Formal Charge

An atom’s formal charge (F.C. or q) in the covalent view of bonding is the charge assigned to it by the covalent view of bonding, which assumes that electrons in all chemical bonds are shared equally across atoms, regardless of relative electronegativity between them. The difference between the number of valence electrons present in an atom in its neutral free state and the number of valence electrons present in that atom in its Lewis structure, to put it simply, is the formal charge. Whenever a molecule’s Lewis structure (also known as predominant resonance structure) is determined, the structure is chosen so that the formal charge on each of the atoms is as close to zero as possible.

It is possible to compute the formal charge of every individual atom in a molecule using the following equation:

F is equal to V-L-(B/2)

For a neutral atom in isolation (in its ground state), V is the number of valence electrons present; L is the number of nonbonding valence electrons present on this atom in the molecule; and B is the total number of electrons shared in bonds with other atoms present in the molecule.

Formal charge: A brief description

The difference between the number of valence electrons presents in each atom and the number of electrons with which the atom is linked is known as the formal charge of FC. For the sake of formal charge, it is assumed that any shared electrons are equally distributed across the two linked atoms.

The following equation is used to compute the formal charge:

FC = eV – eN – eB/2 

Assuming that the atom is isolated from the molecule, eV equals the number of valence electrons in the atom.

eN is the number of valence electrons on the atom in the molecule that are not bound.

eB is the number of electrons shared by the bonds that connect the atoms in the molecule to other atoms.

Calculation of a Formal Charge in an Example

Carbon dioxide, also known as CO2, is a neutral molecule with 16 valence electrons, making it a good example. For the purpose of determining formal charge, there are three possible ways to draw the Lewis structure for the molecule:

Double bonds can be formed between the carbon atom and either or both oxygen atoms (carbon = 0, oxygen = 0, formal charge = zero). One or more oxygen atoms may be in a single bond with the carbon atom and another oxygen atom may be in a double bond with the carbon atom (carbon = +1, oxygen-double = 0, oxygen-single = -1, formal charge = zero).

Single bonds can be formed between the carbon atom and each oxygen atom (carbon = +2, oxygen = -1 each, formal charge = 0) and between the carbon atom and each nitrogen atom.

While each of the possibilities results in a formal charge of 0, the first option is the best because it forecasts that there will be no charge in the molecule. The fact that it is more steady means that it is the most likely.

Formal charge used to predict the structure of molecules

The molecular structure of a molecule or an ion refers to the arrangement of atoms within the molecule or ion. The procedures for constructing Lewis structures may result in more than one conceivable molecular structure in many circumstances – many alternative bonds and lone-pair electron locations, for example, or different groupings of atoms, to name a few examples. There are a few rules of thumb that use formal charge that can be used to assist determine which of the various configurations is the most likely for a given molecule or ion:

In comparison to a molecular structure in which some formal charges are not zero, a structure in which all formal charges are zero is desirable.

If it is necessary for the Lewis structure to have non-zero formal charges, the arrangement with the smallest non-zero formal charges are the preferred configuration.

Lewis structures are preferred when the formal charges between two neighbouring charges are zero or of the opposite sign.

When deciding between various Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is the one that should be preferred.

Consider some possible carbon dioxide (CO2) structures to understand how these rules might be put into action. Although it is well known that the less electronegative atom often occupies the centre position, formal charges provide insight into why this occurs. There are three possible structures that can be imagined: carbon in the centre with two double bonds, carbon in the centre with a single and triple bond, and oxygen in the centre with double bonds.

When comparing the three formal charges, it is possible to determine that the structure on the left is preferred because it contains solely formal charges of zero.

To give another example, the thiocyanate ion, which is an ion generated by the addition of three carbon atoms to one nitrogen atom, and one sulphur atom, can have three different molecular structures: NCS–, CNS–, and CSN–. When looking at each of these molecular structures, the formal charges present can be used to determine the most likely arrangement of atoms. Each of the three possible thiocyanate ion structure types has a Lewis structure in the centre with double bonds, a Lewis structure in the centre with double bonds, and a Lewis structure in the centre with double bonds. In each situation, it is important to note that the sum of all formal charges is equal to the charge of the ion (–1). This configuration of atoms with carbon at its core, on the other hand, is favoured because it has the lowest number of atoms with non-zero formal charges of any of the possible arrangements. Additionally, it places the least electronegative atom in the centre and the negative charge on the element that is more electronegative than the least electronegative.

Conclusion

The electric charge of an atom in a molecule is referred to as the formal charge (FC).

A valence electron is defined as the number of valence electrons minus half of the number of electrons shared in a bond minus the number of electrons not bound in a molecule. A valence electron is computed as

When estimating the way electric charge is distributed in a molecule, formal charge is utilised as a guide.