IB Chemistry HLTopic 2 — Models of Bonding & StructurePaper 1 & 2Higher level~8 min read
Formal Charge
Sometimes you draw a Lewis formula, check it carefully, and it is completely valid — and so is the different one your friend drew. Formal charge is the tie-breaker. It is a bookkeeping tool that tells you which structure chemistry actually prefers, and it takes about thirty seconds to apply.
📘 What you need to know
Formal charge (FC) is the charge an atom would have if every bonding pair were shared perfectly equally.
It deliberately ignores electronegativity. That is what makes it a fair comparison between structures.
The formula is FC = V − ½B − N, where V is valence electrons, B is bonding electrons and N is non-bonding electrons.
The preferred structure is usually the one where the difference between the highest and lowest formal charges is closest to zero.
Any negative formal charge should sit on the most electronegative atom.
The formal charges must add up to the overall charge on the species — a free accuracy check.
Formal charge is a model, not a real measured charge on the atom.
What the formula is actually counting
Start with how many outer electrons the atom brought to the party — that is V, straight off the periodic table. Then work out how many it appears to own in the structure you have drawn: all of its lone pair electrons, plus exactly half of every bonding pair it is involved in. Subtract, and the difference is the formal charge.
Learn this version
FC = V − ½B − N valence − half the bonding electrons − the non-bonding ones
Working out the formal charge on boron in BH₄⁻
Notice the final check in the bottom line. If your formal charges do not add up to the overall charge, you have miscounted somewhere — go back before you go on.
Formal charge is not the real charge. In BH4− the negative charge is spread over the whole ion, and boron is not sitting there with a genuine −1. It is a bookkeeping label that helps you compare structures, nothing more.
Picking between two structures
Once you can calculate formal charges, comparing structures is mechanical. Work out the formal charge on every atom in each candidate, then apply two rules in order.
🧩 How to choose
Calculate FC for every atom in each structure using FC = V − ½B − N.
Check they sum correctly to the overall charge on the molecule or ion.
Find ΔFC for each structure: the highest formal charge minus the lowest.
Prefer the structure with ΔFC closest to zero. Formal charges near zero means the electrons are already sensibly placed.
Use electronegativity as a tie-breaker. Any negative formal charge should sit on the most electronegative atom.
Two ways of drawing sulfur dioxide
Lone pairs are left off both drawings to keep the formal charges readable. In a full answer you would show two lone pairs on each doubly bonded oxygen and three on the singly bonded one.
🤔 Why does a smaller ΔFC mean a better structure?
Separating charge costs energy. A structure that puts +1 on one atom and −1 on another has pulled electrons away from where they naturally sit, and something has to pay for that. A structure where everything comes out at zero has the electrons distributed sensibly to begin with, so it is lower in energy and contributes more to the real molecule.
Formal charge and resonance
This is where formal charge earns its keep. When you have several resonance structures, they do not all contribute equally to the hybrid. The ones with formal charges closest to zero contribute most.
For carbon dioxide, the O=C=O structure gives every atom a formal charge of zero, so it dominates. The alternatives with a triple bond on one side put +1 on one oxygen and −1 on the other, giving ΔFC = 2, so they contribute far less to the real molecule.
Atom in a neutral molecule
Bonds
Lone pairs
Formal charge
Carbon
4
0
0
Nitrogen
3
1
0
Nitrogen
4
0
+1
Oxygen
2
2
0
Oxygen
1
3
−1
Oxygen
3
1
+1
🧠 A pattern worth spotting
For the common atoms, formal charge is just the usual number of bonds compared with what you have drawn. Carbon normally makes 4, nitrogen 3, oxygen 2. One bond more than usual gives +1; one bond fewer gives −1. It matches the formula every time, and it is much faster in an exam.
Worked examples
WE 1
Calculate the formal charge on the nitrogen atom in NH4+
Step 1: find V
Nitrogen is group 15, so V = 5.
Step 2: find B and halve it
Four N–H bonds → B = 8, so ½B = 4Step 3: find N
The lone pair was used to make the fourth bond, so N = 0.
Step 4: substituteFC = 5 − 4 − 0 = +1FC on nitrogen = +1each H is 0, so the total is +1 — matching the ion’s charge
WE 2
Two structures are proposed for CO2: (I) O=C=O and (II) O≡C–O. Use formal charge to decide which is preferred.
Structure I: O=C=O
Carbon: 4 − ½(8) − 0 = 0
Each oxygen: 6 − ½(4) − 4 = 0∆FC = 0 − 0 = 0Structure II: O≡C–O
Carbon: 4 − ½(8) − 0 = 0
Triple bonded O: 6 − ½(6) − 2 = +1
Single bonded O: 6 − ½(2) − 6 = −1∆FC = (+1) − (−1) = 2Apply the rule
Smaller ΔFC wins.
Structure I is preferredboth sum to zero overall, which confirms neither calculation slipped
WE 3
In the thiocyanate ion, SCN−, one structure has a triple C≡N bond. Find the formal charge on sulfur if it has one bond and three lone pairs.
Step 1: V for sulfur
Group 16, so V = 6.
Step 2: B and N
One single bond → B = 2, ½B = 1
Three lone pairs → N = 6Step 3: substituteFC = 6 − 1 − 6 = −1FC on sulfur = −1carbon and nitrogen both come out at 0 here, so the total is −1 as required
💡 Exam tips
Write the formula down first and substitute into it. Method marks are available even if the arithmetic slips.
Always check the sum. Formal charges must add to the overall charge, so this catches most errors instantly.
Halve the bonding electrons, not the bonds. A double bond contributes 4 electrons, so 2 after halving.
State ΔFC explicitly when comparing structures, and say which is smaller.
Remember the second rule: negative formal charge belongs on the most electronegative atom.
Say formal charge is a model if asked to comment on it — it assumes equal sharing, which is not real.
⚠ Common mix-ups
Confusing formal charge with oxidation state. Formal charge assumes equal sharing; oxidation state gives both electrons to the more electronegative atom.
Counting bonds instead of bonding electrons. B is electrons, so a double bond gives 4.
Forgetting to halve. The ½ is the whole point of the model.
Preferring the structure with the biggest ΔFC. It is the other way round — closest to zero wins.
Treating formal charge as a real charge. It is a bookkeeping device, not something you could measure.
Skipping the sum check and handing in an answer that could not possibly be right.
Up next: Sigma and Pi Bonds — a closer look at what a covalent bond really is, and why a double bond is made of two quite different halves.
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