Lewis Structures: Drawing Molecules on Paper

What a Lewis structure shows — and hides

A Lewis structure is chemistry's shorthand for electron bookkeeping. Dots represent valence electrons sitting as lone pairs on an atom; lines represent pairs shared between two atoms — the covalent bonds. Inner-shell electrons never appear, because they take no part in bonding.

What the drawing shows is connectivity: which atoms are joined to which, how many pairs they share, and where the unused electrons sit. What it hides is just as important: it shows nothing about bond angles, molecular shape, or the actual sizes of atoms. Water drawn as H−O−H looks straight on paper, yet the real molecule is bent.

Think of it as the symbolic level made visual. The formula H2O tells you the composition; the Lewis structure adds the wiring diagram. Every later tool in bonding — formal charge, resonance, VSEPR — starts by reading this diagram, so drawing it correctly is the gateway skill.

The five-step drawing method

Follow the same routine every time and the structures almost draw themselves:

1. Count valence electrons. Sum the group numbers of all atoms. For an ion, add one electron per negative charge and subtract one per positive charge.
2. Choose the central atom. Pick the least electronegative atom — never hydrogen, which can only form one bond.
3. Connect with single bonds. Draw one line from the central atom to each outer atom. Each line consumes 2 electrons.
4. Complete the outer octets. Distribute the remaining electrons as lone pairs on the outer atoms first (hydrogen needs only 2).
5. Finish the center. Place any leftover electrons on the central atom. If it still lacks an octet, convert neighboring lone pairs into double or triple bonds until the octet is satisfied.

Work through the steps in order and check the electron total at the end — the count is your built-in error detector.

Worked example: carbon dioxide

Carbon dioxide puts the method to work. First, count: carbon brings 4 valence electrons and each oxygen brings 6, for a total of 16. Carbon is less electronegative than oxygen, so it goes in the center.

Connect with single bonds: O−C−O uses 4 electrons, leaving 12. Give each oxygen its octet: three lone pairs per oxygen consumes all 12 remaining electrons. Now check carbon — it has only 4 electrons around it, well short of an octet.

Convert one lone pair from each oxygen into a bonding pair, creating two double bonds: O=C=O. Recount: each oxygen now has 2 lone pairs (4 electrons) plus 4 bonding electrons = 8; carbon has 8 bonding electrons = 8. All 16 valence electrons are placed, every atom has an octet, and (as a bonus check) every formal charge is zero. The structure is complete and correct.

The octet rule and its famous exceptions

The octet rule — eight electrons around each atom — works beautifully for carbon, nitrogen, oxygen, and the halogens, but several families of exceptions appear on exams:

• Incomplete octets. Hydrogen and helium settle for 2 (the duet rule). Beryllium compounds like BeCl2 stop at 4 electrons, and boron compounds like BF3 stop at 6 — there simply are not enough electrons to go around.
• Expanded octets. Elements in period 3 and below have empty d orbitals available, so they can exceed eight: PCl5 holds 10 electrons around phosphorus, and SF6 holds 12 around sulfur.
• Odd-electron molecules. Some molecules have an odd total, making full pairing impossible. Nitric oxide (NO) has 11 valence electrons, leaving one unpaired — part of why it is so reactive.

When drawing, let the electron count overrule the rule: place every valence electron first, then judge. A correct structure with an exception beats a wrong structure with forced octets.

Reading a structure like a chemist

A finished Lewis structure is a readable map. Lone pairs mark the molecule's reactive sites — they are the electrons most available to form new bonds, so ammonia's lone pair makes it a base and a ligand. Multiple bonds flag regions of high electron density, the places electrophiles attack.

Connectivity alone can distinguish isomers. Ethanol and dimethyl ether share the formula C2H6O, but their Lewis structures wire the atoms differently: ethanol has an O−H group (so it hydrogen-bonds and boils at 78 °C), while dimethyl ether does not (boiling at −24 °C). Same atoms, different wiring, different substance.

The structure also previews the next steps. Count the electron domains around the central atom and VSEPR predicts the shape; compute formal charges and you can judge whether the structure is the best one. The Lewis drawing is the raw data that every other bonding analysis consumes.

Mistakes that cost marks

A few errors account for most lost marks. First, miscounting valence electrons — especially forgetting the charge adjustment for ions. Sulfate is SO4 2−: 6 + 4(6) + 2 = 32 electrons, not 30.

Second, putting hydrogen in the center. Hydrogen forms exactly one bond, so it is always terminal; a central hydrogen is an instant giveaway of a wrong structure.

Third, giving hydrogen more than 2 electrons. Hydrogen's valence shell holds only the 1s orbital — two electrons maximum, no exceptions.

Fourth, leaving the central atom short of an octet without trying multiple bonds. If electrons remain unused and the center is incomplete, lone pairs must be converted before you declare the structure finished.

Finally, drawing a structure whose electron total doesn't match step 1. Recount dots and lines: every line is 2 electrons, every dot is 1. If the total is off, the structure is wrong — find the missing or extra electrons before moving on.

Common mistakes

Real-world applications

• Sketching drug candidates before computer modeling

• Predicting where a molecule will react from lone pairs and pi bonds

• Checking whether a proposed structure obeys electron counts

• Teaching bond types in every chemistry classroom