Lewis structures, formal charge, mesomerism, octet rule,
oxidation number *
Gilbert N. Lewis in 1916: A covalent bond is formed when two neighboring atoms
share an electron pair. Double or triple bonds consist of two or three shared pairs.
Unshared pairs of valence electrons are called free electron pairs or Ione pairs. They
do not contribute directly to the bonding, but have an influence on the shape and the
chemical properties of a molecule.
To construct a Lewis structure, the first step is to count the valence electrons of all
atoms. In Lewis structures, single electrons are painted as dots, electron pairs as a
line (bar).
Examples: H2 and Cl2.
Octet rule (including exceptions)
In covalent molecules the atoms “want to” reach a stable 8-electron noble gas
configuration by sharing binding electron pairs. This tendency is called octet rule.
Elements of the second period must fulfill the octet rule.
Examples: PF3, CH4
Elements of period 3 or higher can cross the octet rule. The central atom has formally
more than 8 electrons. This is named Hypervalency.
Examples: HClO4, SO2
Other exceptions are radicals (e.g. NO, NO2) which have an odd number of
electrons. Radicals are mostly unstable. (e.g. NO and NO2 dimerize at lower
temperature).
Formal charge
To determine the formal charge of atoms in a molecule shared electron pairs are
divided homolytical. The number of this electrons together with the number of lone
pair electrons (ore single electrons) are compared with the number of electrons in the
respective element. In Lewis structures the formal charges should be as small as
possible and the number of atoms with a charge must be minimized.
Example: HNO3, NH4+
But there must be a differentiation between formal an effective charge. If for example
in NH4+ the formal charge is +1 located on the Nitrogen, the whole molecule is a +1
cation but with the charge distributed over all involved atoms.
Oxidation number
The concept of formal charge is different to the concept of oxidation number. To
identify the oxidation number of an atom in a molecule, the whole binding electron
pair “belongs” to the atom with the higher electronegativity (heterolytical division).
For equal atoms it is divided equally. The number of this electrons together with the
number of lone pair electrons (ore single electrons) are compared with the number of
electrons in the respective element.
Oxidation numbers facilitate the formulation of redox equations. They are also good
for the classification of compounds and to explain certain properties. In this
parameter, a bond has a more ionic character.
Mesomerism
Ozon O3 has more than one possible Lewis formula. Actually the structure needs to
be understood as an in-between of these formulas. To express this mesomerism in
Lewis structures, all formulas must be written with a double arrow between them.
There are no differences between the bonds. Normally, a single bond is longer than a
double bond but, as implied in resonance formulas, they are equal (O3: 128 pm).
Examples: O3, CO32References
[1]
[2]
[3]
E. Riedel, Anorganische Chemie, 6. Auflage, de Gruyter, Berlin/New York,
2004
D. F. Shriver, P. W. Atkins, Inorganic Chemistry, 3rd edition, Oxford University
Press, Oxford, 1999
C. E. Mortimer, U. Müller, Chemie, 8. Auflage, Thieme Verlag, Stuttgart, 2003
* summary of Konrad Kiefer, WS 2008/2009