Past the proton: a wider definition
In the previous guide you met the [[inorg-bronsted-lowry-acid-base|Brønsted–Lowry]] picture: an acid donates a proton, H+, and a base accepts it. That view is clean and it explains an enormous amount of aqueous chemistry. But notice a quiet oddity — the Brønsted base is the active partner. When NH3 grabs a proton, what actually does the work is the nitrogen lone pair reaching out and seizing the bare H+. The proton is just a convenient passenger. G. N. Lewis, the same chemist behind the dot structures from the earlier rung, asked the natural follow-up question: what if the proton isn't essential at all, and the real event is simply one species offering an electron pair and another accepting it?
That question is the whole of the [[inorg-lewis-acid-base|Lewis definition]]. A Lewis base is an electron-pair donor — it must carry a lone pair, or sometimes a bonding pair, that it can offer. A Lewis acid is an electron-pair acceptor — it must have somewhere to put that pair, typically a low-lying empty orbital. The two meet and a new covalent bond forms, made entirely from electrons that started out on the base. No proton need ever appear. The proton, in this light, becomes just one Lewis acid among thousands — a tiny, bare, hungry orbital that happens to be very common in water.
The adduct: where acid and base meet
When a Lewis acid and a Lewis base join, the product has a name worth remembering: an [[acid-base-adduct|adduct]]. The classic demonstration is boron trifluoride meeting ammonia. You met BF3 back in the bonding rung as an electron-deficient molecule — boron sits at the centre of three B-F bonds with only six electrons around it and one empty 2p orbital pointing straight up out of the trigonal plane. Ammonia has the opposite problem: a nitrogen with a fat lone pair and nowhere interesting to put it. Bring them together and the nitrogen lone pair slides directly into boron's empty orbital, forming the adduct written F3B-NH3.
Watch what happens to boron's shape as the bond forms. Free BF3 is flat — trigonal planar, all three fluorines and the boron in one plane, leaving that empty 2p orbital sticking out above and below. The moment nitrogen's pair moves in, boron rehybridises from sp2 to sp3 and the molecule puckers up into a tetrahedron, the three fluorines bending down like an umbrella opening. The new B-N bond is an ordinary covalent bond in every measurable way; the only unusual thing is its origin — both electrons came from a single partner. Older books call this a 'dative' or 'coordinate' bond, but once it has formed, an electron cannot remember where it came from.
F F F F H \ / \ | / B : N--H --> B --------N--H / \ / | F F F H trigonal pyramidal adduct: F3B-NH3 planar base (B is now sp3, tetrahedral) sp2, empty 2p orbital
A field guide to the players
Once you know what to look for, Lewis acids and bases jump out everywhere. The classic Lewis acids fall into three big families. First, the electron-deficient covalent molecules — the [[boron-trihalide-lewis-acids|boron trihalides]] BF3 and BCl3, with their empty p orbital. Second, molecules whose central atom can be coaxed to accept more, like AlCl3: in the gas phase two AlCl3 units pair up into the dimer Al2Cl6, each aluminium acting as a Lewis acid toward a chlorine lone pair on its neighbour, forming two bridging Al-Cl-Al bonds — the aluminium chloride dimer you may already have glimpsed. Third, and most important for inorganic chemistry, the metal cations.
A bare metal cation like Fe3+ or Co3+ is a tiny ball of positive charge with empty valence orbitals — a superb electron-pair acceptor. The Lewis bases that surround it are called ligands, and the [[donor-atom|donor atom]] is the particular atom on the ligand that hands over its lone pair. The common Lewis bases are exactly the molecules you would guess: amines and ammonia (donating through nitrogen), ethers and water and alcohols (through oxygen), and the halide ions F-, Cl-, Br-, I- (through the halogen). When six ammonia molecules each push a nitrogen lone pair into a Co3+ ion, you get the complex [Co(NH3)6]3+ — and that, exactly, is the bridge into the next major topic.
How to spot which is which
Faced with an unfamiliar reaction, you can usually label the Lewis acid and base in a few seconds. The trick is to look for who has the pair and who has the hole.
- Find the lone pairs. Any atom with a visible lone pair — the N in an amine, the O in water or an ether, a halide ion — is a candidate Lewis base. If it has no lone pair to give, it cannot be the base.
- Find the empty orbital. Look for an electron-deficient centre (boron in BF3), a positively charged metal ion (Fe3+, Al3+), or a centre that can shuffle to make room. That is your Lewis acid.
- Draw the arrow. The new bond always points from the base's lone pair toward the acid's empty orbital. Both electrons originate on the base; the acid contributes only the empty space.
- Sanity-check the count. The donor atom loses 'sole ownership' of its pair but gains a bond; the acceptor gains a share of two electrons. No electrons are created or destroyed — they are just shared in a new place.
Try it on a reaction you already know from the Brønsted guide. When ammonia dissolves in water and a proton hops across to give NH4+, the Lewis reading is: the nitrogen lone pair (base) attacks the proton of water, which is the smallest possible Lewis acid. The proton's role shrinks from 'the thing being transferred' to merely 'a bare 1s acceptor orbital with a positive charge'. Same event, deeper lens — and now you can see it as one tiny example of the same handshake that builds [Co(NH3)6]3+.
Honest limits, and where this leads
The Lewis definition is powerful precisely because it is broad, but breadth has a cost: it tells you that a reaction can happen, not how strongly or how fast. Saying 'metal cations are Lewis acids' does not tell you that Fe3+ binds fluoride hungrily while Hg2+ prefers iodide and almost ignores fluoride. That extra layer of prediction is the [[hsab-principle|hard–soft acid–base]] idea you will meet in the very next guide: small, high-charge, non-polarisable acids and bases ('hard') pair best with each other, while big, polarisable ones ('soft') pair best together. Lewis tells you the game; hard-soft tells you who wins.
A second honest caveat: 'Lewis acid' is a role a species plays, not a fixed label stamped on it for life. Water is a Lewis base toward Fe3+ (donating an oxygen lone pair) yet acts as a Lewis acid toward a stray proton it picks up. Many species play both parts depending on their partner. And remember that the dative bond in an adduct, despite its special origin story, is an ordinary covalent bond once made — do not imagine it as weaker or somehow second-class just because both electrons came from one side.
Step back and see what you have gained. The Brønsted picture explained acids in water; the Lewis picture explains acids everywhere — in molten salts, in non-aqueous solvents, in the gas phase, on a catalyst surface, inside an enzyme's active site. It links acid-base chemistry, the formation of every metal complex, and a huge slice of reaction mechanism into one idea: electrons flow from where they are abundant to where there is room. That single donor-acceptor sentence is the thread you will follow through the rest of this ladder.