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London Forces or van der Waals Forces Dipole-Dipole Attraction Hydrogen Bonding van der Waals Forces and London Dispersion forces
Simplified Explanation van der Waals Forces (aka London Forces (Fritz London 1930), Weak Intermolecular Forces) are very weak forces of attraction between molecules resulting from:
****This weak intermolecular attraction is why nonpolar gases in our atmosphere can form liquids under extreme pressure and low temperatures. The more electrons that are present in the molecule, the stronger the dispersion forces will be. Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for example, dispersion forces operate between
A More Complex Explanation The attractive forces that exist between molecules
are responsible for many of the bulk physical properties exhibited
by substances. Some compounds are gases, some are liquids, and
others are solids. The melting and boiling points of pure substances
reflect these intermolecular forces, and are commonly used for
identification. Of these two, the boiling point is considered the
most representative measure of general intermolecular attractions.
Thus, a melting point reflects the thermal energy needed to convert
the highly ordered array of molecules in a crystal lattice to the
randomness of a liquid. The distance between molecules in a crystal
lattice is small and regular, with intermolecular forces serving to
constrain the motion of the molecules more severely than in the
liquid state. Molecular size is important, but shape is also
critical, since individual molecules need to fit together
cooperatively for the attractive lattice forces to be large.
Spherically shaped molecules generally have relatively high melting
points, which in some cases approach the boiling point, reflecting
the fact that spheres can pack together more closely than other
shapes. This structure or shape sensitivity is one of the reasons
that melting points are widely used to identify specific compounds.
In general, larger molecules have higher boiling points than smaller molecules of the same kind, indicating that dispersion forces increase with mass, number of electrons, number of atoms or some combination thereof. The following table lists the boiling points of an assortment of elements and covalent compounds composed of molecules lacking a permanent dipole. The number of electrons in each species is noted in the first column, and the mass of each is given as a superscript number preceding the formula.
Two ten electron molecules are shown in the first row. Neon is
heavier than methane, but it boils 84º lower. Methane is composed
of five atoms, and the additional nuclei may provide greater
opportunity for induced dipole formation as other molecules
approach. The ease with which the electrons of a molecule, atom or
ion are displaced by a neighboring charge is called polarizability,
so we may conclude that methane is more polarizable than neon. In
the second row, four eighteen electron molecules are listed. Most of
their boiling points are higher than the ten electron compounds neon
and methane, but fluorine is an exception, boiling 25º below
methane. The remaining examples in the table conform to the
correlation of boiling point with total electrons and number of
nuclei, but fluorine containing molecules remain an exception. Of course, boiling point relationships may be dominated by even stronger attractive forces, such as those involving electrostatic attraction between oppositely charged ionic species, and between the partial charge separations of molecular dipoles. Molecules having a permanent dipole moment should therefore have higher boiling points than equivalent nonpolar compounds, as illustrated by the data in the following table.
In the first row of compounds, ethane, ethene and ethyne have no molecular dipole, and serve as useful references for single, double and triple bonded derivatives that do. Formaldehyde and hydrogen cyanide clearly show the enhanced intermolecular attraction resulting from a permanent dipole. Methyl fluoride is anomalous, as are most organofluorine compounds. In the second and third rows, all the compounds have permanent dipoles, but those associated with the hydrocarbons (first two compounds in each case) are very small. Large molecular dipoles come chiefly from bonds to high-electronegative atoms (relative to carbon and hydrogen), especially if they are double or triple bonds. Thus, aldehydes, ketones and nitriles tend to be higher boiling than equivalently sized hydrocarbons and alkyl halides. The atypical behavior of fluorine compounds is unexpected in view of the large electronegativity difference between carbon and fluorine. from-www.cem.msu.edu/~reusch/
Van der Waals equation Johannes
Van der Waals was interested in the kinetic theory of
gases and fluids, and his primary work was to develop an
equation which applied to real gases, unlike that of
Robert Boyle which assumes that there are no attractive
forces between molecule and that molecules have zero
volume. (P + a/v2)(v-b) = RT , where a and b are constant for a particular gas. Van der Waals was awarded a Nobel prize in 1910 for his work on the equation of state of gases and liquids. The
weak, electrostatic attractions between atoms were named
Van der Waals forces in his honor. from-http://www.chemsoc.org/timeline/pages/1920.html London Forces or van der Waals Forces Dipole-Dipole Attraction Hydrogen Bonding On to Like Dissolve Like |
