What a Lewis structure shows
A Lewis dot structure is a map of where the valence electrons go in a molecule. Every shared (bonding) pair is drawn as a line between two atoms, and every pair that stays put on a single atom is drawn as a pair of dots — a lone pair. A single bond is one line, a double bond two, and a triple bond three. Pick a molecule in the tool above to draw it and watch the electron count add up.
The structure is more than a picture: it tells you which atoms are connected, how many bonds hold them together, and where the leftover electrons sit. Those lone pairs are exactly what later decide a molecule’s shape in molecular geometry.
Counting valence electrons
Start by adding the valence electrons of every atom, which you can read straight off the group number. Oxygen (group 16) brings 6 and each hydrogen brings 1, so water has 6 + 1 + 1 = 8 valence electrons. Those eight have to appear in the finished drawing: two O–H bonds use 4 electrons (2 per bond) and oxygen’s two lone pairs use the other 4.
The tool checks this for you. It totals the bonding electrons (2 for every bond) plus the lone-pair electrons (2 for every pair), and that sum must match the electrons every atom contributed. If they don’t match, the structure is wrong.
The octet rule
Most main-group atoms are “happiest” surrounded by eight valence electrons — the same full outer shell that makes the noble gases so stable. That is the octet rule. In water the oxygen is surrounded by eight (two bonds plus two lone pairs), and each hydrogen reaches its duet of two. Carbon in CO₂ and nitrogen in N₂ get to eight by forming double and triple bonds rather than extra lone pairs.
Octet-rule exceptions
The octet rule is a strong guideline, not a law, and the tool includes the classic exceptions:
- Electron-deficient: boron in BF₃ forms only three bonds, so it shares just six electrons. It is stable anyway — boron simply has too few valence electrons to reach eight.
- Expanded octet: phosphorus in PCl₅ holds ten electrons and sulfur in SF₆ holds twelve. Only larger atoms in period 3 and below are big enough to fit this many bonding pairs around the central atom. (Older textbooks credit empty d-orbitals; modern models show that effect is tiny, so it’s better thought of as a size effect.)
Watching these side by side with normal molecules makes it clear why “complete,” “deficient,” and “expanded” are the three labels the tool shows.
Using this with a class
Have students predict the number of bonds and lone pairs before selecting a molecule, then check the drawing and the electron tally against their work. Comparing BF₃, PCl₅, and SF₆ to water or methane is a quick way to teach the octet rule and its exceptions in one screen. This interactive is free to embed on your own site or LMS using the snippet below.