Dots and dashes: the chemist's shorthand
Chemists needed a quick way to keep track of where electrons go in a molecule, and over a century ago a simple bookkeeping picture won out. In a Lewis structure, you draw each atom's symbol, put dots around it for its valence electrons, and use a line (a dash) for each shared pair that forms a bond. It looks almost childish — letters, dots, and dashes — yet it captures who is bonded to whom and where the leftover electrons sit.
Two kinds of electron pairs show up. A bonding pair is shared between two atoms — that is the covalent bond, drawn as a line. A lone pair sits on a single atom, not shared with anyone, drawn as two dots. Both kinds count toward that atom's full-shell quota. Keeping bonding pairs and lone pairs straight is the whole skill, and it pays off massively in the next guide when we predict the shapes of molecules.
A recipe you can follow
Drawing a Lewis structure is genuinely a recipe — follow the steps and you almost can't go wrong. Let's take carbon dioxide, CO₂, as our worked example.
- Count all valence electrons. Carbon brings 4, each oxygen brings 6, so 4 + 6 + 6 = 16 electrons to place.
- Put the least greedy (least electronegative) atom in the middle. Carbon goes in the center, with an oxygen on each side.
- Connect with single bonds first, then spread the rest as lone pairs on the outer atoms to fill their shells.
- If the central atom is still short of eight, pull lone pairs in to make double or triple bonds. Carbon ends up double-bonded to each oxygen: O=C=O.
Lopsided bonds: polarity
A Lewis structure draws a shared pair as if it sits exactly in the middle, but you already know better. If the two atoms differ in electronegativity, the greedier atom hogs the shared electrons, building up a slight negative charge while its partner is left slightly positive. That uneven sharing is bond polarity: a bond with a partly-negative end and a partly-positive end, even though no electron was fully transferred.
A polar bond has a tiny built-in arrow pointing from the positive end toward the negative end — chemists call its strength and direction the dipole moment. Water is bent and full of polar O–H bonds whose arrows don't cancel, so the whole molecule is polar; that single fact is why water dissolves salt, why ice floats, and why life as we know it exists. Polarity is the bridge from a static dot drawing to how molecules actually behave.
When one picture isn't enough: resonance
Sometimes the recipe gives you a real headache. Take the ozone molecule, O₃. You can draw it with a double bond on the left and a single bond on the right — but you could equally draw the double bond on the right and the single on the left. Both pictures obey every rule. Which one is correct? The honest answer is: neither, and both.
Real measurements show ozone's two bonds are *identical* — each is halfway between a single and a double bond. The true molecule is a blend, an average, of the pictures you can draw. Chemists call this resonance: when no single Lewis structure tells the truth, the real molecule is a hybrid of several. The shared electrons aren't pinned to one spot — they are spread out, or [[delocalization|delocalized]], over the whole stretch.