Lewis acids and bases

Key structures for this topic — drawn live.

Beyond protons: the electron-pair view

The Brønsted definition is all about protons, but many reactions involve no proton transfer at all. The Lewis definition is more general: it describes acids and bases in terms of electron pairs instead of hydrogen ions. A Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. Because it focuses on where electrons go, this picture underlies almost every mechanism you'll draw in organic chemistry.

Lewis acids: the electron-pair acceptors

A Lewis acid needs somewhere to put an incoming pair of electrons — that is, an empty orbital or the ability to accept electrons into a low-lying orbital. Common examples:

• Electron-deficient boron and aluminum compounds like BF₃ and AlCl₃, whose central atom has only six valence electrons and an empty p orbital.
• H⁺, a bare proton with a completely empty 1s orbital (so every Brønsted acid is also a Lewis acid).
• Carbocations, where a positively charged carbon has an empty p orbital.
• Metal cations such as Mg²⁺, Zn²⁺, and Fe³⁺, which accept lone pairs from ligands.

These species are electron-hungry — often positively charged or having an incomplete octet — which is exactly what makes them accept electron density.

Lewis bases: the electron-pair donors

A Lewis base must have electrons available to give away — a lone pair or, in many organic cases, the electrons of a π bond. Examples include:

• Molecules with lone pairs: NH₃, H₂O, ethers, amines, halide ions (Cl⁻), hydroxide (OH⁻).
• π systems: the C=C double bond of an alkene can donate its π electrons to a strong Lewis acid.

Every Brønsted base is also a Lewis base, since accepting a proton requires donating a lone pair to it. But the Lewis definition is broader — an alkene has no proton to give or take, yet it can act as a Lewis base toward an electrophile.

The curved arrow: base to acid

The single most useful mechanistic habit is this: in a Lewis acid–base reaction, the curved arrow starts at the Lewis base's electron pair and points to the Lewis acid, forming a new covalent bond between them. The product of this combination is called a Lewis adduct (or acid–base complex).

For example, when NH₃ (base) reacts with BF₃ (acid), the nitrogen lone pair attacks boron: the arrow goes from N to B, forming an N–B bond and giving H₃N–BF₃. Nitrogen picks up a formal positive charge and boron a formal negative charge, but both now have full octets. No proton moved — this is pure electron-pair chemistry.

The nucleophile–electrophile connection

Here's why this matters for the rest of the course: a Lewis base is a nucleophile and a Lewis acid is an electrophile. They describe the same electron-rich and electron-poor species from slightly different angles.

• Lewis base / nucleophile: "nucleus-loving," electron-rich, donates electrons, and attacks positive centers.
• Lewis acid / electrophile: "electron-loving," electron-poor, accepts electrons, and is attacked by nucleophiles.

The subtle distinction is usually one of emphasis: chemists say "Lewis acid/base" when talking about acid–base equilibria and adduct formation, and "electrophile/nucleophile" when talking about the bond-forming step of a reaction mechanism. Under the hood it's the same electron flow.

Why this is the most general picture

The three acid–base theories nest inside each other. Arrhenius is narrowest (acids make H⁺, bases make OH⁻ in water). Brønsted–Lowry is broader (any proton transfer, in any solvent). Lewis is broadest of all — every Brønsted acid–base reaction is also a Lewis one, but not vice versa. Because the Lewis view tracks electron pairs rather than protons, it explains reactions of metal ions, boron and aluminum reagents, carbocations, and π bonds that the proton-based definitions simply can't. When you learn to spot the Lewis acid and Lewis base in a reaction — and to draw the arrow from base to acid — you have the master key to organic mechanisms.