Q.Calculate the format charge on each atom of CO3^2- ion.

Why Formal Charge Exists

Imagine you're drawing a Lewis structure. You've placed atoms, connected them with bonds, and distributed electrons. But how do you know which arrangement is better? That's where formal charge comes in.

The core idea is simple: formal charge tells you whether an atom in a molecule has the same number of electrons it "owns" as it would as a free atom. If an atom ends up with more or fewer electrons than it started with, it carries a charge.

Think of it like a budget. A neutral atom starts with a certain number of valence electrons (its "allowance"). When it bonds, it shares some electrons and keeps others to itself. Formal charge is the balance — the difference between what it started with and what it ends up with.

The Precise Definition

Formal charge is defined as:

The key insight: bonding electrons are split equally between the two atoms sharing them. That's why we take half of $B$. Lone-pair electrons belong entirely to the atom, so we subtract all of $L$.

Worked Example: Carbon Dioxide ($\text{CO}_2$)

Let's calculate formal charges for one possible Lewis structure of $\text{CO}_2$: $\text{O}=\text{C}=\text{O}$.

Central carbon:

• $V = 4$ (carbon has 4 valence electrons)
• $L = 0$ (carbon has no lone pairs in this structure)
• $B = 8$ (carbon shares 4 bonds × 2 electrons each = 8 bonding electrons)

Each oxygen:

• $V = 6$ (oxygen has 6 valence electrons)
• $L = 4$ (each oxygen has 2 lone pairs = 4 electrons)
• $B = 4$ (each oxygen shares a double bond = 4 bonding electrons)

All formal charges are zero — this is a very stable structure.

Why We Use Formal Charge

Formal charge helps you choose between multiple possible Lewis structures. The rules are:

1. The sum of formal charges must equal the overall charge of the molecule or ion.
2. Structures with smaller formal charges are preferred (closer to zero).
3. Negative formal charges should be on more electronegative atoms when possible.

Common Mistake to Avoid …