= Octet rule
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The <octet rule> is a useful empirical guide: many <Lewis structures> of <chemical compounds> containing <carbon>, <nitrogen>, and <oxygen> give each of those <chemical atoms> eight <electrons> when nonbonding <electrons> and all incident bonding <electrons> are counted. <Hydrogen> instead normally counts two <electrons>. These are local counts, not independent budgets: a shared <bond pair> counts at both endpoints but only once in the <molecule>'s <valence electron> total.
In <water>, <oxygen> counts $4+4=8$ <electrons>; in <Ammonia>, <nitrogen> counts $2+6=8$; in $\mathrm{O=C=O}$, <carbon> counts $0+8=8$ and each <oxygen> counts $4+4=8$. The charge-separated <carbon dioxide> drawing also satisfies the <octet rule>, so the <octet rule> alone cannot select the dominant contributor. <Formal charge> provides another check.
The <octet rule> is a heuristic rather than a theorem about every <chemical bond>. Odd <valence electron> totals cannot be partitioned entirely into <electron pairs>, and some valid <Lewis structures> have fewer or more than eight <electrons> around a <chemical atom>. See https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures[Chemistry 2e, 7.3: exceptions to the octet rule]. The capacity of the $s$ and $p$ <electron subshells>, $2+6=8$, helps motivate the guide for these <chemical elements>, but does not prove universal compliance. Lewis discussed a stable outer group of eight, together with a two-<electron> arrangement for <hydrogen>, in https://zenodo.org/records/1429068[The Atom and the Molecule (1916), pp. 768 and 774].
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