Chapter12: Dispersion Forces and Hydrogen Bonding
London dispersion forces occur between atoms
or molecules of nonpolar substances. Monoatomic atoms (noble gases), diatomic
molecules (H2, N2, O2, F2, Cl2)
and nonpolar compounds (CH4, CCl4, BF3, BeH2,
etc.) are all characterized by a symmetric sharing of electrons in the atom
or molecule. These compounds do not have permanent dipoles as occur in heteronuclear
diatomic molecules (HCl, HBr, HI, etc) and polar compounds (SO2,
H2S, NCl3, etc).
In class we saw an animation of a collection of monoatomic neutral atoms.
If we take "snapshots" of the electron distribution we would generally see a
symmetric distribution of the electron density. However, occasionally we see,
but not very often, an unequal sharing of the electrons. A "snapshot" an instant
later would reveal a return to an equal distribution of the electrons. Every
once and a while we observe that the electrons are unequally distributed around
the nuclei, when this occurs there is a very small charge separation created
which gives rise to an instantaneous dipole.
Another atom near this instantaneous dipole will also be effected causing
a shift of its electron distribution resulting in a small dipole around it.
When this occurs, even for an instant there is a small attraction between the
two molecules. The strength of the London dispersion forces depends on how easily
the electron cloud is distorted or polarized.
The larger the molecule the further the electrons are from the nucleus and
the easier the electron cloud can be polarized. So the magnitude of the dispersion
forces increases with increasing molecular size. Note: dispersion forces operate
in all molecules whether they are polar or nonpolar.
It has been shown (not in class) that the force of attraction between two
nonpolar molecules is inversely proportional to the seventh power of the distance
and directly proportional to a property of each molecule called polarizability.
Polarizability of an atom or a molecule is a measure of the ease with which
the electrons and nuclei can be displaced from their average positions. When
the electrons occupy a large volume of space, which occurs in an atom or molecule
with many electrons, the polarizability of the substance is large. The units
on polarizability are the units of volume, m3.
When the polarizability is large for a particular atom or molecule the magnitude
of the instantaneous dipole can be large with the result producing a stronger
attraction between particles.
The electrons which are the most easily displaced in an atom or molecule
are the valence electrons, these are the furthest from the nucleus. So valence
electrons make the greatest contribution to the polarizability. The force acting
on the valence electrons depends on their distance from the nucleus and on the
core charge. For any group in the periodic table the core charge remains constant,
so we expect polarizability to increase as the atomic size increases. So the
polarizability of HI is greater than the polarizability of HF.
In molecules with large numbers of atoms the polarizability will be larger
compared to smaller molecules. The polarizability of a molecule increases with
both increasing size and increasing numbers of atoms in the molecule. So we
expect the magnitude of the instantaneous dipoles, and therefore the strength
of the London forces, to be greater the greater the number of atoms in a molecule
and the larger the atoms.
The polarizability of N2 is greater than H2, and that of CCl4 greater than
CH4, and CO2 greater than that of CO. The strength of the intermolecular attractive
forces is reflected in the boiling points of the substances.
Two contributions to the intermolecular attractions between covalent molecules;
1) dipole-dipole forces (present only when the molecule is polar) and, 2) London
forces (present between all molecules and are particularly important for large
molecules.) London forces are often stronger than the dipole-dipole forces between
polar molecules. It is rare that dipole-dipole forces will dominate the properties
of a molecule unless the dipole-dipole forces are particularly strong. This
occurs for hydrogen-bonded systems.
Most solids expand up to 10 % of their volume when they melt. Water expands
by the same amount when it freezes!
Most solids are more dense than the liquid phase. However, ice has a density
of 0.917 g/cm3 while liquid water has a density of 0.998 g/cm3.
Water has a melting point which is 100 degrees C higher than expected
for its group of hydides.
Water has a boiling point 200 degrees C higher than expected for its group
Water has the highest surface tension of any liquid except mercury.
Water is an excellent solvent, dissolving many ionic compounds which are
insoluble in other compounds.
Water has a high heat capacity.
All of these properties can be understood in terms of the intermolecular attractive
forces which exist between water molecules. The particular intermolecular attractive
force is called hydrogen-bonding. Hydrogen bonding is another intermolecular
force, which is stronger than London and dipole-dipole forces. Hydrogen bonding
forces occurs in a particularly special group of polar compounds. These compounds
are characterized by the X-H bond, where X can be O, N, or F. Examples of compounds
that exhibit hydrogen bonding forces are H2O, NH3 and
HF. Oxygen, nitrogen and fluorine are small strongly electronegative atoms.
In a covalent bond with hydrogen these atoms attract the pair of electrons giving
rise to a partial positive charge on the hydrogen atom. This partial positive
charge on the hydrogen atom is very interested in any negative charge in another
adjacent molecule that comes close to it. So when another polar molecule which
contains an atom such as O, N, F, Cl or S come near an X-H bond a hydrogen bond
So what would a hydrogen-bonding interaction look like when two or three molecules
of water were close to each other? To do this lets begin with just a lewis structure
for water. Draw the Lewis structure for water. When you are finished look at
Bulk Properties of Liquids
As an introduction to the need to define and discuss intermolecular attractive
forces we began our discussion of the bulk properties of gases, liquids and
solids. Some of these are discussed in Section 12.4 on pages 440 - 447 in Silberberg.
We talked about the following bulk properties;
* molar volume
* thermal expansion
* surface tension
Molar volumes of gases are much larger compared to the molar volume of liquids
and solids. For gases, assuming ideal behavior, the molar volume is 22,400 mLs
at 0 degrees C and 1 atmosphere. For liquids and solids the molar volume ranges
between 10 and 100 mLs at room temperature and pressure. Molar volume is inversely
proportional to the density of the substance.
Compressibility is a measure of the extent to which the volume of a given
amount of substance will shrink when it is squeezed. Liquids and solids are
nearly incompressible, while gases are highly compressible. The high compressibility
of gases is a result of the large amount of empty space which exists between
gas particles. In liquids and solids very little free space is present between
particles. Compressing liquids or solids, even at high pressures, results in
very small changes in volume.
Thermal expansion is the increase in volume of a substance with increasing
temperature. For ideal gases the coefficient of thermal expansion is 1/273 at
0 degrees C. A temperature increase of 1 degrees C changes the volume of a gas
by 0.366 %. At 25 degrees C the volume of 1 g of water is 1.00296 mL, at 26
degrees C the volume is 1.00323 or a change of about 0.026 %.
Viscosity measures the resistance to flow for a liquid. The higher the viscosity
the slower a liquid will flow. From a atomic view viscosity can be understood
in terms of the ability of particles to slide over and past each other. If a
fluid has a high viscosity particles "stick" to each other and do not slide
past each other easily. One wonders what causes this stickiness in the liquid
Surface tension measures how much energy is required to move a particle from
the middle of a fluid to the surface. The easier this can occur the lower the
surface tension. Liquids under the influence of gravity and on a surface for
which the adhesive forces are small, e.g. water on wax paper, will cause the
liquid to "bead-up". In an environment where the gravity is low or absent fluids
will form a spherical shape as this shape minimizes the surface area. Water
has a comparatively high surface tension. It is difficult to move a liquid from
within the liquid to the surface. One wonders why it is difficult to move a
molecule of water from within the middle of the liquid to the surface. Water
molecules must "stick" to each other.