Dipole Moments
Dipole Moments
When a
covalent bond is formed between two identical atoms, e.g. H-H, Cl-Cl, etc., the
two electrons forming the covalent bond may be regarded as being symmetrically
disposed between the two atoms. The centres of gravity of the electrons and
nuclei therefore coincide. With two dissimilar atoms the two electrons
are no longer symmetrically disposed, because each atom has a different electronegativity,
i.e., attraction for electrons. Chlorine has a much greater electronegativity
than hydrogen that when chlorine and hydrogen combine to form covalent hydrogen
chloride, the electrons forming the covalent bond are displaced towards the
chlorine atom without any separation of the nuclei :
The hydrogen
atoms will, therefore, be slightly positively charged, and the chlorine atom
slightly negatively charged. Thus, owing to the greater electronegativity of
the chlorine atom, the covalent bond in hydrogen chloride is characterized by
the separation of small charges in the bond. A covalent bond such as this, in
which one atom has a larger share of the electron-pair, is said to possess partial
ionic character.
In analogy
with a magnet, such a molecule is called a dipole, and the product of
the electronic charge, e, and the distance d, between the charges (positive and
negative centres) is called the dipole moment, μ ; i.e., μ= e✕d. e is the order of 10-10
e.s.u. ; d, 10-8 cm. Therefore μ is of the order of 10-18 e.s.u., and this unit is known as the
Debye (D), in honour of Debye, who did a large amount of work on dipole moments.
The dipole
moment is a vector quantity, and its direction is often indicated by an arrow
parallel to the line joining the points of charge, and pointing towards the
negative end.
e.g., H-Cl. The
greater the value of the dipole moment, the greater is the polarity of
the bond. The terms polar and non-polar are used to describe
bonds, molecules and groups, and the reader is advised to make sure he
appreciates how the terms are applied in each case under consideration.
Not only do polar bonds contribute to the dipole moment of a molecule, but so do long pairs, e.g., the lone pair in ammonia makes a large contribution to the dipole moment.
In the table
given below gives some electronegativity values, and it will be seen that
electronegativity increases from left to right and decreases from top to bottom
of the periodic table. The values are not absolute values; they are relative
values and this is a satisfactory scheme since their use involves differences
in electronegativities. Furthermore, it should be noted that electronegativity
deals with electron attraction within a molecule, whereas electron
affinity deals with the attraction of an electron from outside the atom and has
an absolute value.
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