🔗 Bonding & Molecular Structure · Lesson 2 of 5
Lewis Structures
A dot-and-line map of every valence electron in a molecule — the blueprint you draw before you can predict shape, polarity, or reactions.
⏱️ About 22 min
Illustrative image (AI-generated) — decorative only.
Before a chemist can say what a molecule looks like or how it will react, they draw its Lewis structure — a simple bookkeeping diagram that tracks where every valence electron lives. Master this one skill and molecular shape, polarity, and reactivity all become readable.
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The big idea: A Lewis structure shows a molecule's valence electrons as shared pairs (bonds) and lone pairs (dots). Draw it by counting valence electrons, connecting the atoms, and distributing the rest so each atom reaches an octet — then check your work with formal charge.
Bonds and lone pairs: two homes for electrons
In a covalent molecule, valence electrons live in one of two places: in a bonding pair shared between two atoms (drawn as a line), or in a lone pair that belongs to just one atom (drawn as two dots). A Lewis structure is simply a picture that accounts for all of a molecule's valence electrons across these two homes.
Only valence electrons — the outer shell — take part. That's why counting valence electrons correctly is step one of every structure.
🔑 The octet rule
Atoms are most stable with eight electrons in their outer shell (a full 'octet', matching a noble gas). Hydrogen is the exception that wants only two (a 'duet'). Count both bonding and lone-pair electrons around an atom toward its octet — a shared pair counts for both atoms it joins.
The step-by-step recipe
- Count valence electrons. Add up the group-number valence electrons of every atom. For an ion, add one electron per negative charge, subtract one per positive charge.
- Pick the central atom. Usually the least electronegative atom (never hydrogen). Place the others around it.
- Draw single bonds from the central atom to each outer atom (2 electrons each).
- Distribute the remaining electrons as lone pairs, filling the outer atoms' octets first, then any left over on the central atom.
- If the central atom is short of an octet, convert a lone pair on a neighbour into a double or triple bond.
- Check with formal charge to confirm the best arrangement.
📝 Worked example: Draw the Lewis structure of carbon dioxide, CO₂.
- Count valence electrons: carbon (group 14) has 4, each oxygen (group 16) has 6. Total = 4 + 6 + 6 = 16.
- Central atom: carbon is least electronegative, so put C in the middle: O C O.
- Single bonds: draw O–C–O. That uses 2 pairs = 4 electrons, leaving 12.
- Fill outer atoms: give each oxygen 3 lone pairs (6 electrons each) — that uses all 12. But now carbon has only 4 electrons around it (two bonds) — not an octet.
- Make double bonds: move one lone pair from each oxygen into a second bond, giving O=C=O. Carbon now has 8 electrons; each oxygen has 2 bonds + 2 lone pairs = 8.
- Formal charge check: every atom's formal charge is 0 — a clean, correct structure.
✓ O=C=O: two C=O double bonds, with two lone pairs on each oxygen. All octets satisfied, all formal charges zero.
\[ \ce{O=C=O} \]
The finished carbon dioxide structure: each double bond is a shared pair of electron pairs, and each oxygen keeps two lone pairs (not shown here).
Formal charge: your error-check
When more than one structure looks possible, formal charge tells you which is most reasonable: the best structure keeps formal charges as close to zero as possible, and puts any negative charge on the most electronegative atom.
\[ FC = (\text{valence e}^-) - (\text{lone-pair e}^-) - \tfrac{1}{2}(\text{bonding e}^-) \]
Formal charge = the atom's own valence electrons, minus the ones it 'keeps' (all lone-pair electrons + half of each shared pair).
⚠️ When the octet rule bends
The octet rule is a guide, not a law. Common exceptions: Hydrogen wants 2, not 8. Boron and beryllium are often happy with fewer (BF₃ leaves boron with only 6). Period-3+ atoms like phosphorus and sulfur can hold an expanded octet (PCl₅, SF₆). Odd-electron molecules (like NO) simply can't pair everything up.
✨ Resonance: one picture isn't always enough
Sometimes several equally good Lewis structures exist — for ozone (O₃) the double bond could go on either side. The real molecule is a blend, or resonance, of them, and both oxygen–oxygen bonds are actually identical. Draw the equivalent structures and connect them with a double-headed arrow.
✏️ Practice: How many valence electrons in total must the Lewis structure of a water molecule (H₂O) account for? (Oxygen is group 16; hydrogen is group 1.)
electrons
Solution
- Oxygen (group 16) contributes 6 valence electrons.
- Each hydrogen (group 1) contributes 1, and there are two of them: 2 × 1 = 2.
- Total = 6 + 2 = 8 valence electrons (two O–H bonds plus two lone pairs on oxygen).
Check your understanding
1. What is the first step in drawing any Lewis structure?
Everything else depends on knowing how many valence electrons you have to place, so counting them is always step one.
2. Which atom can legitimately hold more than eight electrons (an expanded octet)?
Sulfur is in period 3 and can accommodate an expanded octet (as in SF₆). Period-2 atoms like C, N and O are limited to eight.
3. When choosing between possible Lewis structures, the best one generally…
Lower (near-zero) formal charges mean a more stable, more reasonable structure — and any negative charge should sit on the most electronegative atom.
✅ Key takeaways
- A Lewis structure accounts for every valence electron as bonding pairs (lines) or lone pairs (dots).
- Recipe: count valence electrons → central atom → single bonds → fill octets → make multiple bonds if short → check formal charge.
- Octet rule: 8 electrons for most atoms, 2 for hydrogen — a shared pair counts for both atoms.
- Exceptions: H (2), B and Be (fewer), period-3+ atoms (expanded octets), odd-electron molecules.
- Formal charge = valence − lone-pair − ½(bonding); the best structure keeps it near zero.
➡️ A Lewis structure is flat on the page, but molecules are 3-D. The lone pairs and bonds you just mapped actually repel each other in space — and that repulsion sets a molecule's real shape. That's VSEPR, next.
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