Octet Rule

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Octet Rule

We have discussed in chemical bonding that atoms combine together to form molecules which are more stable. During the formation of a molecule, atoms try to attain their nearest noble gas configuration. In most cases, they need to have eight electrons in their outermost shell. This property of atoms is known as octet rule.

Atoms can combine in the following ways to have octet in their outermost shell :

  1. By transfer of valence shell electrons.
  2. By sharing of electrons.

Transfer of electrons

A bond between two atoms can be formed by transfer of valence electrons between them. This type of bond is known as electrovalent bond or ionic bond.
Example : Formation of sodium chloride (NaCl) which is discussed below :

  • octet rule, transfer of electrons exmple : NaClThe atomic number of sodium (Na) is 11, so its electronic configuration is 2,8,1. Sodium has only one electron in its outermost shell. Sodium loses one electron to attain its nearest noble gas configuration (neon).

    Na → Na+ + e−

  • The atomic number of chlorine (Cl) is 17, so its electronic configuration is 2,7. Clearly, chlorine requires one electron to attain its nearest noble gas configuration (argon).

    Cl + e− → Cl−

  • The complete reaction may be written as :

    Na+ + Cl− → NaCl or Na+Cl−

Sharing of electrons

The bond formed by sharing of electrons is called a covalent bond. Sharing of electrons can be understood using the example of Cl2.

Formation of dichlorine (Cl2) : The atomic number of chlorine is 17 and has 7 valence electrons, which means Cl needs one more electron to complete its octet. In order to complete its octet , each chlorine atom contributes one electron to the shared pair. By doing this, both the chlorine atoms attain their nearest noble gas configuration (which is argon).

Sharing of electrons between two Cl atoms

Octet rule example : Chlorine atom + Chlorine atom = Dichlorine Molecule (Cl2)
The structures where dots are used to represent valence electrons are known as Lewis dot structures. The above structure is an example of Lewis dot structures.

Single covalent bond : When one pair of electrons takes part in bond formation, the bond is known as single covalent bond.
Example : In H2O molecule, each hydrogen-oxygen bond is formed by sharing a pair of electrons.

Sharing of electrons between Hydrogen and oxygen atoms in H2O molecule

formation of H2O molecule : octet rule
Hydrogen completes its duplet instead of octet to attain its nearest noble gas configuration (i.e., helium) because the outermost shell of hydrogen atom contains only the s orbital which can contain up to 2 electrons.

Double covalent bond : When two pairs of electrons are involved in bond formation, the bond is known as double covalent bond.
Example : In CO2 molecule, each carbon-oxygen bond is formed by sharing two pairs of electrons.

Sharing of electrons between Carbon and oxygen atoms in CO2 molecule

Octet rule, lewis dot structure of CO2 molecule

Triple covalent bond : A triple covalent bond is formed when three pair of electrons are shared.
Example : In N2 molecule, three pair of electrons are shared between nitrogen atoms.

Sharing of electrons between two Nitrogen atoms in N2

Octet rule triple bond example : N2 molecule lewis dot structure

Dashes −, = and ≡

For simplicity, each shared pair of electrons can be represented by a dash (−) i.e., one dash is used to represent a single bond, two dashes are used for double, and three for triple bond. Example : the above structure of N2 can be simplified as :N≡N:.

Steps involved for writing Lewis structures

  1. Count the total number of valence electrons of combining atoms. In case of ions, add one electron for each -ve charge and subtract one electron for each +ve charge.
  2. In general, the least electronegative atom occupies the central position while other atoms are placed at the terminal.
  3. Draw a single bond between atoms and check if all the atoms have complete octets. If each atom has complete octet then our structure is complete.
  4. If some atoms have incomplete octets then use multiple bonds until they have complete octets. (Note : the number of dots i.e., valence electrons must not change)

Question : Draw the Lewis structure of nitrite (NO2−) ion.

Answer : The steps are given below:

1. The total number of valence electrons in NO2− ion (nitirite ion) is 18 ( 5 in Nitrogen atom + 2x6 = 12 in oxygen atoms + 1 for negative charge).

Total number of valence electrons in NO2− ion = 18.

2. In NO2− ion, N being the least electronegative will be in the central.

O    N    O

3. Draw single bonds between nitrogen and oxygen atoms.

Lewis structure example : nitrite ion with incomplete octet of nitrogen

In the above structure, both the oxygen atoms have complete octets; the nitrogen atom has less than eight electrons. That means we are not done yet.

4. Draw multiple bonds in such a way that all atoms have complete octets.

All atoms have complete octets now

Lewis structure example : nitrite ion with complete octet

The above structure can be simplified as :

simplified Lewis structure example 1 : nitrite ion

or

simplified Lewis structure example 2 : nitrite ion
More Practice Questions
In Lewis dot structures, the net charge is possessed by the ion as a whole and not by a particular atom.

Formal Charge

Lewis dot structures do not tell us about the charge possessed by individual atoms but sometimes it is useful to assign a formal charge to each atom in a molecule or ion.

The formal charge of an atom in a molecule or ion is defined as the difference between the number of valence electrons in free state and the number of electrons assigned to that atom in the Lewis structure. i.e.,

Cf = Nv − Ue − ½(Bn)

Where,

Cf = Formal charge on an atom in a molecule/ion
Nv = Total number of valence electrons in free state
Ue = Total number of non bonding electrons (lone pairs)
Bn = Total number of bonding electrons (shared)

The counting of valence electrons in the Lewis structure is based on the assumption that the atom in the molecule/ion own one electron of each shared pair and both the electrons of a lone pair.

Let us take an example of ozone (O3)

Formal charge on ozone molecule

The atoms have been marked as a, b, c.

The formal charge on the atom marked 'a' = 6 - 2 -1/2(6) = +1

The formal charge on the atom marked 'b' = 6 - 4 -1/2(4) = 0

The formal charge on the atom marked 'c' = 6 - 6 -1/2(2) = -1

Hence, O3 along with the formal charges can be represented as :

Formal charge on ozone molecule

Exceptions to the Octet Rule

  1. Formation of compounds like BF3, AlCl3 etc.

    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.

    Limitations of the octet rules example : BF3

    B has 6 surrounding electrons in BCl3.

    Limitations of the octet rules example : AlCl3

    Al has 6 surrounding electrons in AlCl3.

    H : Be : H

    Be has 4 surrounding electrons in BeH2.

    Some other examples are BeCl2, LiCl etc.

  2. Odd electrons molecules

    Compounds having odd number of electrons such as NO, NO2 do not obey the octet rule.

  3. The expanded octet

    Elements beyond the third period of the periodic table also have d orbitals. In a number of such compounds, more than eight valence electrons have been seen.

    Some of the examples of such compounds are : PF5, SF6, H2SO4 etc.

  4. Formation of noble gas compounds

    Octet rule is based on the chemical inertness of the noble gases. However, some noble gases combine with other atoms to form a number of molecules. e.g. XeF2, KrF2, XeF4, XeF6 etc.

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