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Holding Together with Van der Waals
Name: Justin
Status: student
Grade: 9-12
Location: VA
Question:
How could Van der Waals forces be strong enough to hold
paraffin molecules together to form a solid?
Replies:
Justin,
Remember that London Forces (or Van der Waals forces) depend on the
center of negativity being offset from the center of positivity. This
means that, if we consider electrons as particles, whenever the center
of mass of all the electrons do not coincide with the nucleus (the
center of mass of all the protons), then there is a momentary dipole.
This now means that the more nuclei and electrons there are, the more
possibility for these momentary dipoles to appear. Since the number
of nuclei and electrons increase with molar mass, the collective
strength of the London Forces should increase as molar mass increases.
Look up the trend of molar mass and boiling point of the following
compounds: methane, ethane, propane, butane, pentane, etc.
Greg (Roberto Gregorius)
Paraffin waxes are long molecules, 20-30 carbons long, which means each
molecule is quite large and heavy. Paraffin wax is a long "alkane" --
"alkane" is a chemical term that refers to molecules with the molecular
formula C(n)H(2n+2), in other words just carbon and hydrogen, no other
elements. You might know that the bigger the molecule, the more likely it is
to be a solid at a given temperature. In the case of alkanes, the smallest
ones, like methane CH4 or ethane C2H6, are gases at room temperature and
pressure. Some larger alkanes, like octane or hexane, are liquids at room
temperature and pressure. Very large ones, like the 25-carbon paraffin wax,
are solids at room temperature and pressure.
Generally speaking, van der Waals forces are any non-covalent forces that
attract molecules together. In terms of why a given substance is a gas or a
liquid or a solid, you have figure out which state the substance is
"happiest" in -- in this case "happy" means lowest energy. With very large
molecules, there are simply more spots to be attracted, and because the
molecule is larger, it is vibrating comparatively less than smaller
molecules. The reason these molecules are attracted to each other is because
it takes less energy to be next to another thing that is like them (you may
have heard the saying "like attracts like") that being next to something
different. So paraffin would rather be next to paraffin than it would next
to air or water, or something else different. Pressure plays a big role here
-- at higher pressures, the molecules are pushed closer together, and if
they are attractive, this makes them more likes to condense instead of be a
gas. Butane is a liquid inside a lighter (it is pressurized), but becomes a
gas when the pressure is released. Add up all the effects together, and you
can figure out if the molecules are "happiest" as a gas, liquid, or solid at
a given set of conditions.
With paraffin, you have a substance that is "happier" being a solid than it
is being a liquid or gas at room temperature. However, if you add a little
energy -- make it jiggle a touch more -- you can get it to liquefy. Paraffin
wax melts easy, at relatively low temperatures. What happens here is the
thermal vibrations overcome the intermolecular attractions, and the solid
melts.
Hope this helps,
Burr Zimmerman
Hi Justin,
One source of Van de Waal's forces is that momentary
asymmetry in the electron distribution of even a symmetric
molecule such as the long -CH2-CH2-CH2- chains that make up
paraffin, act like temporary dipoles that induce similar
dipoles in nearby molecules. Thus, the molecules attract each
other and tend to form an orderly, almost crystalline
arrangement. This is especially prevalent in long thin
molecules like paraffin, and much less so in the shorter (but
otherwise similar) molecules like Ethane, Butane, and so on.
The relative strength of these forces in long chain molecules
like paraffin, accounts for this material's being a solid at
room temperature. Medium length molecules like Cetane are
semi-oily liquids. Even shorter ones like Pentane or Octane
are low viscosity liquids, and even shorter ones like Ethane
or Butane are gases at room temperature.
Regards,
Bob Wilson.
Paraffin molecules are really big, so there are lots of van
der Waals interactions holding them together. If a single
pebble were to land on you, you could walk without difficulty.
If a cubic meter of gravel were to fall on you, you probably
would not be able to move. The van der Waals interactions that
stick paraffin molecules to each other are like those pebbles.
In large numbers, they add up.
Richard Barrans
Department of Physics and Astronomy
University of Wyoming
The origin of the van der Waals attraction is the dipole induced
by the motions of electrons in a molecule (call it A) by all its
neighbors, this dipole in A in turn induces dipoles in all the
neighbors. There are two factors which make this apparently weak
interaction significant. First, it is always attractive. That is
the field of an induced dipole (+)------(-) always induces a
dipole with the opposite direction (-)--------(+). That makes
the interaction attractive (like the poles of a magnet, but one
in which the interaction is turning "on" and "off" very rapidly.
Second, the induced-dipole / induced-dipole interaction (an
alternative name for the van der Waals interaction) occurs
between all electrons, they don't have to have a particular
molecular structure -- so the interaction is always "on" --
whether the molecule is ionic, dipolar, or whatever.
Hydrocarbons, in addition, have three other factors at work.
First if the hydrocarbon has non-symmetrical structure, it has
a permanent dipole, which can interact with similar dipolar
hydrocarbons both by a dipole to dipole interaction. Second,
by an interaction between the permanent dipole and by a
transient dipole that it induces in its neighbors. Thirdly,
if the hydrocarbon has more than 6 carbons in a row. the
chains can get tangled producing a structure (in the extreme)
of a bowl of cooked spaghetti. In high molecular weight waxy
hydrocarbons this is a major contributor to holding the
hydrocarbon together as a solid.
Vince Calder
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Update: June 2012
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