Chapter 3: – Chemical Bonding
B.sc 1st year Book
(Page 5)

Covalent Bond | Octet rule

Introduction

Lowis-Langmuir Concept (1916−19): Lewis proposed that an atom attain inert gas configuration (1s2 or ns2np6) by sharing one or more electron pairs by the combination of similar or dissimilar atoms. Langmuir called these shared electron pairs between two or more atoms as ‘Lewis electron-pair bond’ or ‘covalent bond’ or sometimes as ‘non-polar’ bond because it does not acquire polarity. Thus, the concept of a covalent bond is also known as the Lewis-Langmuir concept and the compounds formed by the sharing of electrons of their atoms are called ‘Covalent compounds‘.
According to Lewis’s concept, the number of electrons that an atom contributes for sharing in a covalent bond is called ‘covalency’. Thus, the covalency of H, O, and N In H2, O2, and N2 are 1,2 and 3 as they contribute 1,2 and 3 electrons respectively in sharing. Similarly, the covalency of C, P, and S in CH4, PCl5, and SF6 are 4.5, and 6, since they share 4,5, and 6 electrons with louring H, five Cl, and six F – atoms respectively.

Types of covalent bonds :

Covalent bonds may be either single, double, or triple depending upon the number of shared electron pairs available between the atoms in the molecules. double and triple covalent bonds are called ‘multiple covalent bonds.
Single covalent bonds are formed when two similar or different atoms share one electron each, whereas double and triple bonds are formed when they contribute 2 and 3 electrons each in sharing e.g.
(i) H2 Molecule: He+KH⟶(H(H) or H−H One electron pair (Single bond) (ii) O2 Molecule : (iii) N2 Molecule: Missing close braceMissing close brace
If sharing of electrons occurs between two like atoms, non-polar covalent bonds are formed. whereas if dissimilar atoms share their electrons the covalent bond formed is called a polar covalent bond. e.g.

one pair of electronPolar and non-polar covalent bonds :

A covalent bond is formed by equal sharing of electrons between two like atoms or atoms having nearly equal electronegativity value and has no ionic character, such a covalent bond is known as a true covalent bond’ or apolar bond. Such covalent molecule is called ‘nonpolar‘ or ‘apolar’ molecule, examples are ,H2,Cl2,Br2,O2, N2 etc. But in the case of covalent molecules formed between two dissimilar atoms having different electronegativity, the magnitude of sharing electrons is not equal in both atoms. One of them has a strong attractive force for shared electron pairs due to which it develops a partial negative charge (−δ) on it.
       As the result of the development of a partial negative charge, an equal partial positive charge (+δ) will develop on the other atom. It is noticed that these fractional charges developed on both atoms should not be considered unit charges. Molecules having partial charges are called ‘polar molecules’ such as HFHCl, NH3, H2O, etc, and bond is called a polar covalent bond. It is interesting to know that all the non-polar covalent molecules show approximately 2% ionic character. This arises due to small contributions of the ionic resonance structure of the H2 molecule. Similarly, the extreme case of ionic compounds like LiCl, NaCl, and CsCl has a small contribution to-wards covalent character.
small contribution towards covalent character.
The degree of polarity is measured in terms of dipole moment (μ)
i.e. μ=e×d debye
where
(μ) = Partial charge on the atom ‘X‘ of the order of 1010 e.s.u. and d= bond length measured in A.

Characteristics of Covalent compounds :

1. The covalent compounds are formed by sharing of electrons hence their crystal lattice of molecules is held together by weak Van der Waal forces.
2. Covalent compounds are generally insoluble in water but soluble in organic solvents like ether CCl4, alcohol, acetone, etc.
3. The covalent compounds may be solids, liquids, or gases whereas electrovalent compounds are solids.
4. They show low values of melting and boiling points as compared to ionic solids.
5. Covalent compounds are generally soft, volatile, and fusible in nature,
6. The bonds formed in covalent compounds are directional due to which they show a different type of isomerism but bonds in ionic compounds are non-directional.
7. Covalent compounds do not dissociate or conduct electricity.
8. The reaction rates of covalent compounds are much slower as compared to ionic compounds.

