A Lewis formula is a drawing that shows every single outer electron in a molecule — the pairs being shared and the pairs that are not. Follow the same four steps every time and these become some of the easiest marks in the paper.
📚 What you need to know
A Lewis formula shows all valence (outer) electrons: both bonding pairs and lone pairs.
A pair can be drawn as two dots, two crosses, a dot and a cross, or a line.
Count valence electrons from the group number: C = 4, N = 5, O = 6, halogens = 7, H = 1.
The central atom is usually the least electronegative one — and it is never hydrogen.
Fill the outer atoms to 8 electrons first (H needs only 2), then whatever is left goes on the central atom.
Be (4 electrons) and B (6 electrons) are stable without a full octet.
What counts as a Lewis formula
Here is the trap. If you draw chlorine as Cl–Cl, you have shown the bond but you have not shown the six lone pairs on each atom. That is a structural formula, not a Lewis formula, and it will not get the mark.
A Lewis formula has to account for every outer electron. You can show a pair however you like:
Two dots ( : ) or two crosses ( ✕✕ )
A dot and a cross — useful because it shows which atom each electron came from
A line — but a line can only replace a bonding pair, and you still have to draw the lone pairs
Dots and crosses do not mean the electrons are actually different. All electrons are identical. We colour them differently only so we can see where each one started.
The four steps
🧩 Drawing any Lewis formula
Count the total valence electrons. Add up the outer electrons of every atom. (For an ion: add 1 for each negative charge, take away 1 for each positive charge.)
Draw the skeleton. Put the least electronegative atom in the middle and arrange the others around it. Hydrogen always goes on the outside.
Add the bonding pairs — one pair between each joined pair of atoms. Then add lone pairs to the outer atoms until each has 8 (hydrogen stops at 2).
Check. Have you used exactly the number of electrons from step 1? Does every atom have a full shell? If the central atom is short, pull a lone pair in to make a double bond.
Worked through: methane, CH4
Step 1 — count. Carbon brings 4, each hydrogen brings 1. Total = 4 + (4 × 1) = 8 electrons, which is 4 pairs.
Step 2 — skeleton. Carbon in the middle (hydrogen can never be central), four hydrogens around it.
Step 3 — add pairs. Four bonding pairs uses up all 8 electrons, so there is nothing left for lone pairs.
Step 4 — check. Carbon has 8 electrons around it. Each hydrogen has 2. Done.
Methane. All eight valence electrons are used up in the four bonding pairs, so there are no lone pairs anywhere.
Molecules with lone pairs
Ammonia and water are built the same way, but this time the bonding pairs do not use up all the electrons — what is left over stays on the central atom as lone pairs.
NH3: 5 + (3 × 1) = 8 electrons = 4 pairs. Three pairs go into N–H bonds, so 1 lone pair is left on the nitrogen.
H2O: 6 + (2 × 1) = 8 electrons = 4 pairs. Two pairs go into O–H bonds, leaving 2 lone pairs on the oxygen.
Ammonia keeps one lone pair on the nitrogen; water keeps two on the oxygen. Both still have exactly four pairs in total.
Why lone pairs matter so much. They are invisible in a formula like H2O, but they decide the shape of the molecule and its polarity. Miss a lone pair here and every answer that follows will be wrong — which is exactly why you always draw the Lewis formula first.
When the outer atoms have lone pairs too
Halogens bring 7 electrons each. One goes into the bond, and the other six sit as three lone pairs. Tetrachloromethane shows this nicely:
CCl4 = 4 + (4 × 7) = 32 electrons = 16 pairs. Four pairs are used for the C–Cl bonds, so 12 pairs are left — three on each chlorine.
Tetrachloromethane. Each chlorine uses one electron for bonding and holds the other six as three lone pairs.
WORKED EXAMPLE
How many valence electrons are there in a molecule of methanol, CH3OH?
Add up the outer electrons, atom by atom1 carbon (group 14), 4 hydrogens, 1 oxygen (group 16).C: 1 × 4 = 4H: 4 × 1 = 4O: 1 × 6 = 6Total = 14 electrons = 7 pairs5 bonding pairs (3 C–H, 1 C–O, 1 O–H) and 2 lone pairs on the oxygen.
When there are not enough electrons to go round
Sometimes you finish step 3 and the central atom is still short. Don’t add more electrons — you don’t have any. Instead, take a lone pair from an outer atom and slide it into the bond to make a double bond.
CO2: 4 + (2 × 6) = 16 electrons. Two single bonds would leave carbon with only 4, so each oxygen shares a second pair → two C=O double bonds.
HCN: 1 + 4 + 5 = 10 electrons. Carbon needs three pairs from nitrogen, giving a C≡N triple bond, and nitrogen keeps 1 lone pair.
When single bonds leave the central atom short, an outer atom shares a second or third pair instead.
The two exceptions you must be able to draw
Beryllium and boron are the famous rule-breakers. Both make stable molecules with an incomplete octet, and both are popular exam questions precisely because students “correct” them by mistake.
BeCl2: 2 + (2 × 7) = 16 electrons. Beryllium ends up with only 4 electrons around it — and that is the right answer.
BF3: 3 + (3 × 7) = 24 electrons. Boron ends up with only 6. Each fluorine still gets its full octet.
Beryllium is stable with four electrons and boron with six. The outer atoms still reach a full octet.
💡 Exam tip
Always finish by counting your electrons and comparing with step 1. If the numbers don’t match, something is missing.
Never put hydrogen in the middle — it can only make one bond, so it cannot hold a molecule together.
Draw lone pairs as proper pairs, touching the atom they belong to. Loose single dots lose marks.
For a charged species, put the whole thing in square brackets with the charge outside, like [NH4]+.
⚠️ Common mix-up
Cl–Cl is not a Lewis formula. The lone pairs must be shown.
Hydrogen never gets lone pairs. Two electrons and it is completely full.
Don’t “fix” BeCl2 or BF3 by adding extra electrons. They really are stable with less than 8.
A line replaces a bonding pair only. You still have to draw every lone pair as dots.
Up next: Double and Triple Bonds — what happens when atoms share two or three pairs, and why that makes bonds shorter and stronger.
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