Chemical Bonding
Chemical Bond: The attractive force which holds together different constituents in a chemical species is called chemical bond.
Kossel Lewis Approach to Chemical Bonding
Lewis proposed that atom is a positively charged ‘Kernel’ (which contains the nucleus and inner electrons) and the outer shell which could accommodate a maximum of eight electrons. He also said that the Kernel is surrounded by a cube, and eight electrons occupy the eight corners of the cube. Lewis postulated that when atoms are linked by chemical bonds, they achieve the stable octet.
Example: The outer shell electron of sodium will occupy one corner of the cube, while the outer shell electrons of a noble gas will occupy all the eight corners of the cube.
Bond can be formed either by transfer of electrons or by sharing of electrons. Once bonds are formed, each atom attains a sable outer octet of electrons.
Lewis Symbols
The electrons in the outer shell are called valence electrons because they are the electrons which take part in bond formation. The electrons in inner shell generally do not take part in bond formation because they are well protected. Lewis proposed simple notations to represent valence electrons in an atom. These notations are called Lewis symbol.
Significance of Lewis Symbols: Number of valence electrons is represented by number of dots around the symbol of atom. Number of valence electrons gives an idea about the common or group valence of an element. Group valence is generally either equal to the number of dots in Lewis symbol or 8 minus the number of dots.
Kossel proposed following ideas about chemical bonding:
- In the periodic table, the highly electronegative halogens and the highly electropositive alkali metals are separated by noble gases.
- Formation of a negative ion from a halogen atom, and that of a positive ion by an alkali metal atom happens because of gain and loss of an electron by the respective atom.
- Thus, the negative and positive ions attain noble gas electronic configuration.
- The negative and positive ions are stabilized by electrostatic attraction.
Example: Formation of NaCl illustrates Kossel’s theory as follows:
Na → Na+ + e-
[Ne] 3s1 [Ne]
Cl + e-1 → Cl-
[Ne] 3s2 3p5 [Ne] 3s2 3p6 or [Ar]
Na+ + Cl- → NaCl or Na+Cl-
Note: Here Ne stands for Neon
The bond formed as a result of the electrostatic attraction between the positive and negative ions is called electrovalent bond. Thus, electrovalence is equal to the number of unit charge(s) on the ion.
Octet Rule: (Kossel and Lewis 1916): Atoms can combine either by transfer of valence electrons or by sharing of valence electrons in order to attain an octet in their valence shells.
Covalent Bond: When chemical bond is formed by sharing of electrons, it is called covalent bond. The sharing of a pair of electrons results in the formation of a single bond, sharing of two pairs of electrons results in double bond and sharing of three pairs of electrons results in a triple bond. Following are some examples.
Lewis Dot Structure
- The total number of electrons required for writing the structure is obtained by adding the valence electrons of the combining atoms.
- For anions, each negative charge means addition of one electron from the total number of valence electrons. For cations, each positive charge means subtraction of one electron from the total number of valence electrons.
- Generally, the least electronegative atom occupies the central position in a molecule or ion.
- After accounting for the shared pairs of electrons for single bonds, remaining pairs of electrons are either utilized for multiple bonding or are left as lone pairs. While doing so, it is ensured that each bonded atom gets an octet of electrons.
Formal Charge
In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom. But, it is feasible to assign a formal charge to each atom. The formal charge of an atom in a polyatomic molecule or ion is defined as the difference between the number of valence electrons of that atom in free-state and the number of electrons assigned to that atom in Lewis structure.
Formal Charge (F.C.) = Total no. of valence electrons in free atom – Total no. of non-bonding (lone pair) electrons – ½ (total number of bonding (shared) electrons
Example: Let us take example of ozone molecule (O3), for which Lewis structure is given below, in which atoms are numbered as 1, 2 and 3.
Formal charge on central O atom (1)
`= 6-2-1/2(6) = +1`
Formal charge on O atom (2)
`=6-4-1/2(4) = 0`
Formal charge on O atom (3)
`=6-6-1/2(2)=-1`
So, O3 along with formal charges can be represented as follows:
Formal charges don not indicate real charge separation within the molecule. Indicating the formal charge only helps in keeping track of the valence electrons in the molecule. It helps in selection of the lowest energy structure from a number of possible Lewis structures for a given species. The lowest energy structure is generally the structure with the smallest formal charges on atoms. Formal charge is a factor bases on a pure covalent view of bonding.
Limitations of Octet Rule
Incomplete octet of the central atom: In some compounds, the number of electrons surrounding the central atom is less than eight. It is especially the case with elements having less than four valence electrons. Examples: LiCl, BeH2 and BCl3
Odd-electron Molecules: Octet rule is not satisfied in molecules with odd number of electrons. Exmaples: NO and NO2
Expanded Octet: Elements in and beyond the third period have 3d orbitals also available for bonding (apart from 3s and 3p orbitals). In a number of compounds of these elements, there are more thatn eight valence electrons around the central atom. Such a condition is called expanded octet. Examples: PF5, SF6, H2SO4, etc.
Other drawbacks of octet rule
- We know that the octet rule is based on the chemical inertness of noble gases. But some noble gases do combine with oxygen and fluorine to form a number of compounds, e.g. XeF2, KrF2, XeOF2, etc.
- This theory does not explain the shape of molecules.
- It does not explain the relative stability of molecules because it does not explain the energy of a molecule.