Source: Chapter 1.4 Libre Text Notes, Organic Chemistry (Purdue University)
Difficulty: Introductory | Prerequisites: Lewis structures (Ch. 1.3), valence electrons, the octet rule
Formal charge sits right at the start of organic chemistry because it governs how you read and draw Lewis structures for the rest of the course. If you can assign formal charges correctly, you can predict which atom in a molecule is the electron donor and which is the electron acceptor, and that directly feeds into reaction mechanisms from Chapter 2 onward. You should already be comfortable drawing Lewis structures and counting valence electrons before tackling this section. Think of formal charge as the bookkeeping system that tells you whether an atom "owns" the right number of electrons once it is inside a molecule.
Formal charge is the difference between how many valence electrons an atom has on the periodic table and how many it "entirely owns" inside a molecule (lone-pair electrons count fully, bonding electrons count as half). A positive formal charge means the atom owns fewer electrons than expected and will tend to attract electrons; a negative formal charge means it owns more and will tend to donate them.
Formal charge
The hypothetical charge assigned to an atom in a molecule by assuming all bonding electrons are shared equally between atoms. It equals the number of valence electrons on the free atom minus the number of electrons "entirely owned" in the molecule.
In simple terms, it is a way of keeping score: does the atom in this molecule have more or fewer electrons than it would if it were sitting alone on the periodic table?
Valence electrons
The electrons in the outermost shell of a free (unbonded) atom, as read directly from the periodic table group number (for main-group elements).
Think of these as the electrons an atom "should" have. They are your starting point for the formal charge calculation.
Lone pairs (non-bonding electrons)
Pairs of valence electrons on an atom that are not shared with any other atom. In a Lewis structure they appear as dots sitting on one atom.
These belong entirely to that atom, so they count fully when you tally up "entirely owned" electrons.
Bonding electrons
Electrons shared between two atoms in a covalent bond (each line in a Lewis structure represents two bonding electrons).
Because these are shared, each atom only "owns" half of them for the purpose of formal charge.
Lewis structure
A diagram showing how valence electrons are distributed among atoms in a molecule, using lines for bonds and dots for lone pairs.
This is the map you read when assigning formal charges. If you cannot draw a correct Lewis structure first, the formal charge calculation will not come out right.
Formal charge helps predict how a molecule behaves in reactions.
Atoms with a positive formal charge tend to attract electrons (they are electrophilic sites).
Atoms with a negative formal charge tend to donate electrons (they are nucleophilic sites).
Formal charge is often present when an atom does not have its usual number of bonds.
Valence rules act as flags: if an atom has an unusual number of bonds, check for a formal charge.
Find the atom's valence electrons from the periodic table. This is the number of electrons the atom "should" have when it is free.
Count the electrons the atom "entirely owns" in the molecule. This is different from checking whether the atom has an octet.
Lone-pair electrons count fully. Every electron in a lone pair belongs completely to that one atom, so add all of them to the "entirely owned" tally.
Bonding electrons count as half. Each bond contains two shared electrons. The atom only owns one of them per bond, so add half the bonding electrons to the tally.
Compare the two numbers.
If the atom entirely owns more electrons than the periodic table says it should have, the atom carries a negative formal charge.
If the atom entirely owns fewer electrons than the periodic table says, the atom carries a positive formal charge.
If the numbers match, the formal charge is zero.
Formal Charge = (Valence electrons) – (Lone-pair electrons) – (1/2 × Bonding electrons)
Or equivalently:
FC = V – L – B/2
Where:
V = number of valence electrons on the free atom (from the periodic table)
L = number of non-bonding (lone-pair) electrons on the atom in the molecule
B = number of bonding electrons around the atom (each single bond = 2, double bond = 4, triple bond = 6)
Total valence electrons: N = 5, O = 6, minus 1 for the positive charge = 10 e⁻.
Lewis structure: [:N=O:]⁺ or [:N≡O:]⁺ (double or triple bond, with lone pairs distributed accordingly).
In the double-bond form: N has a lone pair (2 e⁻) and two bonds (4 e⁻ shared). Entirely owned = 2 + 2 = 4. Valence electrons = 5. FC on N = 5 – 4 = +1.
O has two lone pairs (4 e⁻) and two bonds (4 e⁻ shared). Entirely owned = 4 + 2 = 6. Valence electrons = 6. FC on O = 6 – 6 = 0.
