Table of Contents
Lewis acid–base theory extends the idea of acids and bases far beyond substances that donate or accept protons. Instead, it focuses purely on electron pairs and is therefore particularly useful for understanding many inorganic and organic reactions, as well as coordination chemistry and catalysis.
Definition of Lewis Acids and Lewis Bases
In the Lewis concept, acids and bases are defined in terms of electron pairs:
- Lewis acid: an electron pair acceptor
- Lewis base: an electron pair donor
A Lewis acid–base reaction is the formation of a new bond by donation of an electron pair from the base to the acid. The product is often called an adduct or a Lewis acid–base complex.
Symbolically:
$$
\text{Lewis base} : \; + \; \text{Lewis acid} \; \rightarrow \; \text{Lewis adduct}
$$
The colon : is often used to represent a lone pair of electrons on the base.
Relationship to Brønsted Acid–Base Theory
In Brønsted theory, acids donate protons and bases accept protons. Every Brønsted acid–base reaction can also be seen as a Lewis acid–base reaction because:
- The proton $H^+$ has no electrons and acts as an extreme Lewis acid (it can accept an electron pair).
- The Brønsted base donates an electron pair to bind the proton.
Example:
$$
\ce{NH3 + H+ -> NH4+}
$$
Viewed as a Lewis reaction:
- $\ce{NH3}$ is a Lewis base (donates a lone pair on N).
- $\ce{H+}$ is a Lewis acid (accepts the lone pair).
- $\ce{NH4+}$ is the Lewis adduct.
However, Lewis theory also covers many reactions in which no proton is involved and which therefore lie outside Brønsted theory.
Typical Examples of Lewis Acids and Bases
Common Lewis Bases
Lewis bases are species that can donate an electron pair. Typical features:
- Lone pairs on electronegative atoms (N, O, S, halogens)
- $\pi$-electron systems
- Anions with excess electron density
Examples:
- Neutral molecules with lone pairs:
- $\ce{H2O}$, $\ce{NH3}$, $\ce{RNH2}$ (amines)
- $\ce{PH3}$, $\ce{CO}$, $\ce{RS-}$ donors in organic chemistry
- Anions:
- $\ce{Cl-}$, $\ce{Br-}$, $\ce{F-}$, $\ce{OH-}$, $\ce{CN-}$, $\ce{S^{2-}}$, $\ce{NO2-}$, etc.
- Organic $\pi$-systems:
- Alkenes ($\ce{C=C}$), alkynes, aromatic rings such as benzene can act as weak Lewis bases by donating $\pi$-electron density to strong Lewis acids.
In Lewis structures, lone pairs are often drawn explicitly to highlight their donor ability:
- $\ce{:NH3}$ (one lone pair on N)
- $\ce{:O(H)2}$ (two lone pairs on O)
Common Lewis Acids
Lewis acids are species that can accept an electron pair. Typical features:
- Electron-deficient atoms (often with incomplete octet)
- Positively charged ions (cations)
- Atoms with low-lying empty orbitals
Typical classes:
- Electron-deficient main-group compounds
- $\ce{BF3}$, $\ce{BCl3}$ (boron trihalides, 6 valence electrons)
- $\ce{AlCl3}$ and other aluminum halides
- Some group 13 (IIIA) compounds in general
- Metal cations
- $\ce{Na+}$, $\ce{Mg^{2+}}$, $\ce{Al^{3+}}$, $\ce{Fe^{3+}}$, $\ce{Cu^{2+}}$, etc.
- Transition metal cations especially act as Lewis acids toward ligands in coordination complexes.
- Molecules with polar multiple bonds
- Carbonyl compounds $\ce{R2C=O}$: the carbonyl carbon is electron-poor
- $\ce{SO3}$, $\ce{CO2}$ (electrophilic centers at the central atom)
- Protonated species like $\ce{H3O+}$: the central atom can still accept electron density.
- Proton $\ce{H+}$
- The simplest and strongest Lewis acid conceptually: completely electron-deficient, accepts an electron pair from a base.
Thus, the class of Lewis acids is much broader than the class of Brønsted acids.
Lewis Acid–Base Adduct Formation
A typical Lewis reaction is adduct formation by coordinate (dative) bond formation. In a coordinate bond, both electrons in the new bond originate from the Lewis base.
A classic example:
$$
\ce{BF3 + :NH3 -> F3B{\ bond}NH3}
$$
- $\ce{BF3}$ is a Lewis acid (electron-deficient B).
- $\ce{NH3}$ is a Lewis base (lone pair on N).
- The N lone pair is donated into the empty orbital on B, forming a coordinate bond.
The product is often written as:
- $\ce{F3B \leftarrow NH3}$ to emphasize the direction of donation, or
- $\ce{F3B-NH3}$ once formed, since afterward the bond is indistinguishable from other covalent bonds.
Another example:
$$
\ce{AlCl3 + Cl- -> [AlCl4]-}
$$
- $\ce{AlCl3}$: Lewis acid
- $\ce{Cl-}$: Lewis base
- The chloride donates a lone pair to aluminum, forming the tetrahedral complex ion $\ce{[AlCl4]-}$.
Coordinate (Dative) Bonds
The bond formed in Lewis acid–base reactions is often called a coordinate bond or dative bond. At the moment of bond formation:
- Both electrons come from the donor (Lewis base).
- The acceptor (Lewis acid) contributes an empty orbital.
Notation variants:
- $\ce{BF3 + :NH3 -> BF3 \leftarrow NH3}$ (arrow shows electron donation direction)
- Sometimes drawn with an arrow on the bond itself: $\ce{BF3 <-: NH3}$ (not standard chemical formula, but emphasizes direction)
Once formed, a coordinate bond is generally treated like any other covalent bond in structure and reactivity.
