= Lone pair
{wiki}
A <lone pair> is an <electron pair> assigned to one <chemical atom> without a connecting line to another <chemical atom> in a <Lewis structure>. It is usually shown as two dots. <Water> has two <lone pairs> on <oxygen>, accounting for four <valence electrons>; <Ammonia> has one <lone pair> on <nitrogen>, accounting for two. In $\mathrm{O=C=O}$, each <oxygen> has two <lone pairs>, and <carbon> has none.
A <lone pair> can be reactive. In <protonation> of <Ammonia>, the <nitrogen> <lone pair> becomes a new <bond pair> to a <proton>:
$$
\mathrm{NH}_3+\mathrm H^+\longrightarrow\mathrm{NH}_4^+.
$$
The <proton> supplies no <electron>, so the eight <valence electrons> are retained: three <bond pairs> plus one <lone pair> become four <bond pairs>. The resulting <nitrogen> has <formal charge> $5-0-8/2=+1$; each <hydrogen> has <formal charge> zero. This is why identifying an available <lone pair> helps explain the <chemical base> behaviour of <Ammonia>, rather than merely decorating its formula.
Lewis already used the <Ammonia> to $\mathrm{NH}_4^+$ example in https://zenodo.org/records/1429068[The Atom and the Molecule (1916), pp. 777–778]. In modern terms the dots identify an allocation of <electrons>, not two stationary particles or an isolated patch unaffected by the rest of the <molecule>. A <molecular orbital> description can spread electronic amplitude over several <chemical atoms> while a localized <Lewis structure> remains useful for <acid-base chemistry>.
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