In the triple-bond form: N has one lone pair (2 e⁻) and three bonds (6 e⁻ shared). Entirely owned = 2 + 3 = 5. FC on N = 5 – 5 = 0.
O has one lone pair (2 e⁻) and three bonds (6 e⁻ shared). Entirely owned = 2 + 3 = 5. Valence electrons = 6. FC on O = 6 – 5 = +1.
Either way, the net charge on the ion is +1.
Total valence electrons: C = 4, N = 5, plus 1 for the negative charge = 10 e⁻.
Lewis structure: [:C≡N:]⁻
C has one lone pair (2 e⁻) and a triple bond (6 e⁻ shared). Entirely owned = 2 + 3 = 5. Valence electrons = 4. FC on C = 4 – 5 = –1.
N has one lone pair (2 e⁻) and a triple bond (6 e⁻ shared). Entirely owned = 2 + 3 = 5. Valence electrons = 5. FC on N = 5 – 5 = 0.
Net charge = –1.
Total valence electrons: C = 4, 3H = 3, O = 6, plus 1 for the negative charge = 14 e⁻.
Lewis structure: H–C(–H)(–H)–O: with three lone pairs on O.
C has no lone pairs and four bonds (8 e⁻ shared). Entirely owned = 0 + 4 = 4. Valence = 4. FC = 0.
Each H has no lone pairs and one bond (2 e⁻ shared). Entirely owned = 0 + 1 = 1. Valence = 1. FC = 0.
O has three lone pairs (6 e⁻) and one bond (2 e⁻ shared). Entirely owned = 6 + 1 = 7. Valence = 6. FC = 6 – 7 = –1.
Net charge = –1.
Total valence electrons: C = 4, 3H = 3, minus 1 for the positive charge = 6 e⁻.
Lewis structure: H–C(–H)(–H) with no lone pairs on C (only three bonds).
C has no lone pairs and three bonds (6 e⁻ shared). Entirely owned = 0 + 3 = 3. Valence = 4. FC = 4 – 3 = +1.
Each H: FC = 0 (same reasoning as above).
Net charge = +1. Notice C has only three bonds instead of its usual four, which flags the formal charge.
Total valence electrons: N = 5, 3O = 18, H = 1 = 24 e⁻.
Lewis structure: O=N(–O⁻)(–O–H), with one N=O double bond, one N–O single bond to O⁻, and one N–O single bond to OH.
N has no lone pairs, one double bond and two single bonds (8 e⁻ shared). Entirely owned = 0 + 4 = 4. Valence = 5. FC = 5 – 4 = +1.
Double-bonded O has two lone pairs (4 e⁻) and a double bond (4 e⁻ shared). Entirely owned = 4 + 2 = 6. Valence = 6. FC = 0.
Singly bonded O (with H) has two lone pairs (4 e⁻) and two single bonds (4 e⁻ shared). Entirely owned = 4 + 2 = 6. FC = 0.
Terminal O (no H) has three lone pairs (6 e⁻) and one single bond (2 e⁻ shared). Entirely owned = 6 + 1 = 7. Valence = 6. FC = 6 – 7 = –1.
Net charge = +1 + (–1) = 0 (neutral molecule, as expected).
Total valence electrons: 2C = 8, 2O = 12, 3H = 3, plus 1 for the negative charge = 24 e⁻.
Lewis structure: H₃C–C(=O)(–O:⁻)
Methyl C: four bonds, no lone pairs. FC = 0.
Carbonyl C: three bonds (one double to O, one single to O, one single to CH₃), no lone pairs. Entirely owned = 0 + 4 = 4 (double bond counts as one full bond for the "entirely owned" tally since it contributes 2 electrons: half of 4). Wait, let us be precise: double bond = 4 bonding e⁻, single bond = 2 bonding e⁻ each. Carbonyl C has 4 + 2 + 2 = 8 bonding electrons, half = 4. Valence = 4. FC = 0.
Double-bonded O: two lone pairs (4 e⁻) and one double bond (4 e⁻ shared). Entirely owned = 4 + 2 = 6. FC = 0.
Singly bonded O: three lone pairs (6 e⁻) and one single bond (2 e⁻ shared). Entirely owned = 6 + 1 = 7. Valence = 6. FC = –1.
Net charge = –1.
