Chemistry Covalent Bond, Formal Charge and Limitations of Octet Rule

Topics Covered :

● Covalent Bond
● Formal Charge
● Limitations of Octet Rule
● Other Drawbacks of the Octet Theory

Covalent Bond :

`=>` Langmuir (1919) refined the Lewis postulations by abandoning the idea of the stationary cubical arrangement of the octet, and by introducing the term covalent bond.
`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
The bond formed between the two atoms by mutual sharing of electrons between them so as to complete their octets or duplets in case of elements having only one shell is called `color{red}("covalent bond")` and the number of electrons contributed by each atom is known as `color{red}("covalency")`.

`=>` The Lewis-Langmuir theory can be understood by considering the formation of the chlorine molecule,`Cl_2`.

● The `Cl` atom with electronic configuration, `[Ne]3s^2 3p^5`, is one electron short of the argon configuration.

● The formation of the `Cl_2` molecule can be understood in terms of the sharing of a pair of electrons between the two chlorine atoms, each chlorine atom contributing one electron to the shared pair.

● In the process both chlorine atoms attain the outer shell octet of the nearest noble gas (i.e., argon).

● The dots represent electrons. Such structures are referred to as `text(Lewis dot structures)`.

● The Lewis dot structures can be written for other molecules also, in which the combining atoms may be identical or different.

`=>` The important conditions being that :

● Each bond is formed as a result of sharing of an electron pair between the atoms.

● Each combining atom contributes at least one electron to the shared pair.

● The combining atoms attain the outershell noble gas configurations as a result of the sharing of electrons.

● Thus in water and carbon tetrachloride molecules, formation of covalent bonds can be represented as shown in fig.

`=>` Wen two atoms share one electron pair they are said to be joined by a single covalent bond.

`=>` In many compounds we have multiple bonds between atoms.

● The formation of multiple bonds envisages sharing of more than one electron pair between two atoms.

`=>` If two atoms share two pairs of electrons, the covalent bond between them is called a `color{red}("double bond")`.

● `color{red}("Example :")` (i) In the carbon dioxide molecule, we have two double bonds between the carbon and oxygen atoms.

(ii) Similarly in ethene molecule the two carbon atoms are joined by a double bond. See fig..

`=>` When combining atoms share three electron pairs as in the case of two nitrogen atoms in the `N_2` molecule and the two carbon atoms in the ethyne molecule, a triple bond is formed. See fig
`color{purple}♣ color{Violet} " Just for Curious"`
`=>` The covalent compounds exist in all three states:solid,liquid,gas.
`=>` They consist of molecules.
`=>` They have low melting and boiling point.(held by weak vander waals forces ).
`=>` They are generally soluble in organic solvents but insoluble in water and other polar solvents.
`=>` They are bad conductor of electricity.
`=>` They are directional in nature.

Lewis Representation of Simple Molecules (the Lewis Structures) :

`=>` The Lewis dot structures provide a picture of bonding in molecules and ions in terms of the shared pairs of electrons and the octet rule.

`=>` It may not explain the bonding and behaviour of a molecule completely, it does help in understanding the formation and properties of a molecule to a large extent.

`=>` Writing of Lewis dot structures of molecules is, therefore, very useful.

`=>` The Lewis dot structures can be written by adopting the following steps :

(i) The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms.

`color{red}("Example :")` In the `CH_4` molecule there are eight valence electrons available for bonding (`4` from carbon and `4` from the four hydrogen atoms).

(ii) For anions, each negative charge would mean addition of one electron.

● For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons.

`color{red}("Example :")` (a) For the `CO_3^(2–)` ion, the two negative charges indicate that there are two additional electrons than those provided by the neutral atoms.

(b) For `NH_(4)^+` ion, one positive charge indicates the loss of one electron from the group of neutral atoms.

(iii) Knowing the chemical symbols of the combining atoms and having knowledge of the skeletal structure of the compound (known or guessed intelligently), it is easy to distribute the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.

(iv) In general the least electronegative atom occupies the central position in the molecule/ion.

`color{red}("Example ":)` In the `NF_3` and `CO_3^(2–)`, nitrogen and carbon are the central atoms whereas fluorine and oxygen occupy the terminal positions.

(v) After accounting for the shared pairs of electrons for single bonds, the remaining electron pairs are either utilized for multiple bonding or remain as the lone pairs.

● The basic requirement being that each bonded atom gets an octet of electrons.

