Chemistry: Atoms First 2e · Equilibria of Other Reaction Classes

Lewis Acids and Bases

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

The Arrhenius and Brønsted–Lowry definitions tie acid–base behavior to protons, but many clearly acid–base reactions involve no H⁺ at all. The Lewis definition focuses on electrons: a is any species that can accept an electron pair, and a is any species that can donate one. The shared pair forms a , and the product is an — or, when a metal ion is involved, a . This one idea unifies proton transfer, the acidic behavior of electron-deficient molecules like BF₃ and AlCl₃, the acidity of hydrated metal ions, and all of coordination chemistry. Because the Lewis definition is the most general of the three, every Brønsted–Lowry acid–base reaction is also a Lewis reaction — but many Lewis reactions involve no proton transfer at all.

Why this matters

Lewis acid–base chemistry explains reactions the other definitions cannot touch. Hemoglobin binds O₂ because the Fe²⁺ center is a Lewis acid accepting O₂'s lone pair; vitamin B₁₂ centers on a cobalt– complex; enzymes use metal-ion cofactors as Lewis acids to polarize substrates. Chelation therapy for heavy-metal poisoning uses EDTA — a six-toothed Lewis base — to bind Pb²⁺ or Hg²⁺ and escort it out of the body. In industry, AlCl₃ and BF₃ are Lewis-acid catalysts for Friedel–Crafts reactions. In the lab, metal ions are Lewis acids, which is why solutions of salts like FeCl₃ are acidic even though no H⁺ was added. The formation constant K_f gives this chemistry a quantitative handle — the same equilibrium mathematics as the K_sp of the previous topic.

The college version

Core Concepts

The Lewis definitions

Lewis acid: electron-pair acceptor — H⁺, BF₃, AlCl₃, Fe³⁺, Ag⁺, SO₃.
Lewis base: electron-pair donor — NH₃, H₂O, OH^-, CN^-, halide ions, any lone-pair species.

In BF₃ + NH₃ → F₃B-NH₃, the nitrogen lone pair fills the empty orbital on boron, forming a new B–N bond. Both electrons come from the base, which is why the bond is called coordinate covalent — though once formed, it is identical to an ordinary covalent bond.

How Lewis theory relates to the other definitions

The definitions nest: every Arrhenius base is a Brønsted–Lowry base, and every Brønsted–Lowry base is a Lewis base. Every Brønsted acid is a Lewis acid, because H⁺ must accept an electron pair from the base. But the Lewis definition includes species that never transfer a proton — BF₃ is a Lewis acid that is not a Brønsted acid, and AlCl₃ contains no hydrogen at all. When H₃O⁺ forms, you are watching a Lewis reaction: H⁺ (acid) accepts the lone pair of H₂O (base).

Recognizing Lewis acids and bases from electron structure

A species with an empty orbital and fewer than an octet is a Lewis acid: BF₃ (boron has only 6 valence electrons), AlCl₃, H⁺ (no electrons), and transition-metal cations (vacant orbitals).
A species with an available lone pair is a Lewis base: NH₃, H₂O, OH^-, CN^-, Cl^-.

Ligands are Lewis bases by definition; the metal cation is the Lewis acid. In [Ag(NH₃)₂]⁺, two ammonia molecules donate lone pairs to Ag⁺, and the coordination number is 2.

Metal ions in water: why salt solutions can be acidic

A hydrated metal cation like Fe(H₂O)₆³⁺ is a Lewis acid: the highly charged Fe³⁺ pulls electron density away from the O–H bonds of the bound waters, making those protons easier to release:

$Fe(H 2 O) 6 3+ (aq) + H 2 O (l) ⇌ Fe(H 2 O) 5 (OH) 2+ (aq) + H 3 O + (aq)$

The stronger the Lewis acidity of the cation (higher charge, smaller radius), the more acidic the solution — which is why FeCl₃(aq) and AlCl₃(aq) are acidic.

