Each element in universe has tendency to acquire stability (state of minimum energy). That is done by combining one atom with another leading to the formation of a compound. This can be achieved in 2 ways:
(1) Electron transfer giving rise to an ionic bond via formation of anion and cation. Force of attraction are electrostatic interaction.
(2) By sharing of electrons which gives rise to a chemical bond. While the formation of chemical bond more than 40kJ/mole of energy is released.
(3) Special Case: Coordinate bond or Dative bond: The bond formed between two atom in which contribution of an electron pair is made by one of them and the sharing is done by both. For the formation of coordinate bond it is necessary to have atleast one covalent bond in the compound. Chemical bonding
THE OCTET RULE
For an atom there must be maximum of eight electron in the valence shell.
Deviation from the Octet Rule: A number of cases are known where the combining species or atoms have less than eight (i.e. incomplete octet) or, more than eight (i.e. expansion of octet) electrons in the covalent bonded molecules.
Incomplete Octet: In the molecules such as BeCl2, BCl3 and NO (: N g = O), the central atoms, i.e. Be, B and N bear four, six and seven electrons respectively. Here it is noteworthy that other atoms except the central one in the above compounds follow the octet rule. As a matter of fact, a large number of compounds formed by Be (quarted), M (sextet) where M = B, Al, Ga are known to have incomplete octets.
These compounds are very often referred to as electron deficient compounds characterised by a profound tendency to receive back a lone pair (for Gr 13 or III A elements, e.g. B, Al, Ga etc.) or two lone pairs (for Gr 2 or II A elements, e.g. Be) to attain the octet. This is why, the electron deficient compounds act as Lewis acids.
Expansion of Octet: In the compounds like, PCl5, ClF3, SF6, SiF62– etc. the central elements, i.e. P, Cl, S and Si, are bearing ten, ten, twelve and twelve electrons respectively to display the expansion of octet. Similarly, in OsF8, there are sixteen electron around Os. Here also, except the central atoms, all other atoms satisfy the octet rule. Be forms bond with 2 Cl atom to complete its octet RULES FOR LEWIS STRUCTURE DRAWING
(1) Calculate n1 = no. of valence shell electrons of all atoms + no. of negative charge (if any) – no. of positive charge (if any)
(2) Calculate n2 = (no. of H-atom × 2) + (no. of atoms other than H-atoms × 8)
(3) Calculate n3 = n2 – n1 ⇒ no. of shared electrons; i.e. No. of no. bonds = .......(a)
(4) Calculate n4 = n1 – n3 ⇒ no. of unshared electrons; i.e. no. of lone pairs = .......(b)
Using the informations (a) & (b) the structure is to be assigned as follows :
(i) Find out the central atom first (i.e. either least in number or more electro positive)
(ii) Allocate the surrounding atoms around the central atom with the help of bonds available in (a).
(iii) To fulfill the octet of each atom, utilise the lone pairs available in (b).
(iv) Finally calculate the formal charge for each atom and assign the atoms according to the formula given
F.C. (on an atom) = Valence shell electron of that atom – no. of bonds associated with it – no. of unshared electrons on it.
In many molecules, the choice of atoms which are connected by multiple bonds is arbitrary. When several choices exist, all of them should be drawn. For example, as shown in Figure, three drawings (resonance structures or canonical forms) of CO32– are needed to show the double bond in each of the three possible C–O positions. Example of resonance
In fact, experimental evidence shows that all the C–O bonds are identical, with bond lengths (129 pm) between double-bond and single bond distances (116 and 143 pm, respectively); none of the drawing alone is adequate to describe the molecular structure, which is a combination of all three.
Thus resonance signifies that there is more than one possible way in which the valence electrons can be placed in a Lewis structure. Note that in resonance structures such as those shown for CO32– in Figure, the electrons are drawn in different places but the atomic nuclei remain in fixed positions.
The species CO32–, NO3–, and SO3 are isoelectronic (have the same electronic structure). Their Lewis diagram are identical, except for the identity of the central atom.
When a molecule has several resonance structures, its overall electronic energy is lowered, making it more stable.
Bond order =
Resonance vs. tautomerism: The proposed canonical forms should not differ in the atomic arrangements. Only the position of the pi-electrons is changed from one position to another (generally the sigma-electrons do not get delocalised) . Shifting of an atom leads to tautomerism not resonance and such species are related through the equilibrium sign ( ). Some examples of tautomerism are shown below.
(a) The nitrous acid may have two forms in equilibrium.
The coloured organic nitrites and some inorganic nitrites such as silver nitrite (pale yellow), mercurous nitrite (pale yellow) probably exist in nitro forms.
(b) phosphorous acid (H3PO3) and hypophosphorous acid (H3PO2) exist in tautomeric equilibria.
Phosphorous acid →Phosphonic acid Hypophosphorous acid→Phosphinic acid
minor tautomer→Major tautomer minor tautomer →major tautomer
(c) Similarly, tautomerism in sulphurous acid is also reported.
All the examples of tautomerism cited above involve the shifting of a proton and this type of tautomerism is referred to as prototropic tautomerism which is very important in organic chemistry (e.g. keto-enol tautomerism.)
Important points in resonance
(i) The canonical forms should have the same number of unpaired electrons.
(ii) In the case of charge separation, if the adjacent atoms bear the same change then the electrostatic repulsion will destabilise the structure.
(ii) Placement of opposite charge on the adjacent atoms stabilises the system through an electrostatic interaction.
(iv) The canonical forms in whch the negative charge resides on the electrnogative atoms contribute more.
The formal charge is the charge which an atom would have if electron pairs were shared equally; Lewis structures with low formal charges typically have the lowest energy.
Formal charge of an atom is the difference between the number of valence electrons in an isolated atom (i.e. free atom) and the number of electrons assigned to that atom in a Lewis structure.