Lewis Acids and Bases: Electron Pairs, Not Protons

Boron trifluoride has no hydrogen atoms at all. It can’t donate a proton, and it doesn’t make hydroxide ions. Yet it reacts eagerly with ammonia, forming a stable compound and releasing heat, in exactly the way an acid reacts with a base. Under the Arrhenius and Brønsted–Lowry definitions, this reaction is impossible to classify.

Gilbert N. Lewis solved the puzzle in 1923 (the same year Brønsted and Lowry published) by asking a different question. Instead of following the proton, he followed the electrons.

The Lewis definitions

• A Lewis acid is an electron-pair acceptor.
• A Lewis base is an electron-pair donor.

When a Lewis acid meets a Lewis base, the base shares one of its lone pairs with the acid and a new covalent bond forms. Both electrons in that bond come from the base. This kind of bond is called a dative or coordinate bond, and the product is called an adduct (or a complex).

The classic example: BF₃ and NH₃

BF₃ + :NH₃ → F₃B←NH₃

In BF₃, the boron atom has only six electrons in its outer shell: three bonding pairs, no lone pairs. It has an empty orbital and is “electron-deficient”. Nitrogen in ammonia has a lone pair.

The nitrogen donates its lone pair into boron’s empty orbital. A B–N bond forms, both boron and nitrogen end up with eight electrons around them, and the shape at boron changes from flat trigonal planar to tetrahedral.

• BF₃ accepts the pair → Lewis acid
• NH₃ donates the pair → Lewis base

No protons were transferred. This is acid–base chemistry based purely on electrons.

How Lewis includes the other definitions

The Lewis definition is the broadest of the three. Every Brønsted–Lowry reaction is also a Lewis reaction, when you look at the electrons.

Take ammonia accepting a proton:

H⁺ + :NH₃ → NH₄⁺

The H⁺ has an empty 1s orbital. Ammonia’s lone pair forms the new N–H bond. So H⁺ is a Lewis acid and NH₃ is a Lewis base.

Or hydroxide neutralising an acid:

H⁺ + :OH⁻ → H₂O

Hydroxide donates a lone pair from oxygen to the proton. Again, OH⁻ is a Lewis base.

Definition	Acid	Base	Scope
Arrhenius	Gives H⁺ in water	Gives OH⁻ in water	Aqueous only
Brønsted–Lowry	Proton donor	Proton acceptor	Any reaction with proton transfer
Lewis	Electron-pair acceptor	Electron-pair donor	Any reaction forming a dative bond

One subtlety: under Lewis’s definition, the Brønsted acid HCl isn’t itself the Lewis acid. The Lewis acid is the proton it hands over. In practice, chemists keep using “acid” for HCl in the Brønsted sense and use “Lewis acid” mainly for electron-pair acceptors that don’t involve protons.

Spotting Lewis acids

Lewis acids have somewhere to put an extra pair of electrons. Common types:

1. Molecules with an incomplete octet. BF₃, BCl₃, AlCl₃. Boron and aluminium compounds with three bonds have only six outer electrons.

2. Metal cations. Fe³⁺, Al³⁺, Cu²⁺, Zn²⁺, Ag⁺ and others have empty orbitals and attract lone pairs. When a metal ion dissolves in water, water molecules bond to it through their oxygen lone pairs, forming ions like [Fe(H₂O)₆]³⁺. The metal ion is a Lewis acid and each water molecule is a Lewis base.

3. Molecules with polar multiple bonds. In CO₂, the carbon is electron-poor. When CO₂ meets hydroxide, OH⁻ donates a pair to carbon and forms hydrogencarbonate:

CO₂ + OH⁻ → HCO₃⁻

4. Atoms that can expand their valence shell. SiF₄ accepts two fluoride ions to form SiF₆²⁻. SO₃ acts as a Lewis acid towards water in forming sulfuric acid.

