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. Important in constructing bonds between atoms is the number of valence electrons. The first step is to count the valence electrons. In Lewis structures, single electrons are painted as dots, electron pairs as a line. 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 an atom 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 to 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 to the number of electrons in the respective element. Oxidation 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). 2- Examples: O3, CO3

References

[1] E. Riedel, Anorganische Chemie, 6. Auflage, de Gruyter, Berlin/New York, 2004 [2] D. F. Shriver, P. W. Atkins, Inorganic Chemistry, 3rd edition, Oxford University Press, Oxford, 1999 [3] C. E. Mortimer, U. Müller, Chemie, 8. Auflage, Thieme Verlag, Stuttgart, 2003

Questions

2- 1) Please write down all resonance formulas of CO3 , N2O and O3. Which don’t come into consideration? Give the oxidation numbers for all elements in these molecules. 2) Explain the difference of formal and effective charge and oxidation number 2- using the example of CO3 .