Formal Charge
Formal charge is a bookkeeping tool used to assess the distribution of electrons in a Lewis structure by assigning each atom a hypothetical charge, assuming electrons in covalent bonds are shared equally. It is calculated as the number of valence electrons minus the number of lone pair electrons minus half the number of bonding electrons for each atom. The sum of formal charges equals the overall charge of the molecule or ion, and the most stable Lewis structure is generally the one with the smallest formal charges and any negative formal charges on the more electronegative atoms.
Key Formula
FC = V − L − B/2
LaTeX: FC = V - L - \frac{B}{2}
| Symbol | Meaning | Unit |
|---|---|---|
| FC | Formal charge of the atom | dimensionless |
| V | Number of valence electrons of the atom in its neutral state | electrons |
| L | Number of lone pair (non-bonding) electrons on the atom | electrons |
| B | Number of bonding (shared) electrons involving the atom | electrons |
Worked Example
Problem
Calculate the formal charge on each atom in the carbonate ion (CO₃²⁻). The central carbon is double-bonded to one oxygen and single-bonded to two oxygens; single-bonded oxygens each have 3 lone pairs, double-bonded oxygen has 2 lone pairs.
Solution
Step 1: Carbon formal charge. V(C) = 4, L(C) = 0, B(C) = 8 (2 single + 1 double = 4+4 bonds) FC(C) = 4 – 0 – 8/2 = 4 – 4 = 0 Step 2: Double-bonded oxygen formal charge. V(O) = 6, L(O) = 4 (2 lone pairs), B(O) = 4 (double bond) FC(O) = 6 – 4 – 4/2 = 6 – 4 – 2 = 0 Step 3: Single-bonded oxygen formal charge (×2). V(O) = 6, L(O) = 6 (3 lone pairs), B(O) = 2 (single bond) FC(O) = 6 – 6 – 2/2 = 6 – 6 – 1 = –1 Step 4: Check: 0 + 0 + (–1) + (–1) = –2 = overall charge of CO₃²⁻. Correct.
Answer
C: 0, double-bonded O: 0, each single-bonded O: –1. Total = –2, consistent with the carbonate ion charge. This structure shows the three equivalent resonance structures each place –1 on an oxygen.
Formal Charge Calculation Examples in Common Molecules
| Molecule/Ion | Atom | V | L | B | Formal Charge |
|---|---|---|---|---|---|
| H₂O | O | 6 | 4 | 4 | 0 |
| NH₃ | N | 5 | 2 | 6 | 0 |
| NH₄⁺ | N | 5 | 0 | 8 | +1 |
| CO₃²⁻ | C | 4 | 0 | 8 | 0 |
| CO₃²⁻ | O (single) | 6 | 6 | 2 | –1 |
| NO₂⁻ | N | 5 | 2 | 6 | 0 |
Interactive Tools
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Related Terms
Resonance Structure
Resonance structures (also called canonical forms or contributing structures) are two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms), used collectively to represent the actual electron distribution. No single resonance structure accurately depicts the molecule; the true structure is a hybrid (weighted average) of all contributing structures, with electron density delocalised across multiple bonds. The concept is essential for understanding the stability of molecules like benzene, ozone, and carbonate ion, where observed bond lengths are intermediate between single and double bonds.
Covalent Bond
A covalent bond is a type of chemical bond formed when two atoms share one or more pairs of electrons, resulting in a stable arrangement for both atoms. This sharing occurs most commonly between non-metal atoms that have similar electronegativities, allowing each atom to achieve a full valence shell without complete electron transfer. Covalent bonds are the foundation of organic chemistry and molecular biology, governing the structure of molecules ranging from water (H₂O) to complex proteins.
Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a shared pair of electrons toward itself within a covalent bond, expressed on a dimensionless scale. The most widely used scale is the Pauling scale, where fluorine is assigned the highest value of 3.98, making it the most electronegative element, and caesium the lowest at 0.79. Electronegativity determines bond polarity, the character of chemical bonds (ionic vs. covalent), and influences molecular properties such as reactivity, acid strength, and solubility.
From Latin "formalis" (relating to form or structure) + Latin "carricare" (to load, charge). The term "formal" emphasises that this charge is a conceptual assignment for bookkeeping purposes, not a true physical charge. The concept was systematised alongside Lewis structure theory in the early-to-mid 20th century.