Chemistry · General chemistry I · Worked example
Use formal charge to choose a structure
Two Lewis structures can be drawn for CO₂: one with two double bonds and two lone pairs on each oxygen, and one with a triple bond and a single bond, where the triple-bonded oxygen has one lone pair and the other has three. Which is better?
Check the electron count
Carbon brings 4 and each oxygen 6, so both structures must use 16 electrons. Each does: four bonds hold 8 and the lone pairs hold 8.
Formal charges with two double bonds
Each oxygen: 6 − 4 − 4/2 = 0. Carbon: 4 − 0 − 8/2 = 0. Every formal charge is zero.
Formal charges with a triple bond
The triple-bonded oxygen: 6 − 2 − 3 = +1. Carbon: 4 − 0 − 4 = 0. The single-bonded oxygen: 6 − 6 − 1 = −1. They add to zero, but oxygen, a very electronegative atom, carries a positive formal charge.
Choose
Prefer the structure whose formal charges are closest to zero: the one with two double bonds.
Result
The structure with two double bonds, where every formal charge is zero.
Your turn
In the cyanate ion, OCN⁻, oxygen is single-bonded to carbon, which is triple-bonded to nitrogen. Oxygen has three lone pairs and nitrogen one. Find each formal charge.
Show the answer and explanation
Oxygen −1, carbon 0, nitrogen 0.
Oxygen: 6 − 6 − 1 = −1. Carbon: 4 − 0 − 4 = 0. Nitrogen: 5 − 2 − 3 = 0. They add to −1, the charge of the ion, and the negative charge sits on the most electronegative atom.
Keep exploring
In Lewis electrons & resonance, make one bond triple and the other single, moving lone pairs so the total stays 16: the count still matches, and the formal charges come out +1, 0 and −1.
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