Formation constants K_f

Complex formation is an equilibrium:

$Ag + + 2 NH 3 ⇌ [Ag(NH 3) 2] + K f = [[Ag(NH 3) 2] +] [Ag +][NH 3] 2 = 1.7 \times 107$

Large K_f values (often 10¹⁰–10³⁰ for chelating ligands) mean the metal is almost entirely tied up in the complex — which is exactly why complexation can dissolve precipitates, the subject of the next topic.

How It Works / Step-by-Step Process

Classifying a reaction partner as Lewis acid or base

Draw the Lewis structure of each species.
Look for an empty orbital or incomplete octet (acid) versus an available lone pair (base).
Confirm by writing the electron-pair movement: a lone pair from the base forms the new bond to the acid.

Using K_f to find free metal-ion concentration

Write the complex-formation equilibrium and its K_f expression.
When K_f is large, assume near-complete complexation; compute the complex concentration and the excess ligand concentration.
Solve the K_f expression for the tiny free metal-ion concentration.

Common Confusions

Do Not Confuse	With	The Difference
Lewis acid	Brønsted–Lowry acid	Lewis acids accept electron pairs; they need not contain or donate H⁺ (BF₃, AlCl₃, metal cations)
Lewis base	Brønsted–Lowry base	All Brønsted bases are Lewis bases, but any lone-pair donor (e.g., CN⁻ binding a metal) is a Lewis base without proton chemistry
Coordinate covalent bond	A special, weaker bond	Once formed, a coordinate bond is identical to any covalent bond; the label only records where the electrons came from
Acid = donor / base = acceptor	Acid = acceptor / base = donor	Lewis acid accepts the pair; Lewis base donates it. Getting this backwards fails every classification
"No H⁺, so not acid–base"	Lewis acid–base chemistry	Proton-free reactions like BF₃ + NH₃ are acid–base reactions under the Lewis definition
Complex ion	Precipitate	Complex ions are soluble, charged species held by coordinate bonds; precipitates are solids governed by K_sp

Eli explains

The same idea, in plain words

Explain it like I’m 10

A Lewis acid is a hungry catcher and a Lewis base is a ball-thrower: the base tosses a pair of electrons and the acid catches them to make a new bond. Some catchers have no protons to give — like a bouncer with an empty hand — but they still catch electrons, which is why this rule covers many more reactions than acids that simply give protons.

Worked example

Example 1: Classifying and drawing the BF₃ + NH₃ reaction

Boron has 3 valence electrons; in BF₃ it forms three single bonds and holds no lone pair, so boron is surrounded by only 6 electrons — a sextet, one pair short of an octet. BF₃ is electron-poor: a Lewis acid with an empty p orbital. Ammonia's nitrogen has a lone pair, making NH₃ a Lewis base. The reaction:

$BF 3 + NH 3 \to F 3 B - NH 3$

The ammonia lone pair fills the empty orbital on boron, completing boron's octet with a coordinate covalent bond. This electron-poor behavior makes BF₃ (and AlCl₃) useful Lewis-acid catalysts. Note that BF₃ never donates or accepts a proton, so it is not a Brønsted acid — only the Lewis definition covers it.

Example 2: How much free Ag⁺ remains in an ammonia solution?

The formation constant for Ag⁺ + 2 NH₃ ⇌ [Ag(NH₃)₂]⁺ is K_f = 1.7 × 10⁷. Suppose 0.010 mol of Ag⁺ is added to 1.0 L of 1.0 M NH₃. Find the free [Ag⁺] at equilibrium.

