Formal Charge Calculator

What is formal charge?

Formal charge is a bookkeeping tool chemists use to assign a hypothetical charge to each atom in a Lewis structure, assuming that all bonding electrons are shared equally between the two atoms in a bond.

It does not represent the true charge on an atom (which is influenced by electronegativity and bonding), but it is extremely useful for:

• Checking whether a Lewis structure is reasonable
• Comparing alternative Lewis structures
• Choosing the most important resonance structure(s)
• Tracking where positive and negative charge is located in a mechanism

How to calculate formal charge step‑by‑step

1. Draw a complete Lewis structure. Include all lone pairs and all bonds (single, double, triple).
2. Pick an atom and count electrons:
   • Valence electrons, V: from the periodic table (e.g. C = 4, N = 5, O = 6, F = 7).
   • Nonbonding electrons, N_nonbonding: all lone‑pair and unshared electrons on that atom.
   • Bonding electrons, N_bonding: total electrons in bonds around that atom (2 per single bond, 4 per double, 6 per triple).
3. Apply the formula. \[ \text{FC} = V - N_\text{nonbonding} - \frac{1}{2}N_\text{bonding} \]
4. Repeat for every atom. The sum of all formal charges must equal the overall charge of the molecule or ion.

Worked example 1: nitrate ion, NO₃⁻

One common resonance structure of nitrate has:

• Central N with three O atoms
• One N=O double bond, two N–O single bonds
• Each single‑bonded O has three lone pairs; the double‑bonded O has two lone pairs

Nitrogen (central atom)

• \(V = 5\) (group 15)
• \(N_\text{nonbonding} = 0\) (no lone pairs on N)
• \(N_\text{bonding} = 8\) (three N–O bonds: 2 + 2 + 4 electrons)

\[ \text{FC}_\text{N} = 5 - 0 - \frac{8}{2} = 5 - 4 = +1 \]

Double‑bonded oxygen (N=O)

• \(V = 6\)
• \(N_\text{nonbonding} = 4\) (two lone pairs)
• \(N_\text{bonding} = 4\) (one double bond)

\[ \text{FC}_\text{O(double)} = 6 - 4 - \frac{4}{2} = 6 - 4 - 2 = 0 \]

Each single‑bonded oxygen (N–O)

• \(V = 6\)
• \(N_\text{nonbonding} = 6\) (three lone pairs)
• \(N_\text{bonding} = 2\) (one single bond)

\[ \text{FC}_\text{O(single)} = 6 - 6 - \frac{2}{2} = 6 - 6 - 1 = -1 \]

Check the total:

\[ (+1) + 0 + (-1) + (-1) = -1 \] which matches the overall charge of NO₃⁻.

Worked example 2: hydronium ion, H₃O⁺

Hydronium has O in the center with three O–H single bonds and one lone pair on O.

Oxygen

• \(V = 6\)
• \(N_\text{nonbonding} = 2\) (one lone pair)
• \(N_\text{bonding} = 6\) (three O–H single bonds)

\[ \text{FC}_\text{O} = 6 - 2 - \frac{6}{2} = 6 - 2 - 3 = +1 \]

Each hydrogen

• \(V = 1\)
• \(N_\text{nonbonding} = 0\)
• \(N_\text{bonding} = 2\) (one O–H bond)

\[ \text{FC}_\text{H} = 1 - 0 - \frac{2}{2} = 1 - 1 = 0 \]

Total: \(+1 + 0 + 0 + 0 = +1\), as expected for H₃O⁺.

How to choose the best Lewis structure using formal charge

When you have multiple possible Lewis structures, use these guidelines:

1. Minimize formal charges. Structures with all atoms at 0, or with the smallest magnitude of formal charges, are usually preferred.
2. Put negative charge on more electronegative atoms. For example, O and F are better places for −1 than C or B.
3. Put positive charge on less electronegative atoms. Positive charge on C is usually more reasonable than on O or F.
4. Keep the total formal charge equal to the overall charge. If the species is neutral, the sum must be 0; if it is an ion, the sum must equal the ionic charge.

Common mistakes with formal charge

• Forgetting lone pairs. Missing lone pairs changes both nonbonding and bonding counts and leads to wrong formal charges.
• Confusing formal charge with oxidation state. Formal charge assumes equal sharing; oxidation state gives all bonding electrons to the more electronegative atom.
• Ignoring resonance. In many ions (e.g. NO₃⁻, CO₃²⁻, benzene), the “real” structure is a resonance hybrid. Formal charges may be spread out over several atoms.
• Assuming formal charge is the actual charge density. Real electron density is better described by resonance and molecular orbital theory; formal charge is just a simple model.

FAQ

What is a good formal charge pattern?

A “good” Lewis structure typically has all atoms with formal charges of 0, or at most ±1, and places any negative charge on the most electronegative atom. Large charge separations like +2 and −2 on adjacent atoms are usually a sign that a better structure exists.

Do I always need to draw all resonance structures?

For simple problems, you often only need the major resonance contributors (those with the best formal charge patterns). However, when predicting reactivity or charge delocalization, it is helpful to draw all reasonable resonance forms that obey the octet rule and match the overall charge.

Can formal charge explain reactivity?

Often, yes. Atoms with negative formal charge tend to be nucleophilic or basic; atoms with positive formal charge tend to be electrophilic or acidic. But always consider resonance and inductive effects as well.