereas the C-O-C bond in
dimethyl ether is not.
• The strong attraction between these –OH
groups results in the massive difference in
boiling point between these substances.
ethanol
BP: 78.3oC
dimethyl ether
BP: -22.0oC
Structure and Density of Ice
• Water is one of the few compounds
that is less dense in its solid phase
than its liquid phase.
• This is due to hydrogen bonding.
• In liquid water, 80% of the atoms are
H-bonded. In ice, 100% are Hbonded.
• To maximize H-bonding, the water
molecules in ice spread out.
• Therefore, we have the same mass of
water, with a larger volume. Since
ρ=(mass/volume), ρ decreases.
London Dispersion Forces
• With nonpolar molecules, there are no dipoles, so we would not expect to
see dipole-dipole interactions. Despite this, intermolecular interactions
have still been observed.
• For example, nonpolar gases like Helium can be liquified, but how can this
happen? What force brings the He atoms together?
• Fritz London, a physicist, proposed that the motion of electrons in a
nonpolar molecule can create instantaneous dipoles
London Dispersion Forces
• Lets take a Helium atom. At some moment in time, the electrons are spread
out within the atom
e-
2+
e-
• However, because electrons are constantly moving, electrons can
temporarily end up on the same side of the atom, creating an instantaneous
dipole.
δ
ee-
2+
δ
+
London Dispersion Forces
• These dipoles induce other dipoles on neighboring atoms, the cycle
repeats indefinitely. This leads to the formation of a condensed phase.
e-
δ
repel
2+
e-
ee-
2+
-
δ
+
δ δ
ee-
2+
ee-
2+
+
δ
attract
+
δ δ-
e-
ee-
2+
2+
e-
+
δ
London Dispersions
• The ease of the electron distortion is called
polarizability. Bigger molecules are more
polarizable, so they are more prone to
instantaneous dipoles.
• Hence, London dispersion forces increase with
increasing molar mass because heavier
atoms/molecules have more electrons, and are
therefore are more polarizable.
All substances have dispersion forces.
C5H12
C12H26 C18H38
Boiling Points Increase With Increasing Strength of London
Dispersion Forces
Other Things to Consider
• Show to the right are two isomers of
pentane, n-pentane and
neopentane. Both have the formula
C5H12 (same mass), but the boiling
points are very different. Why?
• The long, straight shape of npentane allows for more atoms to
participate in LD interaction, as
compare to the compact
neopentane molecule.
n-pentane
BP: 36.1o C
neopentane
BP: 9o C
Like Substances Mix
• Polar substances are soluble in polar substances.
• Nonpolar substances are soluble in nonpolar substances.
• Polar and nonpolar substances DO NOT MIX (ex. oil and water)
• E.G. Oil and water; Dipole-dipole and H-bond
forces between water molecules are MUCH
stronger than the London dispersion forces
between oil and water
• Water molecules would rather associate with
other water molecules than to associate with
oil.
• Salts are soluble in water and other highly
polar solvents, but insoluble in nonpolar and
weakly polar solvents.
Molecular Geometry of Fats And Oils
• A typical nutritional label breaks down the types of
fats present in a food to include saturated fats and
unsaturated fats
• A saturated fat is a long chain fatty acid
containing no C=C double bonds
• Unsaturated fats contain one (mono) or more
(poly) double bonds
• Chemically, fats and oils differ on this basis,
although these terms are used interchangeably in
nutrition. Fats are saturated fatty acids and are
solid at room temperature. Unsaturated fats are
techically oils, as they are liquids.
• Oils can be converted into fats by breaking the C=C
double bonds by a process called hydrogenation (ex.
Margarine)
Molecular Geometry of Fats And Oils
• Unsaturated fats can be separated into two classes, trans and cis. These differ
by the orientation of terminal atoms about the C=C double bond. Trans fat are
a byproduct of hydrogenation.
trans
cis
• The structure of trans oils allows the molecules to stack, whereas the cis
structure does not. This stacking results in enhancement of the London
dispersion force, which leads to the formation of a solid fat. This causes
clogging of arteries.
trans
stacking
cis
stacking
Molecular Geometry of Fats And Oils
• Saturated fats are considered to be worse for heart health than unsaturated
fats because of the lack of “kinks” in the carbon chain, which makes it easier for
the molecules to align and stack.
n
clogged artery
Which of the following substances mix?
•
•
•
•
•
O2 and CH3CH2CH2CH3 (butane)
SO2 and CCl4
H2O and CH3OH
K2S and H2O
H2O and CH3CH2NH2
Intermolecular Forces At Work: Surface Tension
• Imagine a body of liquid. The liquid molecules
in the bulk have attractive intermolecular forces
on all sides, pulling in all directions. These
interactions minimize the energy of these
molecules.
• Molecules at the surface have less interactions,
which makes them less stable due to their
higher potential energy. The net force on the
surface molecules is inward, and they cluster
close together to minimize surface area.
Intermolecular Forces At Work: Surface Tension
• The surface tension of a liquid is the energy required to increase its
surface area by a unit amount (J/cm2 or mJ/m2). The surface tension of a
liquid is what causes its surface to create a “skin” that resists penetration,
or assume a droplet shape.
Intermolecular Forces At Work: Viscosity
• Viscosity is a measure of a liquid’s resistance to flow and is a temperature
dependent property (e.g. Maple syrup is more viscous than water).
Stronger intermolecular forces lead to higher viscosity.
• C.I.R.L.: Viscosity is an important property of motor
oil. It must be thick enough to lubricate and protect
the engine, but thin enough not to clog the
components.
• Multi-grade oils allow drivers to keep the same
oil all year long without a loss of performance.
• Ex. An oil with a 10W-30 rating has a viscosity of
10 (thin) in the winter and 30 (thick) in the
summer. How?
Intermolecular Forces At Work: Capillary Action
• Closely related to surface tension,
capillary action is the ability of a liquid to
flow against gravity up a tube.
• This is caused by a combination of
intermolecular forces between
neighboring liquid molecules (cohesive
forces), as well as intermolecular forces
between the liquid molecules and the
molecules of the surface of the tube
(adhesive forces).
• If the adhesive forces are stronger than
the cohesive forces, the liquid rises.
Intermolecular Forces At Work: Capillary Action
• A consequence of surface tension and
capillary action is the formation of a
meniscus (the shape formed by a liquid in a
tube).
• In systems with strong adhesion forces, the
liquid molecules rise up the sides of the
tube, forming a concave meniscus.
• In systems with strong cohesion forces, the
liquid molecules crowd around the interior
to maximize their interactions, creating a
convex meniscus.
Summary
• We can arrange the intermolecular forces by relative strength:
1. Coulombic attraction
2. Ion-dipole/Hydrogen Bonding
3. Dipole-Dipole
4. *London Dispersion
• It is important to note that molecules with large nonpolar portions
and polar end groups will not mix with polar solvents
(ex. CH3CH2CH2CH2CH2CH2CH2OH) because the LD forces dominate.
• The list above is a general trend. For very large molecules, LD forces
can exceed the strength of dipole-based forces.
Example
• Arrange the following in order of increasing boiling point:
– H2O, Xe, CH4, NaCl (s), CH3OH, SO2, SO3
• To do this, we must consider the intermolecular forces that exist between
the molecules. Those with the strongest forces between them will be the
hardest to separate, and thus, have the highest boiling temperature.
– DRAW THE LEWIS STRUCTURES (where applicable)
– Separate polar and nonpolar molecules, then sort by strength of
attraction
CH4 < SO3 < Xe < SO2 < CH3OH < H2O < NaCl (s)
LD
DD
HB
CA