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Chapter - 4
Chemical bonding and Molecular Structure
Some Important Points and Terms of the Chapter
1. Lewis dot structures are shorthand to represent the valence electrons of an atom. The structures
are written as the element symbol surrounded by dots that represent the valence electrons.
2. Covalent Bonds- The bond formed between two atoms by mutual sharing of electrons between
them so as to complete their octets or duplets. When two atoms share one electron pair they are
said to be joined by a single covalent bond.e.g H₂ If two atoms share two electron pairs of
electrons, the covalent bond between them is called a double bond. e.g O₂ If two atoms share
three electron pairs of electrons, the covalent bond between them is called a double bond. e.g N₂
3. Octet Rule-Kossel 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 or by sharing of
valence electrons in order to attain their octet. This is known as octet rule.
4. Limitations of octet rule-
a) 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- LICI,BeCl₂, BCI,
b) Odd-electron molecules: In molecules with an odd number of electrons like nitric oxide, NO and
nitrogen dioxide, the octet rule is not satisfied for all the atoms.
c) The expanded octet: Elements in and beyond the third period of the periodic table have, apart
from 38 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. Some of examples of such compounds are: PFs, SFG-
d) This theory does not account for the shape of molecules.
5. Electrovalent bond or Ionic Bond: The che
