2. A chemical bond formed by sharing of electrons between two
elements is called as covalent bond.
3. A ̶ A (Single bond) : When 2 electrons are shared between the two
combining elements.
(Double bond) : When 4 electrons are shared between the two combining
elements.
(Triple bond) : When 6 electrons are shared between the two combining
elements.
To explain nature of chemical bond, different theories are given
(i) Octet rule
(ii) Valence bond theory
(iii) Valence shell electron pair repulsion theory
(iv) Molecular orbital theory.
4. LEWIS OCTET RULE
(I) Every atom has a tendency to complete its octet outermost.
(II) H has the tendency to complete its duplet.
(III) To acquire inert gas configuration atoms loose or gain electron or
share electron.
(IV) The tendency of atoms to achieve eight electrons in their outer
most shell is known as Lewis octet rule.
5. EXCEPTION OF OCTET RULE
(a) Incomplete octet molecules : - or (electron defficient molecules)
Compound in which octet is not complete in outer most orbit of central
atom.
Examples - Halides of IIIA groups, BF3, AlCl3, BCl3, hydride of III
A/13th group etc.
Other examples - BeCl2 (4e-), ZnCl2(4e-), Ga(CH3)3 (6e-)
6. (b) Expansion of octet or (electron efficient molecules):
Compound in which central atom has more than 8e. in outermost orbits.
Example - In PCl5, SF6, IF7, the central atom P, S and I contain 10, 12,
and 14 electrons respectively.
7. (c) I-Pseudo inert gas configuration : -
(i) Cations of transition metals, which contains 18 electrons in
outermost orbit.
Examples : Ga+3, Cu+, Ag+, Zn+2, Cd+2, Sn+4, Pb+4 etc.
(d) Odd electron molecules : -
Central atom have an unpaired electron or odd no (7e., 11e. etc) of
electrons in their outer most shell.
Examples : NO, NO2, ClO2 etc.
8. Formal charge & limitations of octet rule
Formal Charge : The formal charge of an atom in a molecule or ion is
defined as the difference between the number of valence electrons of that
atom in an isolated or free state and the number of electrons assigned to
that atom in the lewis structure.
9.
10. Limitations of the Octet Rule
The incomplete octet of the central atom : In some compounds, the
number of electrons surrounding the central atom Is less than eight. This
is especially the case with elements having less than four valence
electrons. Examples are LiCl, BeH2 and BCl3.
BeF2, BF3, AlCl3
Odd-electron molecules : In molecules with an odd number of
electrons like nitric oxide, NO and nitrogen dioxide (NO2), the octet
rule is not satisfied for all the atoms.
e.g. NO, ClO2 , ClO3
11. The expanded octet : Elements in and beyond the third period of the
periodic table have, apart from 3s and 3p orbitals, 3d orbitals also
available for bonding. In a number of compounds of these elements
there are more than eight valence electrons around the central atom.
This is termed as the expanded octet. Obviously the octet rule does not
apply in such cases.
Some of the examples of such compounds are: PF, SF6, PCl5, HNO3,
SO3, SO2, H2SO4 and a number of coordination compounds.
12. Other drawbacks of the octet theory
(i) It is clear that octet rule is based upon the chemical inertness of
noble gases. However, some noble gases (for example xenon and
krypton) also combine with oxygen and fluorine to form a number
of compounds like XeF2, KrF2, XeOF2 etc.,
(ii) This theory does not account for the shape of molecules.
(iii) It does not explain the relative stability of the molecules being
totally silent about the energy of a molecule.