5. The proton itself. H⁺ is the simplest Lewis acid of all.

Spotting Lewis bases

Lewis bases have a lone pair (or occasionally a π bond) available to donate:

• Neutral molecules with lone pairs: H₂O, NH₃, amines, ethers, alcohols, carbon monoxide.
• Anions: OH⁻, F⁻, Cl⁻, CN⁻, and any other negative ion with lone pairs.
• Alkenes and aromatic rings can donate π electrons in some reactions.

Every Brønsted base is a Lewis base, because to accept a proton a particle needs a lone pair to bond to it.

Where Lewis acid–base chemistry shows up

Coordination complexes. Transition metal complexes are Lewis adducts. In [Cu(NH₃)₄(H₂O)₂]²⁺, the deep-blue ion that forms when ammonia is added to copper(II) sulfate solution, the copper ion is the Lewis acid and the ammonia and water molecules (called ligands) are Lewis bases. Haemoglobin carries oxygen using the same idea: Fe²⁺ at the centre of each haem group acts as a Lewis acid.

Organic chemistry. Many reactions use Lewis acid catalysts. In Friedel–Crafts reactions, aluminium chloride accepts a lone pair from a chlorine atom, generating a highly reactive carbocation that attacks a benzene ring. Understanding nucleophiles (Lewis bases) and electrophiles (Lewis acids) is the key to organic reaction mechanisms.

Acidic metal ions in water. Solutions of iron(III) chloride or aluminium sulfate are acidic even though the salts contain no acidic hydrogen. The highly charged metal ion pulls electron density from its bound water molecules, weakening the O–H bonds so that a proton is released:

[Fe(H₂O)₆]³⁺ + H₂O ⇌ [Fe(H₂O)₅(OH)]²⁺ + H₃O⁺

This is why iron(III) salt solutions have a pH of around 2–3 and why aluminium sulfate is used to lower soil pH for blue hydrangeas.

Industrial catalysis. Zeolites and solid acid catalysts used in oil refining rely on Lewis acid sites to crack large hydrocarbons into petrol-sized molecules.

Hard and soft acids and bases

Once you think in terms of electron pairs, you can ask which acids prefer which bases. Ralph Pearson’s HSAB principle (1963) sorts them:

• Hard acids and bases are small, highly charged and not easily polarised: H⁺, Na⁺, Mg²⁺, Al³⁺ (acids); F⁻, OH⁻, H₂O, NH₃ (bases).
• Soft acids and bases are large, less highly charged and easily polarised: Ag⁺, Hg²⁺, Pt²⁺ (acids); I⁻, S²⁻, CN⁻, CO (bases).

The rule of thumb: hard likes hard, soft likes soft. It explains why mercury and lead occur in nature as sulfide ores (soft acid + soft base), while aluminium and magnesium occur as oxides and silicates (hard + hard).

Worked examples

Classify the Lewis acid and base in each reaction:

1. AlCl₃ + Cl⁻ → AlCl₄⁻ Acid: AlCl₃ (aluminium has an empty orbital). Base: Cl⁻.

2. Ag⁺ + 2NH₃ → [Ag(NH₃)₂]⁺ Acid: Ag⁺. Base: NH₃ (two molecules each donate a lone pair).

3. CO₂ + H₂O → H₂CO₃ Acid: CO₂ (carbon accepts a pair). Base: H₂O (oxygen lone pair).

4. Ni + 4CO → Ni(CO)₄ Acid: Ni atom. Base: CO (lone pair on carbon).

Key takeaways

• Lewis acids accept electron pairs; Lewis bases donate them. The product contains a dative covalent bond.
• The Lewis definition includes all Brønsted–Lowry reactions and many more that involve no protons.
• Typical Lewis acids: electron-deficient compounds (BF₃, AlCl₃), metal cations and H⁺. Typical Lewis bases: anything with a lone pair.
• Lewis acid–base ideas underpin coordination chemistry, organic mechanisms and catalysis.
• For the bigger picture, see acid–base definitions and Brønsted–Lowry acids and bases.