Because K_f is very large, assume the complexation goes essentially to completion: [[Ag(NH₃)₂]⁺] ≈ 0.010 M. Each complex consumes 2 NH₃, so:

$[NH 3] eq = 1.0 - 2(0.010) = 0.98 M$

Solve the K_f expression for the free silver:

$[Ag +] = [[Ag(NH 3) 2] +] K f [NH 3] 2 = 0.010 (1.7 \times 107)(0.98) 2 = 0.010 1.63 \times 107 = 6.1 \times 10 -10 M$

Free Ag⁺ falls to roughly 6 × 10^-10 M — more than seven orders of magnitude below the total silver. This near-total sequestration is why ammonia can dissolve silver chloride (the next topic's coupled-equilibria idea).

Example 3: Why a solution of FeCl₃ is acidic

Iron(III) chloride dissolves to give hydrated Fe(H₂O)₆³⁺ ions. The small, highly charged Fe³⁺ center is a Lewis acid that withdraws electron density from the O–H bonds of its six bound waters, weakening them:

$Fe(H 2 O) 6 3+ (aq) + H 2 O (l) ⇌ Fe(H 2 O) 5 (OH) 2+ (aq) + H 3 O + (aq)$

The released hydronium makes the solution acidic even though the salt contains no "obviously acidic" hydrogen. The same reasoning explains the acidity of AlCl₃, CuSO₄, and CrCl₃ solutions: the smaller and more highly charged the cation, the stronger its Lewis acidity.

Key takeaways

Lewis acid = electron-pair acceptor; Lewis base = electron-pair donor.
Every Brønsted–Lowry acid–base reaction is a Lewis reaction; the reverse is false (BF₃, AlCl₃ are Lewis acids with no H⁺).
The donated pair forms a coordinate covalent bond; the product is an adduct or complex ion.
Electron-poor species (incomplete octet, cations, H⁺) are Lewis acids; lone-pair species are Lewis bases.
Hydrated metal ions are Lewis acids: small, highly charged cations (Fe³⁺, Al³⁺) make their solutions acidic by hydrolysis.
K_f quantifies complex stability; chelating ligands (like EDTA) give very large K_f values used in therapy and analysis.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

State the Lewis definitions of acid and base.
Show answer
Lewis acid: electron-pair acceptor. Lewis base: electron-pair donor.
Classify each species as a Lewis acid or base: NH₃, BF₃, H⁺, OH^-, Fe³⁺, CN^-.
Show answer
Acids: BF₃, H⁺, Fe³⁺. Bases: NH₃, OH^-, CN^- (all lone-pair donors).
Why is every Brønsted–Lowry acid–base reaction also a Lewis reaction?
Show answer
Proton transfer always involves the base donating its electron pair to H⁺ — the definition of a Lewis acid–base reaction. The Lewis view is simply more general.
Why is BF₃ a Lewis acid even though it has no hydrogen?
Show answer
Boron in BF₃ has only 6 valence electrons (a sextet) and an empty orbital; it can accept an electron pair, the defining ability of a Lewis acid.
Explain why a solution of AlCl₃ is acidic.
Show answer
Al³⁺ is a small, highly charged Lewis acid. In water it forms Al(H₂O)₆³⁺, whose bound waters have weakened O–H bonds; hydrolysis releases H⁺, acidifying the solution.
In Example 2, why does sequestering free Ag⁺ cause AgCl to dissolve?
Show answer
Complexation consumes free Ag⁺, dropping [Ag⁺] far below the saturation value. With Q = [Ag⁺][Cl^-] < K_sp, the dissolution equilibrium shifts right and AgCl dissolves.

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Study toolsKey vocabulary

Key vocabulary

Lewis acid: Any species that accepts an electron pair
Lewis base: Any species that donates an electron pair
coordinate covalent bond: A bond whose shared pair is donated entirely by one atom
adduct: The product of a Lewis acid–base reaction, e.g., F₃B-NH₃
ligand: A Lewis base that binds a metal ion in a complex
complex ion: A metal cation surrounded by ligands, e.g., [Ag(NH₃)₂]⁺
formation constant (Kf): Equilibrium constant for complex-ion formation

Sources & references

openstax.org — Chemistry Atoms First 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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