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Lewis Structures of Molecules

To draw the Lewis structures of polyatomic species, follow the given sequence:
  1. First calculate n1:
    n1 = Sum of valence electron of all the atoms of the species
      ± Net charge on the species.
    For a negatively charged species, electrons are added, while for positively charged species, the electrons are subtracted. For an uninegatively charged species, add 1 to the sum of valence electrons; for a dinegatively charged species, add 2 and so on.
  2. Then calculate n2:
    n2 = (8 × Number of atoms other than H) + (2 × Number of H atoms)
  3. Subtract n1 from n2, which gives n3.
    n3 = n2n1 = Number of electrons shared between atoms
    = Number of bonding electrons
    Description: 14975.png = Number of shared (bonding) electron pairs
    = Number of bonds
  4. Subtracting n3 from n1 gives n4.
    n4 = n1n3 = Number of unshared electrons or non-bonding electrons
    Description: 14984.png = Number of unshared electron pairs = Number of lone pairs
  5. Identify the central atom. Generally, the central atom is the one which is least electronegative of all the atoms, when the other atoms do not contain hydrogen. When the other atoms are hydrogen only, then the central atom would be the more electronegative atom. However, some exceptions are possible, for example Cl2O.
  6. Now around the central atom, place the other atoms and distribute the required number of bonds (as calculated in step 3) and required number of lone pairs (as calculated in step 4), keeping in mind that every atom gets an octet of electrons, except hydrogen.
  7. Then calculate the formal charge on each atom of the species.
    Formal charge on an atom = Number of valence electrons of the atom
    – Number of bonds formed by that atom
    – Number of unshared electrons
    (2 × lone pairs) of that atom
  8. When two adjacent atoms get opposite formal charges, then charges can be removed by replacing the covalent bond between the atoms by a dative (co-ordinate) bond. This bond will have the arrowhead pointing towards the atom with positive formal charge. It is not mandatory to show the dative bonds unless required to do so.
  9. The given Lewis structure should account for the factual aspects of the molecule such as resonance (delocalization), bond length, and pπ–dπ back bonding.
This method would be applicable to only those molecules/species which follow octet rule, except hydrogen.
There are three kinds of molecules/species which do not follow octet rule:
  1. Molecules which have contraction of octet: Such molecules are electron deficient.
    For example, BH3, BF3, BCI3, AlCI3, and GaCl3.
  2. Molecules which have expansion of octet: Such species have more than eight electrons in their outermost shell. This is possible in those molecules which have vacant d orbitals; thus they can expand their octet.
    For example, PCI5 and SF6.
  3. Molecules containing odd number of electrons (in total): Such molecules cannot satisfy octet rule. Such species are called odd electron species and are paramagnetic in nature due to the presence of unpaired electron.
    For example, NO, NO2, and ClO2.

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