Formal charge is a fundamental tool used to predict a molecule’s chemical behavior and determine the most stable Lewis structure. It is a theoretical concept, acting as an essential bookkeeping method for valence electrons, and does not represent the actual charge on an atom. This article explains how to calculate formal charge and use these calculations to assess the stability of different resonance structures.
Defining Formal Charge
Formal charge (FC) is the hypothetical charge an atom would possess if all bonding electrons were shared equally, ignoring electronegativity differences. This calculation is crucial when drawing Lewis structures, particularly when multiple valid resonance forms exist. The goal is to minimize formal charges across the structure; a structure where all atoms have a formal charge of zero is generally the most stable.
The Formal Charge Formula
The calculation of formal charge is straightforward, requiring knowledge of the atom’s valence electrons and the electron assignment in the Lewis structure. The formula for calculating formal charge (FC) on a specific atom is:
FC = V – N – B/2
V represents the number of valence electrons the atom brings to the molecule. N is the number of non-bonding electrons (lone pairs). B is the total number of bonding electrons. The sum of the formal charges for all atoms in a neutral molecule must equal zero. For a polyatomic ion, the sum must equal the overall charge of the ion, which serves as a useful check.
Step-by-Step Calculation of Formal Charge
To calculate the formal charge for an atom within a molecule, follow these steps:
- Determine the total number of valence electrons for all atoms in the molecule or ion. Adjust this total by adding electrons for anions or subtracting electrons for cations.
- Draw the Lewis structure for the molecule or ion. Place the least electronegative atom in the center, connect atoms with single bonds, and distribute the remaining electrons as lone pairs to satisfy the octet rule.
- Count the number of non-bonding electrons (N) and the number of bonding electrons (B) for the specific atom being analyzed.
- Apply the formal charge formula (FC = V – N – B/2) to each atom individually.
Assessing Structure Stability Using Formal Charge
Formal charge is the primary tool used to evaluate the relative stability of different resonance structures. When comparing multiple valid structures, follow these three rules in order of priority to determine the most stable structure:
Rule 1: Minimize Formal Charges
The most stable structure is the one that minimizes the magnitude of formal charges on all atoms. Ideally, all atoms should have a formal charge of zero. Structures with formal charges of +1, -1, or higher are less stable.
Rule 2: Place Negative Charges on the Most Electronegative Atom
If formal charges cannot be avoided, the most stable structure places the negative formal charge on the most electronegative atom. Electronegative atoms (such as Oxygen or Fluorine) are better able to accommodate and stabilize a negative charge. Conversely, positive formal charges should be placed on the least electronegative atoms.
Rule 3: Avoid Adjacent Charges of the Same Sign
Structures that have adjacent atoms bearing formal charges of the same sign (e.g., +1 next to +1) are highly unstable. This electrostatic repulsion significantly destabilizes the molecule and should be avoided.
Example: Carbon Dioxide (CO2)
Consider the molecule carbon dioxide (CO2). Carbon has 4 valence electrons, and Oxygen has 6. The total valence electrons equal 16.
Structure A (Standard): O=C=O
Carbon: FC = 4 – 0 – 8/2 = 0.
Oxygen (left): FC = 6 – 4 – 4/2 = 0.
Oxygen (right): FC = 6 – 4 – 4/2 = 0.
Total Formal Charge = 0.
Structure B (Less Common): O≡C-O
Carbon: FC = 4 – 0 – 8/2 = 0.
Oxygen (triple bond): FC = 6 – 2 – 6/2 = +1.
Oxygen (single bond): FC = 6 – 6 – 2/2 = -1.
Total Formal Charge = 0.
Comparing Structure A and Structure B, Structure A is the most stable because all atoms have a formal charge of zero (Rule 1). Although Structure B satisfies the octet rule, the presence of non-zero formal charges makes it less favorable.
Limitations of Formal Charge
While formal charge is an excellent tool for predicting the most likely Lewis structure, it has limitations. It assumes perfect covalent sharing, ignoring the reality of electronegativity differences. For example, in water (H2O), the formal charges are zero, but the oxygen atom is significantly more negative than the hydrogen atoms due to its high electronegativity, resulting in a polar molecule. Formal charge is primarily an electron bookkeeping model, not a measure of actual charge distribution.