Electron dot structures & formal charge
For a few years after obtaining his doctorate, Lewis worked and studied both in the United
States and abroad (including Germany and the Philippines) and he was even a professor at
M.I.T. from 1907 until 1911. He then went on to U.C. Berkeley in order to be Dean of the
College of Chemistry in 1912.
In 1916 Dr. Lewis formulated the idea that a covalent bond consisted of a shared pair
of electrons. His ideas on chemical bonding were expanded upon by Irving Langmuir and
became the inspiration for the studies on the nature of the chemical bond by Linus Pauling.
In 1923, he formulated the electron-pair theory of acid-base reactions. In the so-called Lewis
theory of acids and bases, a ”Lewis acid” is an electron-pair acceptor and a ”Lewis base” is
an electron-pair donor.
In 1926, he coined the term ”photon” for the smallest unit of radiant energy.
Lewis was also the first to produce a pure sample of deuterium oxide (heavy water) in 1933.
By accelerating deuterons (deuterium nuclei) in Ernest O. Lawrence’s cyclotron, he was
able to study many of the properties of atomic nuclei.
During his career he published on many other subjects, and he died at age 70 of a heart
attack while working in his laboratory in Berkeley. He had one daughter and two sons; both
of his sons became chemistry professors themselves.
9.2 Formal Charge
Theformal chargeof an atom is the charge that it would have if every bond were 100%
covalent (non-polar). Formal charges are computed by using a set of rules and are useful
for accounting for the electrons when writing a reaction mechanism, but they don’t have
any intrinsic physical meaning. They may also be used for qualitative comparisons between
different resonance structures (see below) of the same molecule, and often have the same
sign as the partial charge of the atom, but there are exceptions.
The formal charge of an atom is computed as the difference between the number of valence
electrons that a neutral atom would have and the number of electrons that ”belong” to it
in the Lewis structure when one counts lone pair electrons as belonging fully to the atom,
while electrons in covalent bonds are split equally between the atoms involved in the bond.
The total of the formal charges on an ion should be equal to the charge on the ion, and the
total of the formal charges on a neutral molecule should be equal to zero.
For example, in the hydronium ion, H 3 O+, the oxygen atom has 5 electrons for the purpose
of computing the formal charge—2 from one lone pair, and 3 from the covalent bonds with
the hydrogen atoms. The other 3 electrons in the covalent bonds are counted as belonging
to the hydrogen atoms (one each). A neutral oxygen atom has 6 valence electrons (due to
its position in group 16 of the periodic table); therefore the formal charge on the oxygen
atom is 6 – 5 = +1. A neutral hydrogen atom has one electron. Since each of the hydrogen
atoms in the hydronium atom has one electron from a covalent bond, the formal charge on
the hydrogen atoms is zero. The sum of the formal charges is +1, which matches the total
charge of the ion.