= Hydrogen bond acceptor
A <hydrogen bond acceptor> supplies the electron-rich atom or group $Y$ interacting with the attached <hydrogen> of a <hydrogen bond donor> in $X-\mathrm H\cdots Y$. An <oxygen> <lone pair> in <water> is the familiar example. Acceptance describes the association, not necessarily acquisition of a <proton>. A single <molecule>, such as <water> or <ammonia>, can offer both <hydrogen bond donor> and <hydrogen bond acceptor> sites. See https://people.reed.edu/~glasfeld/Chem391/notes/ChIMF_2018.pdf[Glasfeld: water and nitrogen-containing accepting sites].
A <lone pair> drawn in a <Lewis structure> is a useful clue, not a universal test. At an ordinary neutral <amide> <nitrogen>, the nominal <lone pair> participates in <chemical resonance> with the <carbonyl group>, so this <nitrogen> is ordinarily excluded from lists of <hydrogen bond acceptors>; the <carbonyl group>'s <oxygen> remains an accepting site. <Protonation> changes the test again: $\mathrm{NH}_4^+$ has four N–H <covalent bonds> and no <nitrogen> <lone pair>, whereas <ammonia> has three and one. These distinctions follow the actual chemical form, and accepting ability also depends on the partner and whether it can approach the site. See https://openstax.org/books/organic-chemistry/pages/24-3-basicity-of-amines[McMurry: amide resonance and protonation] and https://people.reed.edu/~glasfeld/Chem391/notes/ChIMF_2018.pdf[Glasfeld: limitations of nitrogen accepting sites].
Conversely, an accepting group need not supply an atom-centred <lone pair>. The following <phenol>–<benzene> experiment probes an O–H donor interacting with the ring's distributed $\pi$ <electron density>.
\Image[/hydrogen-bond-phenol-infrared-controls.png]
{title=<Infrared spectroscopy> distinguishes free and associated <phenol>: schematic experimental logic, no plotted or simulated spectra}
{description=Original toy layout of the comparisons reported by Saggu, Levinson and Boxer (2011). The connector identifies an O–H donor and a ring accepting region; it does not specify a measured <bond length> or the calculated contact geometry.}
Saggu, Levinson and Boxer (2011) used <infrared spectroscopy> to compare $200\,\mathrm{mmol\,L^{-1}}$ <phenol> in different <solvents>. The O–H stretching peak was at $3611\,\mathrm{cm^{-1}}$ in $\mathrm{CCl}_4$ and $3555.9\,\mathrm{cm^{-1}}$ in <benzene>. To distinguish association from a general <solvent> effect, they diluted the ring-containing <chemical solutions> with $\mathrm{CCl}_4$: two peaks appeared, with the free peak nearly fixed while the associated peak depended on the partner. Replacing the O–H <hydrogen> with <deuterium> reproduced the ordering of shifts for different partners. Thus the perturbed O–H group identifies the <hydrogen bond donor>; the ring supplies the <hydrogen bond acceptor>. A peak shift alone does not specify a unique <molecular geometry>. https://pmc.ncbi.nlm.nih.gov/articles/PMC3202663/[Saggu, Levinson and Boxer, Direct measurements of electric fields in weak OH···π hydrogen bonds, JACS 133 (2011), Figure 1, Table 1 and dilution controls]{ref}
The roles make recognition between <molecules> intelligible: exposed <hydrogen bond donors> need compatible <hydrogen bond acceptors> in reachable positions. A count of sites is a starting description, not a prediction of binding strength. In <water>, forming a contact between two partners can also replace their existing <hydrogen bonds> to <water>; the <Gibbs free energy> of association concerns the whole change. See https://people.reed.edu/~glasfeld/Chem391/notes/ChIMF_2018.pdf[Glasfeld, pp. 15–16: competition with solvent interactions].
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