Formal Charge Formula
Determine charge distribution in molecules
Understanding Formal Charge
Formal charge is a bookkeeping method used to track electron distribution in Lewis structures. It helps determine which resonance structure is most stable and predicts reactive sites in molecules. Unlike oxidation states which assume complete electron transfer, formal charge assumes equal sharing of bonding electrons.
The concept of formal charge was developed to evaluate different Lewis structures for the same molecule. When multiple valid structures exist (resonance), the one with formal charges closest to zero, and negative charges on more electronegative atoms, is typically most stable and best represents the actual electron distribution.
Formal charge calculations guide chemists in drawing correct Lewis structures, predicting molecular geometry with VSEPR theory, understanding reactivity patterns, and explaining why certain bonding arrangements are preferred over others in organic and inorganic molecules.
Formula
FC = V - N - B/2
- FC = formal charge on atom
- V = valence electrons in free atom (group number)
- N = non-bonding electrons (lone pair electrons)
- B = bonding electrons (shared electrons in bonds)
Alternative Form:
FC = (# valence e⁻) - (# lone pair e⁻) - (# bonds) or
FC = (# valence e⁻) - (# non-bonding e⁻) - (½ × # bonding e⁻)
Detailed Example: CO (Carbon Monoxide)
Lewis structure: :C≡O:
Step 1: Count valence electrons
Carbon (Group 14): V = 4
Oxygen (Group 16): V = 6
Step 2: Analyze Carbon atom
N = 2 electrons (one lone pair on left)
B = 6 electrons (triple bond = 3 bonds × 2 electrons)
FC(C) = 4 - 2 - 6/2 = 4 - 2 - 3 = -1
Step 3: Analyze Oxygen atom
N = 2 electrons (one lone pair on right)
B = 6 electrons (triple bond)
FC(O) = 6 - 2 - 6/2 = 6 - 2 - 3 = +1
Answer: FC(C) = -1, FC(O) = +1
Total formal charge: -1 + 1 = 0 (matches neutral molecule)
Rules for Best Lewis Structures
1. Minimize Formal Charges
Structures with formal charges closest to zero are most stable. Ideal structure has all atoms with FC = 0.
2. Negative Charge on Electronegative Atoms
When formal charges must exist, place negative charges on more electronegative atoms (O, N, F) and positive charges on less electronegative atoms.
3. Minimize Charge Separation
Structures with less charge separation (fewer and smaller formal charges) are more stable than those with large or many formal charges.
4. Adjacent Same Charges Unfavorable
Avoid structures with like charges on adjacent atoms due to electrostatic repulsion.
Common Examples
H₂O (Water)
O: FC = 6 - 4 - 4/2 = 0; H: FC = 1 - 0 - 2/2 = 0. All atoms have zero formal charge.
NH₄⁺ (Ammonium)
N: FC = 5 - 0 - 8/2 = +1; H: FC = 0 each. Total = +1 (matches ion charge).
NO₃⁻ (Nitrate)
In best resonance structure: N: FC = +1; two O: FC = 0; one O: FC = -1. Total = -1.
SO₄²⁻ (Sulfate)
S can expand octet. Best structure has S: FC ≈ 0 by forming S=O double bonds.
Important Tips & Common Mistakes
- Sum Rule: Sum of all formal charges must equal overall molecular or ionic charge.
- Valence Electrons: Use periodic table group number for main group elements (C=4, N=5, O=6, F=7).
- Counting Bonds: Count each bond stick as 2 electrons. Double bond = 4 electrons, triple = 6 electrons.
- Lone Pairs: Count all unshared electrons. Each lone pair = 2 electrons.
- Best Structure: Lowest formal charges doesn't always mean all zeros, but closest to zero with proper electroneg ativity placement.
- Expanded Octets: Period 3+ elements (P, S, Cl) can have more than 8 electrons, affecting formal charge calculations.
- Not Oxidation States: Formal charge ≠ oxidation number. Different electron allocation methods.