Organic Chemistry · Polar Covalent Bonds; Acids and Bases

Acids and Bases: The Lewis Definition

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Lewis acid–base definitions, the boron trihalide acidity trend, and the Friedel–Crafts mechanism are standard textbook material; verify quantitative claims (e.g., specific acidity rankings in solution) against current primary sources before formal citation.

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

The Brønsted–Lowry definition of acids and bases is built around proton transfer, but many chemical reactions — especially in organic chemistry — involve electron-pair donation without any proton moving. G. N. Lewis proposed a broader definition in 1923: a is an electron-pair acceptor, and a is an electron-pair donor. The product of a Lewis acid–base reaction is called an (or complex), and the new bond formed is a , in which both electrons come from the Lewis base. This definition includes all Brønsted acids and bases (a proton H⁺ is an electron-pair acceptor; hydroxide and amines are electron-pair donors) but also covers species with no protons at all: boron trifluoride (BF₃, electron-deficient), aluminum chloride (AlCl₃), carbocations, and essentially every metal cation. In organic chemistry the Lewis framework is the language of nucleophiles and electrophiles — the electron-rich and electron-poor partners that drive most reactions — making this topic the conceptual bridge between acid–base theory and the reaction mechanisms that follow.

Why this matters

Reaction mechanism vocabulary: every – reaction is a Lewis acid–base reaction. Recognizing which atom is electron-rich and which is electron-poor tells you where bonds will form before you learn any specific mechanism.
Catalysis: BF₃, AlCl₃, FeBr₃, and TiCl₄ are Lewis acid catalysts in industrial and laboratory chemistry — for example, AlCl₃ catalyzes Friedel–Crafts acylation by accepting a chloride ion to generate the reactive .
Biochemistry: metal ions (Mg²⁺, Zn²⁺, Fe³⁺) act as Lewis acids in enzymes; ATP binding to Mg²⁺ and zinc in carbonic anhydrase are classic examples.
Bonding insight: the coordinate covalent bond explains why NH₄⁺, H₃O⁺, and BF₃ · NH₃ exist and why they are stable.

The college version

Core Concepts

The Lewis definitions

$Lewis acid: electron-pair acceptor (electrophile) Lewis base: electron-pair donor (nucleophile)$

A reaction between a Lewis acid A and a Lewis base B: forms an adduct with a coordinate covalent bond:

$A + B \to A : B (adduct)$

The double-headed arrow (curved arrow) in mechanism drawings shows the electron pair moving from the base's lone pair (or π bond) to the acid's empty orbital. For example, BF₃ + NH₃ → F₃B-NH₃: the nitrogen lone pair of ammonia fills the empty p orbital of boron, and the adduct F₃B-NH₃ forms. In the adduct, boron is tetrahedral and both B–N electrons came from nitrogen — a coordinate covalent bond.

Relationship to the Brønsted–Lowry definition

Every Brønsted acid–base reaction is also a Lewis reaction, but not vice versa. The proton H⁺ is the ultimate Lewis acid — an empty 1s orbital — so H⁺ + OH^- → H₂O is a Lewis reaction (OH⁻ donates a lone pair to H⁺). But BF₃ + NH₃ involves no proton transfer at all, so only the Lewis definition covers it. The Lewis definition is therefore the most general; the Brønsted definition is a special case limited to proton transfer.

Common Lewis acids and bases in organic chemistry

Lewis acids (electron-pair acceptors):

Molecules with an incomplete octet: BF₃, AlCl₃, FeBr₃, BBr₃, carbocations R₃C⁺, and H⁺.
Metal cations: Mg²⁺, Zn²⁺, Fe³⁺, Cu²⁺.
Polarized π systems: the carbonyl carbon of aldehydes/ketones/esters acts as a Lewis acid toward nucleophiles.

Lewis bases (electron-pair donors):

Species with lone pairs: amines (R₃N), ethers (R₂O), alcohols, water, halide ions (Cl^-, Br^-, I^-), hydroxide, thiols (RSH).
π-bond electrons: alkenes and aromatic rings donate their π electrons to Lewis acids such as Br₂ or H⁺ — the first step of addition and substitution reactions.

Why some species are stronger Lewis acids: the boron trihalide puzzle

Among the boron trihalides, the Lewis acidity toward a fixed base follows the order BF₃ < BCl₃ < BBr₃. Fluorine is the most electronegative halogen, so you might expect BF₃ to be the strongest acid — but it is the weakest. The reason is : fluorine's lone pairs overlap with boron's empty p orbital, partially filling the octet that the incoming base must fill. The better the halogen's lone-pair overlap with boron (F > Cl > Br), the less the boron "needs" the base, and the weaker the Lewis acidity. This effect (called π donation or back-bonding) is an exam favorite.

