The lewis theory gave the first explanation of a covalent bond in terms of electrons that was generally accepted. If two electrons are shared between two atoms, this constitutes a bond and binds the atoms together. For many light atoms, a stable arrangement is attained when the atom is surrounded by eight electrons.

The octet can be made up from some electrons which are totally owned and some electrons which are `text(shared)`. Thus atoms continue to form bonds until they have made up an octet of electrons. This is called the 'octet rule'.The octet rule explains the observed valences in a large number of cases. There are exceptions to the octet rule; for example, hydrogen is stable with only two electrons.

The conventional Lewis structure representation of a pair of bonded electrons is by means of a `text(dash)` `(-)` usually called a `text(bond)`. Lone pairs or `text(non-bonded)` electrons are represented by `text(dots)`. Some structures are represented below:

Such representations of organic molecules are not usually problematic. However, `text(hit-and-trial)` is generally the method (obviously not very efficient) used by most students in figuring out the structures of inorganic molecules.

The formula of a molecule shows the number of atoms of each element but does not show the bonding arrangement of the atoms. To represent the bonding pattern in a molecule, the electron dot symbols of the elements are arranged such that the shared pairs and unshared pairs (called lone pairs) are shown and the octet rule (or duet for hydrogen) is satisfied. For example


A molecule of fluorine is shown as `: overset( . .) underset( .. )F : overset( .. ) underset( .. )F : ` or

` : overset( . .) underset( .. )F - overset( .. ) underset( .. )F : `

And a molecule of hydrogen fluoride is shown as

` H : overset( . . ) underset( * * ) F : ` or `H- overset ( * * ) underset (* *) F : `

Arrangement of dot symbols used to represent molecules are called Lewis structures . Lewis structures do not convey any information regarding the shape of the molecule. Usually, the shared pairs of electrons are represented by lines between atoms and any unshared pairs are shown as dot pairs.

Lewis structures are written by fitting the element dot symbols together to show shared electron pairs and to satisfy the octet rule. For example,

(i) In water (`H_2O`), one `H` and two `O ` complete their duet and octet respectively

` : overset( * * ) underset( underset H | ) O - H `

(ii) In ammoma `(NH_3)`, three `H` and one `N` fit together and satisfy their duet and octet respectively as

`H- overset (* *) underset ( underset H |)N- H`

(iii) In carbon tetrachloride `(C Cl_4)`, four `Cl` and one `C` complete their octet as shown in fig.

For the given molecules, we have adopted hit & trial method to fit the dot symbols together and satisfy the octet rule. But remember that hydrogen form one bond, oxygen forms two bonds, nitrogen three bonds and carbon forms four bonds. For simpler molecules, the hit & trial method works perfectly but for slightly complicated polyatomic species, this may give us more than one possible structure. Thus, a systematic approach is needed to design the Lewis structures of such polyatomic species. But before proceeding further, let us understand the limitation of this approach.

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.

(a) Molecules, which have contraction of octet. Such molecules are electron deficient. For example, `BH_3`, `BF_3`, `BCl_3`, `AlCl_3`, `GaCl_3` etc.

(b) 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, `PCl_5`, `SF_6` etc.

(c) Molecules containing odd number of electrons (in total) cannot satisfy octet rule. Such species are called odd electron species and are paramagnetic in nature due to presence of unpaired electron. For example, `NO`, `NO_2` and `ClO_2`.

To draw the lewis stryctyres of polyatomic species follow the given sequence

(i) First calculate `n_1`.

`n_1 =` Sum of valence electron of all the atoms of the species `pm` 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 and for a dinegatively charged species, add
`2` and so on.

(ii) Then calculate `n_2`.

`n_2 =` ( `8xx`number of atoms other than `H`) `+` (`2xx`number of `H` atoms)

(iii) Subtract `n_1` from `n_2`, which gives `n_3`.

`n_3 = n_2 - n_1 =` number of electrons shared between atoms = number of bonding electrons.

`n_3 /2 =(n_2-n_1)/2=` number of shared (bonding) electron pairs = number of bonds.

(iv) Subtracting `n_3` from `n_1` gives `n_4`.

`n_4 = n_1 - n_3 =` number of unshared electrons or non-bonding electrons.

`n_4/2 =(n_1-n_3)/2=` number of unshared electron pairs = number of lone pairs.

(v) 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. Here, you are required to know a bit of chemistry, physics or mathematics won't help.

(vi) Now around the central atom, place the other atoms and distribute the required number of bonds (as calculated in step (iii))& required number of lone pairs (as calculated in step (iv) ), keeping in mind that every atom gets an octet of electrons except hydrogen.

(vii) Then calculate the formal charge on each atom of the species. Formal charge is the difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis Structure.

Formal charge on an atom = number of valence electrons of the atom - (number of shared electrons of that atom + number of unshared electrons of that atom).

Formal charge on an atom = number of valence electrons of the atom - number of bonds formed by that atom - number of unshared electrons (`2xx` lone pairs) of that atom.

For every electron of an atom that is shared in a bond, the `text(number of bonds formed by the atom)` is one. Therefore if an atom forms only one bond (`A- B`), one electron of the bond is that of `A` and other is that of `B`. So the "number of bonds" of A and B each is one. But if the bond were a co-ordinate bond (`A-> B`), then two electrons of `A` are involved in it. This makes the number of bonds of `A` to be `2` and that of `B` to be zero.

(viii) 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.

(ix) The given Lewis structure should account for the factual aspects of the molecule like resonance (delocalization), bond length, `p-d` back bonding etc. Sometimes, there are more than one acceptable Lewis structure for a given species. In such cases, we select the most plausible Lewis structure by using formal charges and the following guidelines:

`ast` For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.

`ast` Lewis structures with large formal charges (`+2`, `+3` and/or `-2`, `-3` and so on) are less plausible than those with small formal charges.

`ast` Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.