I'm having trouble understanding the difference between hydrogen bonding and London forces. I know that hydrogen bonding only occurs with hydrogen but what is the difference in their actual bonding?
Molecules can attract each other at moderate distances and repel each other at close range.
The attractive forces are collectively called "van der Waals forces". Van der Waals forces
are much weaker than chemical bonds, and random thermal motion around room temperature can
usually overcome or disrupt them.
Intermolecular forces are feeble; but without them, life as we know it would be impossible.
Water would not condense from vapor into solid or liquid forms if its molecules didn't attract each other. Intermolecular forces are responsible for many properties of molecular compounds, including crystal structures (e. g. the shapes of snowflakes), melting points, boiling points, heats of fusion and vaporization, surface tension, and densities. Intermolecular forces pin gigantic molecules like enzymes, proteins, and DNA into the shapes required for biological activity.
Van der Waals' forces include all intermolecular forces that act between electrically neutral molecules. Several special cases occur.
Permanent forces occur when the interacting molecules contain groups or regions that
are permanently electron-rich or electron poor. For example, the animation at right shows
short range forces acting between molecules of gaseous HCl. The electron-rich region (on the chlorine atom)
is colored red; the electron-poor hydrogen atom is shown in blue. Notice that the molecules align
when they pass close to each other because the positive end of one molecule is attracted
to the negative end of the other. The yellow glow indicates the formation of a weak intermolecular
attraction during a close encounter. Notice that a molecule's momentum is often strong enough to overcome the
attraction and prevent it from being captured in a cluster of other molecules.
When the the molecule has a distinctly positive end and a negative end,
the permanent force is referred to as a dipole-dipole attraction. Weaker (but still noticeable) permanent forces can act between any molecules with polar bonds.
For example, the oxygen atoms in CO2 are electron-rich, while the carbon atom in the center is electron poor,
so the oxygen atom of one CO2 can be attracted to the carbon of another during very close encounters.
Hydrogen bonds are abnormally strong dipole-dipole attractions that involve molecules with -OH, -NH,
or FH groups. Hydrogen atoms are very small (with an atomic
radius of about 37 pm, they're smaller than any other atom but helium). When a
bonded electronegative atom (oxygen, nitrogen, or fluorine) pulls electrons away from the hydrogen atom,
the positive charge that results is tightly concentrated. The hydrogen is intensely attracted to small, electron-rich
O, N, and F atoms on other molecules. (Larger electron-rich groups and atoms (like -Cl, for example) will also attract
the hydrogen, but because their electrons aren't as tightly concentrated, the resulting dipole-dipole attraction is
too weak to be considered a "real" hydrogen bond.) Hydrogen bonds are essential for building biological systems:
they're strong enough to bind biomolecules together but weak enough to be broken, when necessary, at the temperatures that typically exist
inside living cells.
A polar molecule can also induce a temporary dipole in a nonpolar molecule. The electron cloud around a nonpolar molecule responds
almost instantaneously to the presence of a dipole, so this "dipole-induced dipole" force isn't as orientation-dependent as the dipole-dipole interaction.
Transitory forces arise when electron clouds oscillate in step on two molecules at close range.
Bond vibrations in molecules may produce the oscillations or they may be triggered by random, instantaneous pile-ups of electrons in atoms.
The electron-rich and electron-poor regions on the molecule may not persist for more that 10-14 or 10-15 seconds,
but if they can polarize the
electron distribution on an adjacent molecule,
electron clouds on the two molecules may begin to oscillate cooperatively with each other.
The dipoles are transitory but aligned, and a net attractive
force pulls the molecules together. At closer range, the oscillation becomes even more effective.
In the animation at left, the electrons in H2 molecules are oscillating in time to produce a tiny crystal of solid hydrogen.
The electron-rich (red) ends are aligned with electron-poor (blue) ends through nearly all of the cycle. The arrows
show the direction of the transitory dipole moment during the oscillation.
Transitory forces are sometimes called "London forces" in honor of their discoverer.
Since all molecules have electron clouds that can oscillate, London forces always contribute to intermolecular
attractions. However, they are usually weaker than the permanent forces that they are usually only invoked to
explain intermolecular forces between nonpolar molecules or noble gas atoms. However, molecules with large, diffuse electron
clouds (like polyatomic anions or molecules containing multiple bonds) can have London forces that are at least as strong as permanent forces are.
More electrons in a molecule or atom means potentially larger electron imbalances and so stronger London forces. This nicely explains the periodic trend in boiling points for the noble gases; Ne boils at a much lower temperature than Xe because, having fewer electrons, its
London forces are more easily overcome by thermal motion. It also explains why high-molecular weight nonpolar compounds tend
to be solids or liquids while light nonpolar compounds tend to gases; high molecular weight generally means more electrons and
more powerful London attractions.