Lewis acids and bases in catalysis

In Friedel–Crafts acylation, aluminum chloride catalyzes the reaction by accepting a chloride ion from acetyl chloride:

$AlCl 3 + CH 3 COCl \to AlCl 4 - + CH 3 CO + (acylium ion)$

The acylium ion, a powerful electrophile (Lewis acid), then accepts electron density from the aromatic ring (a Lewis base). The mechanism is: aluminum chloride pulls a chloride off the acyl chloride (AlCl₃ is the Lewis acid, chloride the Lewis base), generating the reactive acylium cation, which is then attacked by the π electrons of the benzene ring. This single concept — electron-pair acceptance — explains the entire catalytic cycle.

Common Confusions

Do Not Confuse	With	Difference
Lewis acid	Brønsted acid	A Brønsted acid must donate a proton; a Lewis acid only needs to accept an electron pair (BF₃, AlCl₃, carbocations have no proton to donate).
Lewis base	Brønsted base	Brønsted bases accept protons; Lewis bases donate electron pairs — π bonds and halide ions count even when no proton is involved.
Coordinate covalent bond	Ordinary covalent bond	Both have a shared pair; only the origin differs (coordinate: both electrons from one partner). The resulting bond is indistinguishable once formed.
Electrophile/nucleophile	Lewis acid/base	Same concept, different vocabulary: mechanisms use electrophile/nucleophile; acid–base theory uses Lewis acid/base.
BF₃ strength by electronegativity	BF₃ strength by π back-donation	Electronegativity alone predicts BF₃ strongest; π back-donation makes it weakest. The π effect wins.
Adduct	Reaction intermediate	An adduct is an isolable product of Lewis acid + base; intermediates are transient species along a reaction path.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of a Lewis base as someone with spare LEGO bricks (electrons) and a Lewis acid as someone with an empty LEGO board (an empty spot for electrons). The base shares its bricks with the acid, and they click together into one piece. Some acids don't even have protons to give — like boron trifluoride, which just has an empty spot and wants bricks. The Brønsted idea was only about giving away a hydrogen atom; the Lewis idea is about sharing building blocks, which covers many more situations.

Worked example

Example 1: Classifying a reaction — BF₃ + NH₃

Problem: Identify the Lewis acid and Lewis base in BF₃ + NH₃ → F₃B-NH₃, and describe the bonding change.

Plan and formula: boron in BF₃ has 3 valence electrons and forms 3 B–F bonds with 6 shared electrons — an incomplete octet (6 electrons, not 8). It needs two more electrons, so it is the electron-pair acceptor: the Lewis acid. Nitrogen in NH₃ has a lone pair not involved in bonding; it can donate it: the Lewis base.

Answer: BF₃ is the Lewis acid; NH₃ is the Lewis base. The nitrogen lone pair forms a coordinate covalent bond to boron, and boron becomes tetrahedral (4 bonds, complete octet). The adduct F₃B-NH₃ is stable and isolable. This reaction pattern — empty orbital fills with lone pair — is the template for countless organic reactions, including amine–borane reducing agents.

Example 2: Ranking Lewis acidity of the boron trihalides

Problem: Rank BF₃, BCl₃, and BBr₃ by Lewis acidity toward ammonia, and explain the trend.

Plan and formula: Lewis acidity depends on how much the boron "wants" the incoming base, which depends on how well its empty p orbital is already filled by halogen lone pairs (π back-donation). Overlap of halogen lone pairs with boron's empty p orbital: F (2p–2p, best overlap) > Cl (3p–2p) > Br (4p–2p, worst overlap).

$π back-donation: BF 3 > BCl 3 > BBr 3 \Rightarrow Lewis acidity: BF 3 < BCl 3 < BBr 3$

Answer: BBr₃ is the strongest Lewis acid toward ammonia, BF₃ the weakest — the opposite of what electronegativity alone predicts. Fluorine's lone pairs fill boron's empty orbital most effectively, so BF₃ has the least need for ammonia's electrons. This counterintuitive trend is a standard exam question.

Example 3: The carbonyl carbon as a Lewis acid

Problem: The carbonyl group of an aldehyde, R₂C=O, reacts with a hydride ion H^- (as in NaBH₄ reduction). Identify the Lewis acid and base, and describe the product.

