Valence Electrons and the Octet Rule

Valence electrons are the electrons in the highest occupied principal energy level. For main-group elements the group number gives the count directly: group 1 has 1, group 2 has 2, group 13 has 3, and so on up to group 18 with 8. Sulfur is in group 16, so it has 6 valence electrons regardless of which period it sits in.

For main-group atoms the highest-shell count is the useful starting definition. Helium is the group-18 exception: 1s² gives two valence electrons and a full first shell, not eight. The group-number shortcut does not extend unchanged to transition metals, whose d electrons may also participate in bonding. Use the periodic-table group to obtain the neutral atom's contribution before adjusting a whole species for charge.

A Lewis symbol is the element symbol surrounded by one dot per valence electron. Dots are placed singly on the four sides first and then paired up, which is why nitrogen shows three single dots plus one pair and oxygen shows two single dots plus two pairs. Those unpaired dots are the electrons available for bonding.

The four sides of a Lewis symbol are a drawing convention, not a literal map of the isolated atom's orbital directions. To draw oxygen, place one dot on each side until four dots are used, then add the remaining two as partners, producing two lone pairs and two single dots. Rotating the symbol or choosing different sides for the pairs does not change its electron count. When atoms form a shared pair, a bond line replaces two dots between them.

Main-group atoms are most stable when they are surrounded by eight valence electrons, matching the nearest noble gas. Metals reach that count by losing electrons, nonmetals by gaining or sharing them. Hydrogen is the exception that everyone must remember: it is complete with two electrons, matching helium.

A bonding pair counts toward the octet of both bonded atoms but counts only once in the molecule's total electron budget. In H₂O, two O–H bonds contain four electrons and two oxygen lone pairs contain four more, totaling eight electrons. Around O, the two bonds and two lone pairs supply an octet; each H sees the two electrons in its one bond. Adding “eight around O plus two around each H” would double-count the shared electrons. The octet rule is a useful pattern for many main-group compounds, not the physical cause of every stable bond or an exception-free law.

For a neutral nonmetal, the usual number of covalent bonds is eight minus the group valence count. Carbon with 4 valence electrons forms 4 bonds and no lone pairs, nitrogen with 5 forms 3 bonds and 1 lone pair, oxygen with 6 forms 2 bonds and 2 lone pairs, and fluorine with 7 forms 1 bond and 3 lone pairs. Hydrogen forms exactly 1 bond.

These are neutral patterns, not mandatory counts for charged atoms. Nitrogen in NH₄⁺ has four bonds and no lone pair; oxygen in OH⁻ has one bond and three lone pairs. In both cases the central atom still has an octet, with a nonzero formal charge. A multiple bond counts as two or three shared pairs for the octet, even though it later counts as only one electron domain for geometry.

For NH₄⁺, 5 + 4(1) − 1 = 8 electrons are available; four N–H bonds use all eight. Nitrogen has an octet and every hydrogen a duet. For CO₃²⁻, 4 + 3(6) + 2 = 24 electrons must appear across the entire drawing. Adding charge changes the budget once, not once per atom.

Ask two independent questions: did I use the right total, and does each atom have an allowed local pattern? A drawing can satisfy every octet but contain too many electrons for the specified charge. Conversely, a correct budget can still leave carbon short of an octet until lone pairs are converted into multiple bonds. Formal-charge accounting supplies a third check in the later lessons.