Octet rule:

During the formation of covalent bonds, the atoms acquire an inert gas configuration (1s2 or ns2, np6 ). It is known as the octet rule. For example, N2,O2,Cl2,CH4,NH3,H2O,HCN,CO2,  etc. obey octet rules, In these molecules, the central atom, as well as bonded atoms both acquire pure inert gas configuration, and Le. 8 electrons, are present in their valence shell.
Deviation from octet rule:
There are certain covalent molecules that do not follow the octet rule. They have either less than 8 electrons (incomplete octet) like covalent molecules that do not follow the octet ruleor more than 8 electrons (expansion of octet) like ClF3, Cl3,PCl5, SF6, and IF7as shown below :
 atom is surrounded by 6 electrons
In PCl3. a phosphorous atom is linked with five chlorine atoms by the formation of five covalent bonds hence, it is surrounded by 10 electrons (expansion of octet). CIF3 and ICl3 also have 10 electrons around the central atoms Cl and I respectively. SF6 has 12 electrons and IF7 has 14 electrons surrounding the central atoms S and I in its covalent molecules as shown below:
(iv) For PCl5 molecule : P-atom is surrounded by 10 electrons.
Reason for the failure of octet rule

Reason for the failure of the octet rule :

The following concepts are helpful to explain the failure of the octet rule.

1- Sedgwick rule of Maximum covalency:

According to Sidgwick, it is not essential for an element to have 8 electrons in its valence shell for their stability as described earlier. The covalency of an element may be expanded from the octet rule and depends on the period in which the element resides. The maximum covalency for the elements of the first period (n=1) is two, for the second period (n=2) it is 4, for 3 rd and 4th period (n=3 and 4) it is 6, and for the rest of the period (n>4) it is 8. Consequently, the maximum number of electrons shared for these elements is 1×2=2;2×4=8;2×6=12, and2×8=16 respectively.

2- Sugden’s concept of Singlet linkage:

According to this concept, it is not possible to violate the octet rule. Sugden gave a concept of singlet linkage according to which in PCl5, the three Clatoms form three covalent linkages with P-atom by sharing one electron each with them and the P atom forms two singlet linkages with the remaining two Cl-atoms by one-sided donation of an electron from it to them. This type of bond formed is called ‘singlet linkage’ and is represented by a half arrow( ).

Whose head is pointed towards the acceptor atoms (Cl), Such a type of bond is a special type of coordinate bond? In SF6, sulfur also forms four singlet linkages with four F-atoms and two covalent linkages by sharing electrons with two F-atoms as shown below.
Sugden's concept of Singlet linkage

Variable covalency :

Generally, the covalency of an element is equal to the number of unpaired s or p-electrons present in its atom in the ground state. For example, Hydrogen has one unpaired electron and hence its covalency is one. Oxygen has two unpaired p-electrons and its covalency is two and nitrogen has three unpaired p-electrons and its covalency is 3 . These are clear Irom their electronic configurations-
Variable covalency :
Since none of these atoms contain d-orbitals in their outermost shell. It is, therefore, in these atoms no excitation occurs and these atoms do not have other covalencies. On the other hand, the elements which contain empty d-orbitals in their valence shell (outermost shell) show variable covalencies. For example, the Sulphur atom has two unpaired p-electrons in its ground state hence its covalency is 2. whereas it has four and six unpaired electrons in its first and second excitations hence, its covalencies in excited states are 2 and 4 respectively. It is clear from the following electronic configurations:
electron distribution in shells

Illustration :

When sulfur combines with highly electronegative F-atoms, the energy is released during the process of combination which is sufficient to promote the paired 3px electrons to higher secant 3d−orbitals in first excitation and the paired 3 s-electrons to other 3d− orbitals in the second excitation, showing variable covalencies 4 and 6 respectively, Similarly, phosphorus also shows variable covalencies of 3 and 5 respectively. The variable covalencies 3,5 and 7 are observed in the case of halogens Cl, Br, and I (except fluorine). Thus, the maximum covalency shown by an element is equal to the number of unpaired electrons obtained after complete unpairing of paired s and p- electrons.
t is the point to remember that when an atom has a tendency to form covalent molecules with the more electronegative atom in all its covalencies, then the molecule in which the covalency of the central atom is maximum, would be more stable. For example: when sulfur combines with fluorine forms SF2, SF4, and SF6. In which the SF6 molecule is more stable than the remaining two fluorides. Similarly, PCl5 is more stable than PCl3.
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