Plan and formula: in the carbonyl, the oxygen is more electronegative than carbon, so the π bond is polarized: carbon carries a partial positive charge and has an empty p orbital that can accept electron density. The hydride ion carries a lone pair. Therefore:

$Lewis acid: carbonyl carbon (δ+) Lewis base: hydride, H -$

Answer: the hydride lone pair forms a new C–H bond at the carbonyl carbon; the π electrons shift onto oxygen, giving an alkoxide R₂CH-O^-, which is protonated to the alcohol R₂CH-OH on workup. This one Lewis acid–base step — nucleophilic addition to a polarized π bond — is the foundation of carbonyl chemistry in later chapters.

Key takeaways

Lewis acid = electron-pair acceptor (electrophile); Lewis base = electron-pair donor (nucleophile). No proton needed.
Product of a Lewis acid–base reaction = adduct, held by a coordinate covalent bond (both electrons from the base).
Every Brønsted reaction is a Lewis reaction (H⁺ is a Lewis acid); the reverse is false — BF₃ + NH₃ involves no proton transfer.
Incomplete-octet molecules (BF₃, AlCl₃, carbocations) and metal cations are Lewis acids; lone-pair species (amines, ethers, halides) and π bonds are Lewis bases.
Boron trihalide acidity: BF₃ < BCl₃ < BBr₃ because F's lone pairs best fill boron's empty p orbital (π back-donation).
AlCl₃ catalyzes Friedel–Crafts acylation by abstracting Cl⁻ to form the acylium ion CH₃CO⁺, the actual electrophile.
Nucleophile/electrophile terminology in later chapters is Lewis base/Lewis acid in disguise.

Check yourself

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

Define Lewis acid and Lewis base in your own words, without mentioning protons.
Show answer
A Lewis acid is an electron-pair acceptor (it has an empty orbital); a Lewis base is an electron-pair donor (it has a lone pair or π electrons it can share). No proton transfer required.
Why is every Brønsted acid–base reaction also a Lewis acid–base reaction, but not the reverse?
Show answer
Because H⁺ is itself an electron-pair acceptor (empty 1s orbital), every proton transfer is a special case of electron-pair donation. But Lewis reactions like BF₃ + NH₃ involve no proton at all, so they fall outside the Brønsted definition.
In AlCl₃ + Cl^- → AlCl₄^-, which species is the Lewis acid and which is the Lewis base? What is the product called?
Show answer
AlCl₃ is the Lewis acid (incomplete octet on Al), Cl^- is the Lewis base; the product AlCl₄^- is an adduct (complex) with a coordinate covalent Al–Cl bond.
Why is BF₃ a weaker Lewis acid than BBr₃ toward ammonia, even though fluorine is more electronegative than bromine?
Show answer
Fluorine's lone pairs overlap strongly with boron's empty p orbital (π back-donation), partially filling the octet the base would otherwise fill. Bromine's 4p orbitals overlap poorly with boron's 2p orbital, so BBr₃ remains more electron-hungry and is the stronger Lewis acid.
Identify the Lewis acid and Lewis base in the reaction of an alkene with H⁺ to form a carbocation.
Show answer
The alkene's π electrons act as the Lewis base (electron-pair donor); H⁺ is the Lewis acid (electron-pair acceptor). The π bond donates to H⁺, forming a C–H bond and leaving a carbocation.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Lewis acid: Electron-pair acceptor (electrophile)
Lewis base: Electron-pair donor (nucleophile)
Adduct: The product of a Lewis acid–base reaction
Coordinate covalent bond: A bond whose two electrons come from one partner
Electrophile: Electron-loving species; another name for Lewis acid
Nucleophile: Nucleus-loving species; another name for Lewis base
π back-donation: Halogen lone pairs filling boron's empty p orbital
Acylium ion: RCO⁺, a resonance-stabilized carbocation

Sources & references

openstax.org — Organic Chemistry

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

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In 30 seconds

Why this matters

The college version

Core Concepts

The Lewis definitions

Relationship to the Brønsted–Lowry definition

Common Lewis acids and bases in organic chemistry

Why some species are stronger Lewis acids: the boron trihalide puzzle

Lewis acids and bases in catalysis

Common Confusions

The same idea, in plain words

Worked example

Example 1: Classifying a reaction — BF₃ + NH₃

Example 2: Ranking Lewis acidity of the boron trihalides

Example 3: The carbonyl carbon as a Lewis acid

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references