Coordinate bonding is central to:
- Complex formation in coordination chemistry (metal–ligand bonds)
- Many catalytic cycles
- Binding of small molecules (e.g., $\ce{CO}$, $\ce{NH3}$) to metal centers
Beyond Proton Transfer: Reactions Explained Only by Lewis Theory
Because the Lewis definition centers on electron pairs rather than protons, it can describe many reactions that the Brønsted concept does not classify as acid–base reactions.
Example: Reaction of $\ce{BF3}$ and $\ce{F-}$
$$
\ce{BF3 + F- -> [BF4]-}
$$
- No proton transfer occurs.
- Yet this is clearly an interaction of an electron-deficient center with an electron-rich species.
- $\ce{BF3}$ is a Lewis acid, $\ce{F-}$ is a Lewis base, and $\ce{[BF4]-}$ is the adduct.
Example: Metal–Ligand Complex Formation
For a metal ion $\ce{Cu^{2+}}$ in water:
$$
\ce{Cu^{2+} + 6 H2O -> [Cu(H2O)6]^{2+}}
$$
- $\ce{Cu^{2+}}$ is a Lewis acid.
- $\ce{H2O}$ is a Lewis base donating lone pairs on O.
- The complex $\ce{[Cu(H2O)6]^{2+}}$ is a Lewis acid–base adduct with six coordinate bonds.
Such reactions are key to understanding:
- Coordination compounds
- Hydrated ions in solution
- Many biological metal centers
Example: Carbonyl Additions
In organic chemistry, nucleophiles (Lewis bases) attack electrophilic carbonyl carbons (Lewis acids):
$$
\ce{R2C=O + :Nu- -> R2C(ONu)-}
$$
- $\ce{R2C=O}$: the carbonyl carbon is electron-poor, acts as Lewis acid center.
- $\ce{:Nu-}$: nucleophile with lone pair (e.g., $\ce{CN-}$, $\ce{OH-}$, $\ce{R-O-}$).
The first step is a Lewis acid–base interaction: electron pair donation from $\ce{Nu-}$ into the empty antibonding orbital at the carbonyl carbon.
Classification of Lewis Acids and Bases: Hard and Soft
Lewis acids and bases can be further categorized to predict which combinations are particularly stable. A widely used concept is the HSAB principle (Hard and Soft Acids and Bases). The detailed theory belongs elsewhere; here only the basic idea is outlined.
- Hard species: small, high charge, low polarizability.
- Soft species: larger, more polarizable, lower charge density.
Empirically:
- Hard acids prefer to bind hard bases.
- Soft acids prefer to bind soft bases.
Examples (without going into mechanistic detail):
- Hard acids: $\ce{H+}$, $\ce{Li+}$, $\ce{Al^{3+}}$, $\ce{Fe^{3+}}$
- Hard bases: $\ce{F-}$, $\ce{OH-}$, $\ce{H2O}$, $\ce{NH3}$
- Soft acids: $\ce{Ag+}$, $\ce{Hg^{2+}}$, $\ce{Pd^{2+}}$
- Soft bases: $\ce{I-}$, $\ce{S^{2-}}$, phosphines $\ce{PR3}$, alkenes, aromatic rings
This classification helps:
- Rationalize stability of complexes
- Predict reaction pathways and selectivity in inorganic and organic chemistry
Lewis Concept and Acid–Base Strength
Just as with Brønsted acids and bases, Lewis acids and bases can be stronger or weaker relative to each other. However, in Lewis theory, “strength” depends strongly on:
- The specific reaction partner (acid–base pair)
- Solvent and environment
- Geometry and steric hindrance
- Electronic structure (availability of empty or filled orbitals, charge distribution)
Some general trends:
- The more electron-deficient and electrophilic a species is, the stronger a Lewis acid it tends to be.
- The more electron-rich (and less stabilized by resonance or electronegativity) a species is, the stronger a Lewis base it tends to be.
- Resonance and inductive effects can increase or decrease Lewis acidity/basicity by redistributing electron density.
Examples of relative behavior (in a given context):
- $\ce{BF3}$ is a stronger Lewis acid than $\ce{BCl3}$ in some conditions due to differences in $\pi$-backbonding.
- Among halides, $\ce{F-}$ is often less polarizable and sometimes a weaker donor towards soft acids than $\ce{I-}$.
Important: unlike simple pH scales for Brønsted acidity in water, Lewis acidity and basicity cannot be captured by a single universal numeric scale; they are context-dependent.
Applications and Importance of the Lewis Concept
The Lewis acid–base concept is widely used to interpret and design reactions in many areas of chemistry:
- Inorganic chemistry and coordination chemistry
- Description of metal–ligand bonding
- Understanding structures of complex ions and catalysts
- Organic reaction mechanisms
- Nucleophile–electrophile interactions (nucleophiles as Lewis bases, electrophiles as Lewis acids)
- Activation of substrates by Lewis acids (e.g., $\ce{AlCl3}$ in Friedel–Crafts reactions)
- Catalysis
- Many catalysts are Lewis acids (e.g., $\ce{AlCl3}$, $\ce{TiCl4}$, $\ce{BF3}$, metal centers in enzymes)
- Some catalysts are Lewis bases (e.g., tertiary amines) activating substrates or stabilizing intermediates
- Material and solid-state chemistry
- Acidic sites on solid surfaces (e.g., $\ce{Al2O3}$, zeolites) behaving as Lewis acids
- Adsorption of molecules via Lewis acid–base interactions
By shifting attention from protons to electron pairs, Lewis theory provides a unifying language for donor–acceptor reactions throughout chemistry.