Source: Chapter 1, Prof. Forrest Michael, University of Washington
Difficulty: Introductory | Prerequisites: General chemistry (atomic structure, electron configuration, electronegativity trends).
Formal charge and bond order are foundational tools you will use in every topic from here forward in organic chemistry. They let you evaluate Lewis structures, decide which resonance contributor matters most, and predict how molecules behave in reactions. If you took general chemistry, you have seen electrons, bonds, and electronegativity before, but organic chemistry demands you apply those ideas quickly and accurately to carbon-based molecules. Getting comfortable with formal charge calculations and bond-order reasoning now saves enormous time when you reach reaction mechanisms in later chapters.
Formal charge tells you whether an atom in a Lewis structure is carrying more or fewer electrons than it normally would. Bond order tells you how many bonds effectively sit between two atoms (single, double, triple, or something in between). Together, they are the starting point for drawing good Lewis structures, ranking resonance contributors, and predicting molecular properties.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons minus (non-bonding electrons + number of bonds). In simple terms, it is the difference between what an atom "should" have and what it "gets" in a particular structure.
Bond order
The number of chemical bonds between a pair of atoms, calculated as: (bonding electrons in molecular orbitals minus antibonding electrons) divided by 2. Think of it as a way to express whether a bond is single, double, triple, or somewhere in between.
Valence electrons
The electrons in an atom's outermost shell, which participate in bonding. For second-row elements: B has 3, C has 4, N has 3 bonds typically, O has 2, F has 1.
Non-bonding electrons (lone pairs)
Electron pairs on an atom that are not shared in a bond. These count fully toward formal charge because the atom "owns" them entirely.
Lewis structure
A diagram showing how valence electrons are distributed among atoms in a molecule, using dots for lone pairs and lines for bonds. In simple terms, it is the molecular blueprint that tells you where the electrons live.
Formal charge = valence electrons - (non-bonding electrons + number of bonds)
Each bond counts as one bond toward the atom, regardless of whether it is sigma or pi
Apply the formula atom by atom across the structure
Example: CO (carbon monoxide)
Carbon and oxygen share a triple bond
Carbon: 4 valence electrons - (0 lone pair electrons + 3 bonds) = +1 formal charge... but in the Lewis structure shown in lecture, carbon gets a lone pair and a triple bond: 4 - (2 + 3) = -1
Oxygen: 6 - (2 + 3) = +1
This gives the perhaps surprising result that carbon carries the negative formal charge in CO
B: 3 valence electrons, typically forms 3 bonds
C: 4 valence electrons, typically forms 4 bonds
N: 5 valence electrons, typically forms 3 bonds (with one lone pair)
O: 6 valence electrons, typically forms 2 bonds (with two lone pairs)
F: 7 valence electrons, typically forms 1 bond (with three lone pairs)
Bond order = (number of electrons in bonding MOs - number of electrons in antibonding MOs) / 2
A bond order of 1 = single bond, 2 = double bond, 3 = triple bond
Fractional bond orders are possible when resonance is involved (e.g., 1.5 in benzene)
Write out the atoms: H-C with connections to two oxygens and another H
Assign bonds and lone pairs so each atom reaches a full octet (or duet for H)
Calculate formal charges for each atom to verify the structure
The correct Lewis structure has: C double-bonded to one O, single-bonded to the other O (which also bonds to H), with lone pairs on each oxygen
Formal charge:
FC = V - (N + B)
Where V = valence electrons, N = non-bonding electrons, B = number of bonds.
Bond order:
BO = (bonding e⁻ - antibonding e⁻) / 2
Formal charge helps chemists predict which sites on a molecule are electron-rich or electron-poor, which directly determines where reactions happen. Drug designers, for instance, use formal charge patterns to understand how a molecule will bind to a protein target. Bond order determines bond strength and length, which matters everywhere from materials science (why diamond is hard: each carbon forms four strong single bonds in a 3D network) to atmospheric chemistry (the triple bond in N₂ is why nitrogen gas is so unreactive).
Students often count the total electrons in a bond (two per bond) instead of counting each bond as one toward each atom. The formula uses the number of bonds, not the number of bonding electrons.
Formal charge is not the same as oxidation state. Formal charge splits bonding electrons equally between the two atoms; oxidation state gives them all to the more electronegative atom. These are different bookkeeping systems.
A formal charge of zero does not necessarily mean the atom is neutral in a physical sense. It means the electron distribution in that particular Lewis structure matches the atom's expected valence.
Students sometimes assume that the structure with all formal charges of zero is always the "right" one. It is usually the most stable, but resonance means several structures contribute, and some with non-zero formal charges can still matter.
⚠️ You will almost certainly be asked to calculate formal charges on a Lewis structure. Practise until the formula is automatic.
⚠️ Bond order questions often appear alongside resonance: you may be given resonance structures and asked to determine the bond order of a particular bond in the hybrid.
⚠️ The valence/bond-count table for B, C, N, O, F is worth memorising. It lets you spot errors in Lewis structures at a glance: if carbon has five bonds, something is wrong.
True or false: Formal charge equals the number of valence electrons minus the number of bonds.
False. It equals valence electrons minus (non-bonding electrons + bonds).
Fill in the blank: Carbon typically forms ___ bonds in a neutral, stable molecule.
Four.
True or false: A bond order of 1.5 is impossible.
False. Fractional bond orders arise from resonance (e.g., the C-C bonds in benzene).
Fill in the blank: In the bond-order formula, you subtract the number of ___ electrons from the number of bonding electrons, then divide by 2.
Antibonding.
True or false: Fluorine typically forms three bonds.
False. Fluorine typically forms one bond.
Q: Calculate the formal charge on each atom in the Lewis structure of CO where carbon has a lone pair and a triple bond to oxygen (which also has a lone pair).
A: Carbon: 4 - (2 + 3) = -1. Oxygen: 6 - (2 + 3) = +1. Carbon carries a formal charge of -1 and oxygen carries +1.
Q: What is the bond order of the C-O bonds in the formate ion (HCO₂⁻), given that two equivalent resonance structures exist?
A: Each resonance structure shows one C=O double bond and one C-O single bond. Averaging gives a bond order of 1.5 for each C-O bond.
Q: An atom of nitrogen in a Lewis structure has one lone pair and three bonds. What is its formal charge?
A: FC = 5 - (2 + 3) = 0. Nitrogen has a formal charge of zero, which is the expected situation for a neutral nitrogen.
Q: Why can boron form only three bonds and still be stable, even though it does not have a full octet?
A: Boron has only three valence electrons. It forms three bonds and has an empty p orbital. It is electron-deficient but stable enough in many compounds (e.g., BF₃) because the fluorines can donate electron density back into boron's empty orbital via resonance.
Formal charge feeds directly into resonance analysis (Lecture 2): you rank resonance contributors partly by which atoms carry formal charges and whether those charges sit on appropriate atoms. Bond order connects to bond length and bond strength (also Lecture 2): higher bond order means a shorter, stronger bond. Both concepts reappear heavily when you study reaction mechanisms in Chapter 2 and beyond, where electron-pushing arrows track how formal charges shift as bonds form and break.
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