Lewis structure, formal charge, mesomerism and octet rule
The Lewis structures were developed as a well studied idea about covalent bonds by G.N.
Lewis in 1916. These are covalent bonds existing between non metallic atoms. Therefore it’s
a region of relatively high electron density between two nuclei. The two electrons in this bond
are shared by the two partners. The notation of the Lewis structures is described in literature
in different ways. Officially in IUPAC’s gold book it’s written that a Lewis structure is the
electron-dot-structure. This is different from the line formula. The only tradeoff is “Bonding
pairs of the electrons are usually denoted by lines […]” Hence in Lewis’ model the bonding
electrons can be drawn as a line. Dots should be chosen for the notation of the lone pairs.
(Mentioned Examples: H2, N2)
In the Lewis structures one phenomenon can be easily recognized, the octet rule. This rule
describes the tendency of the atoms to get eight valence electrons like the noble gases. For
hydrogen just one additional electron is needed to reach a “noble-gas-configuration”, in this
case the He-configuration. To check if this rule is realized, just count all electrons which are
distributed to one atom. For example in ammonia (NH3) all three hydrogen- atoms have two
valence electrons because of their bonds to nitrogen. They reached the He-configuration.
Nitrogen has eight valence-electrons, six from the three bonds and two from the lone pair. The
Ne-configuration is reached. In carbon monoxide (CO), the octet rule can be applied, too.
Both atoms have eight valence-electrons, six from the three bonds and two from each lone
pair. As for nitrogen in N2, the Ne-configuration is valid.
This rule works pretty fine until the third period. From here on there can be exceptions, like
the phosphorus in the phosphate (PO43-)-anion or the sulfur in the sulfur hexafluoride (SF6)molecule.
To draw a Lewis structure, use three steps. First count all valence electrons, the molecule
would have. For each negative charge of the molecule, add one electron. For each positive
charge, take one electron less. In a second step arrange the molecules in a useful way. The
third and last step is distributing all electrons corresponding to the octet rule.
The effect of non located electrons, like in the PO43—anion, leads to more than one Lewis
structure of the molecule. This is called mesomerism respectively resonance. The
P = O – double bond in the mentioned example isn’t located on one explicit oxygen. All four
oxygen-atoms can have the π-bond. Another example is the nitrate (NO3)-anion. One oxygenatom must have a double bound. This leads to three Lewis formulas of the molecule, with the
lowest energy. The mesomerism further leads to a distribution of the total charge over the
whole molecule. This leads to resonance stabilization and to a lower energy.
Apart from the mentioned total charge, there is also a formal charge to determine. This is the
charge an atom would have, if the electron pairs were separated equally, therefore
homolytically. To calculate the formal charge qf, count the number of valence electrons nv and
subtract the number of lone pairs nlp and half of the number of shared electrons ns:
qf = nv – nlp – ½ · ns. The sum of all formal charges is equal to be the total charge:
Generally you can approximate that the most realistic structure is the one with the lowest
formal charges. The most electronegative atom usually wears the neg. charge, the most
electropositive atom the positive charge. The formal charges are written in Arabic numbers to
separate them from the oxidation states (roman numbers).
(Examples mentioned: NO3- and O3.)
Tasks:
1.)
Please draw the Lewis formulas of BF3 and SF6.
2.)
Please compare the Lewis structures of the cyanate-anion in terms of formal charge and
mesomerism and name the dominant structure.
References:
- D. F. Shriver, P. W. Atkins, Inorganic Chemistry, 3rd edition, Oxford University Press,
Oxford, 1999
- E. Riedel, Anorganische Chemie, 6. Auflage, de Gruyter, Berlin/New York, 2004