`=>` Lewis representations of a few molecules/ ions are given in Table 4.1.

Formal Charge :

`=>` Lewis dot structures, in general, do not represent the actual shapes of the molecules.

`=>` In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by a particular atom.

● It is, however, feasible to assign a formal charge on each atom.

`color{green}("Formal Charge :")` The formal charge of an atom in a polyatomic molecule or ion may be defined as the difference between the number of valence electrons of that atom in an isolated or free state and the number of electrons assigned to that atom in the Lewis structure.

It is expressed as :

Formal charge (F.C.) on an atom in a Lewis structure `color{green}(= [text(total number of valence electrons in the free atom)] - [text{total number of non bonding (lone pair) electrons}] - (1//2) [text{total number of bonding(shared) electrons}])`

`=>` The counting is based on the assumption that the atom in the molecule owns one electron of each shared pair and both the electrons of a lone pair.

`=>` Let us consider the ozone molecule `(O_3)`. The Lewis structure of `O_3` may be drawn as shown in fig.1.

The atoms have been numbered as `1`, `2` and `3`.

● The central `O` atom marked 1

` = 6-2-1/2 (6) = +1`

● The end `O` atom marked 2

`= 6-4-1/2 (4) = 0`

● The end `O` atom marked 3

`= 6-6-1/2 (2) = -1`

● Hence, we represent `O_3` along with the formal charges as shown in fig.2.

`color{red}("Note :")` (i) Formal charges do not indicate real charge separation within the molecule.

(ii) Indicating the charges on the atoms in the Lewis structure only helps in keeping track of the valence electrons in the molecule.

`=>` Formal charges help in the selection of the lowest energy structure from a number of possible Lewis structures for a given species.

`=>` Generally the lowest energy structure is the one with the smallest formal charges on the atoms.

`=>` The formal charge is a factor based on a pure covalent view of bonding in which electron pairs are shared equally by neighbouring atoms.

Limitations of the Octet Rule :

`=>` The octet rule, though useful, is not universal.

`=>` It is quite useful for understanding the structures of most of the organic compounds and it applies mainly to the second period elements of the periodic table.

`=>` There are three types of exceptions to the octet rule :

`color{green}("The Incomplete Octet of the Central Atom :")`

● In some compounds, the number of electrons surrounding the central atom is less than eight.

● This is especially the case with elements having less than four valence electrons. Examples are `LiCl, BeH_2` and `BCl_3`.

`Li : Cl \ \ \ \ \ \ \ H : Be : H \ \ \ \ \ \ \ Cl : overset( overset(Cl)( . .))B : Cl`

● `Li`, `Be` and `B` have `1`, `2` and `3` valence electrons only.

● Some other such compounds are `AlCl_3` and `BF_3`

`color{green}("Odd-electron Molecules :")`

● In molecules with an odd number of electrons like nitric oxide, `NO` and nitrogen dioxide, `NO_2`, the octet rule is not satisfied for all the atoms.

`underset(.)overset(. .)N = underset(. .) overset(. .)O \ \ \ \ \ underset(. .) overset(. .)O= overset(.+)N - underset(. .) overset( . . -)O :`

`color{green}("The Expanded Octet :" )`

● Elements in and beyond the third period of the periodic table have, apart from `3s` and `3p` orbitals, `3d` orbitals also available for bonding.

● In a number of compounds of these elements there are more than eight valence electrons around the central atom. This is termed as the expanded octet. Obviously the octet rule does not apply in such cases.

● Some of the examples of such compounds are : `PF_5`, `SF_6`, `H_2SO_4` and a number of coordination compounds.

`color{red}("Note :")` Sulphur also forms many compounds in which the octet rule is obeyed. In sulphur dichloride, the `S` atom has an octet of electrons around it.

`: underset(. .) overset(. .)(Cl) - underset(. .) overset(. .)S - underset(. .) overset(. .)(Cl):` or `: underset(. .) overset(. .)(Cl) : underset(. .)overset(. .)S : underset(. .)overset(. .)(Cl) :`

Other drawbacks of the octet theory :

`=>` It is clear that octet rule is based upon the chemical inertness of noble gases.

● However, some noble gases (for example xenon and krypton) also combine with oxygen and fluorine to form a number of compounds like `XeF_2, KrF_2, XeOF_2` etc.

`=>` This theory does not account for the shape of molecules.

`=>` It does not explain the relative stability of the molecules being totally silent about the energy of a molecule.