Structure & basicity — Amines

Amines are the most reliably basic functional group in organic chemistry, and almost everything about them follows from one question: how available is the nitrogen lone pair? That single question sets basicity, nucleophilicity, and whether the nitrogen can act as a leaving group — and the answers span more than ten orders of magnitude.

Nitrogen's lone pair

An amine nitrogen is sp³ hybridized with three bonds and a lone pair, giving it a trigonal pyramidal shape with bond angles near 107°. That lone pair makes it a Brønsted base, able to accept a proton, and a Lewis base, able to donate a pair to any electrophile — the two descriptions are of the same property.

A typical alkylamine's conjugate acid, R–NH₃⁺, has a pKa around 10–11. That makes amines vastly more basic than water (conjugate acid pKa −1.7) or an alcohol. In practical terms it means an amine is protonated at physiological pH, which is why most amine-containing drugs are formulated and administered as their hydrochloride salts: the salt is water-soluble where the free base is not.

An amine nitrogen with three different groups is not a resolvable stereocenter. It inverts through a planar transition state millions of times per second — nitrogen inversion — so the two forms cannot be separated. A quaternary ammonium salt, with four groups and no lone pair, cannot invert and is a genuine stereocenter.

Alkyl substitution: a small and complicated effect

You would expect more alkyl groups to donate more electron density inductively and make the amine progressively more basic. In the gas phase, that is exactly what happens: NH₃ < MeNH₂ < Me₂NH < Me₃N, cleanly.

In water the ordering breaks down. A more substituted ammonium ion has fewer N–H bonds available to hydrogen-bond with water, so it is less well stabilized by solvation — and that effect works against the inductive one. The measured aqueous pKaH values are methylamine 10.6, dimethylamine 10.7, trimethylamine 9.8: a rise, then a fall.

This is worth more than its exam weight, because it is a clean demonstration that solvation can compete with and even reverse an intrinsic electronic effect. The gas-phase answer is the "pure" one and the aqueous answer is the one you measure.

Despite the complication, every simple alkylamine — primary, secondary or tertiary — clusters in the same pKaH range of roughly 9–11. The differences among them are small compared with the effect described next, which is worth six orders of magnitude.

Resonance donation: the dramatic basicity killer

Basicity in amines is one question asked three times: is the nitrogen lone pair still there to donate? A lone pair that is delocalized somewhere else is not available to pick up a proton, and every drop in the row below is that and nothing else.The amide is the case worth remembering, because the collapse is enormous — around ten orders of magnitude from an ordinary alkylamine. The lone pair is conjugated into the carbonyl and spends its time on oxygen, which is also why an amide C–N bond is short, planar and does not rotate freely. An amide nitrogen is not a weak base; it is not usefully a base at all.

Amine basicity is not a list to memorise — it is one question asked five times: how busy is the lone pair? Where nothing competes for it, the nitrogen is a good base. Where resonance can pull it into a ring or a carbonyl, it is partly or wholly spent before a proton ever arrives, and the basicity collapses by orders of magnitude. Pyridine is the instructive case: it is fully aromatic and still a decent base, because its lone pair sits in the ring plane and was never part of the sextet.Basicity is one question: is the lone pair free to take a proton, or is it busy? Alkyl groups nudge it slightly one way; RESONANCE moves it by orders of magnitude, because a delocalised pair is a pair already spent. Pyridine is the case worth understanding: aromatic, yet still basic, because its lone pair was never in the ring.

Compare aniline, where the nitrogen is attached to a benzene ring, with cyclohexylamine, where it is attached to a saturated one. Cyclohexylamine has pKaH 10.7; aniline has pKaH 4.6. Aniline is about a million times less basic.

The reason is that aniline's lone pair is delocalized into the aromatic ring — it is part of the pi system, not sitting available on nitrogen. Protonating it would destroy that delocalization, which costs energy, so the equilibrium shifts away from the protonated form.

An amide is the extreme case. Its nitrogen lone pair is so heavily delocalized into the adjacent carbonyl that the conjugate acid's pKa is around −1: essentially non-basic. This is the same delocalization that makes an amide planar, that makes amide nitrogen a hopeless leaving group, and that makes peptide bonds stable. One structural feature, four consequences.

Compound	pKaH	Lone pair
alkylamine	10–11	fully available
ammonia	9.2	available, no donors
pyridine	5.2	available but sp²
aniline	4.6	into the ring
amide	~−1	into the carbonyl
pyrrole	~−4	part of the aromatic sextet

Two nitrogen heterocycles worth contrasting

Pyridine (pKaH 5.2) has its lone pair in an sp² orbital in the plane of the ring, pointing outward — not part of the aromatic pi system at all. It is available, but sp² orbitals hold electrons more tightly than sp³ ones, so pyridine is a weaker base than an alkylamine. This is the Orbital factor from Module 3 acting on basicity.

Pyrrole (pKaH about −4) is different in kind. Its lone pair is the two electrons that make the ring aromatic, contributing to the six-electron sextet. Protonating the nitrogen would destroy the aromaticity entirely, at a cost of some 20 kcal/mol, so it simply does not happen — pyrrole is protonated on carbon instead, if at all.

Same element, same ring size, two nitrogens differing only in whether the lone pair is in the pi system. Nine orders of magnitude apart.

Worked example — ranking four nitrogens

Rank by basicity: p-nitroaniline, aniline, p-methoxyaniline, cyclohexylamine.

Cyclohexylamine is clearly most basic — no delocalization at all (10.7). The three anilines all delocalize into a ring, so they sit far lower, and the substituent then fine-tunes.

p-Methoxyaniline (5.3) is the most basic of the three, since the donating methoxy pushes density back toward nitrogen. Aniline is 4.6. p-Nitroaniline (1.0) is the least, since the nitro group withdraws strongly and pulls the lone pair further away.

Order: cyclohexylamine > p-methoxyaniline > aniline > p-nitroaniline. Resonance sets the tiers; substituents order within them — exactly the pattern from the carboxylic acids section.

The question to ask

Whenever you need to rank two nitrogen-containing compounds, ask first: is the lone pair delocalized, and if so into what? Into a carbonyl is worst, into an aromatic sextet is worse still, into a ring is bad, and not at all is best. Hybridization and induction then break ties. Almost every basicity question in this module is answered by that one ordering.

What carries forward

The availability of the lone pair determines everything in the next section — amines are nucleophiles when the lone pair is free and unreactive when it is not, which is the basis of the acylation strategy. The same delocalization argument returns in Module 13, where an amino group on a ring is one of the strongest activating substituents precisely because it donates so well.