Slide contain information on Octet rule and its Drawback Rules to draw Lewis structure Calculation of formal charge

1. Nature of Chemical
Bonding
(NOCB)
• Kossel and Lewis
theory
• Valance Bond
theory(VBT)
• Molecular Orbital
theory(MOT)

2. KÖSSEL-LEWIS
APPROACH
TO CHEMICAL
BONDING
Lewis pictured the atom in terms of a
positively charged ‘Kernel’ (the nucleus plus
the inner electrons) and the outer shell that
could accommodate a maximum of eight
electrons
This octet of electrons, represents a
particularly stable electronic arrangement.
Lewis postulated that atoms achieve the
stable octet when they are linked by
chemical bonds.
E.g. F2, O2 etc

3. Lewis symbol
• outer shell electrons take part in chemical combination and they are known
as valence electrons.
• The inner shell electrons are well protected and are generally not involved in
the combination process.
• Thus to represent the valence elcetrons lewis developed an representation
method called the Lewis symbol

4. Significance
• This number of valence electrons helps to calculate the common or group valence of the
element.
• The group valence of the elements is generally either equal to the number of dots in Lewis
symbols or 8 minus the number of dots or valence electrons.
• E.g. Na  has 7e- in valence shell so valency is 1

5. Kossel- Lewis
• First time explained a type of bond that is the electrovalent/ionic bond
• They explained how the highly electropositive Grp 1 elements bond with the
group 17 electronegative elements.
• They bond in order to achieve the stable noble gas configuration
• The duplet or octet state in their valence orbitals.
• Thus the bonding formed in order to achieve it is called electrovalent bond
• The electrovalence is thus equal to the number of unit charge(s) on the ion

6. Octet rule
Kössel and Lewis in 1916 developed an important theory of chemical
combination between atoms known as electronic theory of chemical
bonding.
According to this, atoms can combine either by transfer of valence
electrons from one atom to another (gaining or losing) or by sharing of
valence electrons in order to have an octet in their valence shells. This
is known as octet rule

7. Langmuir’s theory of Covalent bond
• Each bond is formed as a result of sharing of an electron pair between
the atoms.
• Each combining atom contributes at least one electron to the shared
pair.
• The combining atoms attain the outershell noble gas configurations
as a result of the sharing of electrons
E.g. Formation of Cl2

8. • If two atoms share two pairs of electrons, the covalent bond
between them is called a double bond. For example, in the carbon
dioxide molecule
• When combining atoms share three electron pairs as in the case of
two nitrogen atoms in the N2 molecule and the two carbon atoms in
the ethyne molecule, a triple bond is formed

9. Rules to write Lewis structure – E.g.
CO32-
1. Write the symmetrical skeleton for the polyatomic ions
O C O 2- C1s22s22p2 O1s22s22p4
O
2. Calculate the number of electrons available in the valence shell of all atoms (A)
A= 1x4(C)+3x6(O)+2(for extra 2 e-)=24 e-
3. Calculate the total number of electrons needed by atoms to accqire the noble gas config(N)
N=1x8+3x8=32 e-
4. Calculate total number of electrons shared(S) i.e S=N-A
S=32-24=8 e- i.e 8/2=4 pairs of electrons

10. 5. Place the shared electrons in the skeleton.
Use = and triple bonds wherever necessary
[O:C::O]2-
..
O

11. E.g. Ozone(O3)
1. Skeleton - O1s22s22p4
O
O O
2. Calculate A
A=3x6=18e-
3. Calculate N
N=3x8=24e-
4. S=N-A=24-18=6e- i.e 6/2=3 electron pairs

12. Bond order
• Bond Order is given by the number of
bonds between the two atoms in a
molecule
• Isoelectronic molecules and ions have
identical bond orders; for example, F2
and O2 have bond order 1.
• N2 , CO and NO+ have bond order 3

13. Formal charge
• The formal charge of an atom in a polyatomic molecule or ion may be 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
• Formal charge (F.C.) on an atom in a Lewis structure =
total number of valence electrons in the free atom
— total number of non bonding (lone pair) electrons
— (1/2) total number of bonding(shared) electrons
• F.C = VE-LP-1/2 x BE

14. Example

16. Significance of formal charge
• It helps to calculate the charge on the atoms in a Lewis structure
• It helps to calculate the number of valence electrons
• Formal charges help in the selection of the lowest energy structure
from a number of possible Lewis structures for a given species

17. Drawbacks of
octet rule
The incomplete octet of the
central atom
Odd electron molecules
Expanded octet

18. Incomplete
octet
• 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

19. 2. 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

20. 3. 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
• examples of such compounds are: PF5 , SF6 , H2SO4 and a number
of coordination compounds.

21. Other drawbacks
• 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.,
• This theory does not account for the shape of molecules.
• It does not explain the relative stability of the molecules being totally
silent about the energy of a molecule.