Formal charge is how chemists pinpoint reactive sites in drug molecules, deciding which atom on a compound will bond to a biological target. It also matters in materials science: the charge distribution in a polymer or crystal determines properties like conductivity and solubility, and formal charge analysis is the first step in mapping that distribution.
Students often confuse formal charge with oxidation state. They are different bookkeeping systems. Formal charge assumes electrons in a bond are shared equally; oxidation state assigns them entirely to the more electronegative atom.
Students sometimes think formal charge means the atom is actually carrying a full ionic charge. It does not. Formal charge is a formalism for electron bookkeeping, not a measurement of real electron density.
A common error is counting all bonding electrons instead of half. Remember: you only "own" half the electrons in each bond.
Students forget to check for formal charges on atoms that have an unusual number of bonds. If nitrogen has four bonds instead of three, or oxygen has one bond instead of two, a formal charge is almost certainly present.
⚠️ Formal charge calculation is a near-guaranteed exam question in introductory organic chemistry. You will be asked to draw a Lewis structure and assign formal charges, or to identify which atom carries the charge in a given ion.
⚠️ Expect questions where the molecule is neutral overall but individual atoms carry non-zero formal charges that cancel out (as in HNO₃). These trip students up because they assume a neutral molecule has no formal charges anywhere.
⚠️ Know the shortcut: if an atom has its "normal" number of bonds (C = 4, N = 3, O = 2, H = 1), its formal charge is zero. Any deviation from that number is a red flag.
⚠️ Be ready to compare resonance structures using formal charge: the structure with the smallest formal charges (closest to zero) on each atom is generally preferred.
True or False: Formal charge equals the actual electric charge on an atom in a molecule.
Fill in the blank: To calculate formal charge, lone-pair electrons count ______ and bonding electrons count as ______.
True or False: A nitrogen atom with four bonds and no lone pairs has a formal charge of +1.
Fill in the blank: If an atom entirely owns more electrons in a molecule than it has as a free atom, its formal charge is ______.
True or False: In a neutral molecule, all atoms must have a formal charge of zero.
Answers: 1. False (it is a formalism, not a measured charge). 2. Fully; half. 3. True (valence 5, entirely owns 4, FC = +1). 4. Negative. 5. False (individual atoms can carry non-zero formal charges that cancel to give a net charge of zero).
Q: What is the formal charge on nitrogen in the ammonium ion (NH₄⁺)?
A: Nitrogen has 5 valence electrons. In NH₄⁺ it has four bonds (8 bonding electrons) and no lone pairs. Entirely owned = 0 + 4 = 4. FC = 5 – 4 = +1.
Q: In the cyanide ion (CN⁻), which atom carries the negative formal charge?
A: Carbon. C has 4 valence electrons but entirely owns 5 in the molecule (one lone pair + half of a triple bond). FC = 4 – 5 = –1. Nitrogen's formal charge is zero.
Q: Draw the Lewis structure of nitric acid (HNO₃) and assign formal charges to every atom.
A: The nitrogen is bonded to three oxygens (one double bond, two single bonds). N has FC = +1 (no lone pairs, four bonds worth of shared electrons, owns 4, valence is 5). The double-bonded O has FC = 0. The O bonded to H has FC = 0. The terminal O with three lone pairs has FC = –1. Net charge on the molecule = 0.
Q: A carbon atom has three bonds and one lone pair. What is its formal charge?
A: C has 4 valence electrons. Entirely owned = 2 (lone pair) + 3 (half of 6 bonding electrons) = 5. FC = 4 – 5 = –1. This carbon carries a formal charge of –1 (it is a carbanion).
Q: Why is checking for an unusual number of bonds a useful shortcut for spotting formal charges?
A: Each common atom has a "normal" bond count where FC = 0 (C = 4, N = 3, O = 2, H = 1). When an atom deviates from this count, the electron bookkeeping will not balance, and a non-zero formal charge results. Spotting the unusual bond count is faster than running the full calculation every time.
Formal charge connects directly to resonance structures (Ch. 1.5): when you draw multiple valid Lewis structures for the same molecule, you use formal charges to judge which contributor is more stable (fewer and smaller formal charges win). It also feeds into acid-base chemistry and reaction mechanisms, where identifying the electron-rich (nucleophilic) and electron-poor (electrophilic) sites on a molecule is the first step in predicting what reacts with what. If you go on to study molecular orbital theory, formal charge gives you the classical-model baseline that MO theory refines.
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