2.6.1: 2.5.1 Lewis Acid-Base Neutralization

Lewis Acid-Base Neutralization Involving Electron-Pair Transfer

Just as any Arrhenius acid is also a Brønsted acid, any Brønsted acid is also a Lewis acid, so the various acid-base concepts are all "upward compatible". Although we do not really need to think about electron-pair transfers when we deal with ordinary aqueous-solution acid-base reactions, it is important to understand that it is the opportunity for electron-pair sharing that enables proton transfer to take place.

This equation for a simple acid-base neutralization shows how the Brønsted and Lewis definitions are really just different views of the same process. Take special note of the following points:

• The arrow shows the movement of a proton from the hydronium ion to the hydroxide ion.
• Note that the electron-pairs themselves do not move; they remain attached to their central atoms. The electron pair on the base is "donated" to the acceptor (the proton) only in the sense that it ends up being shared with the acceptor, rather than being the exclusive property of the oxygen atom in the hydroxide ion.
• Although the hydronium ion is the nominal Lewis acid here, it does not itself accept an electron pair, but acts merely as the source of the proton that coordinates with the Lewis base.

The point about the electron-pair remaining on the donor species is especially important to bear in mind. For one thing, it distinguishes a Lewis acid-base reaction from an oxidation-reduction reaction, in which a physical transfer of one or more electrons from donor to acceptor does occur. The product of a Lewis acid-base reaction is known formally as an "adduct" or "complex", although we do not ordinarily use these terms for simple proton-transfer reactions such as the one in the above example. Here, the proton combines with the hydroxide ion to form the "adduct" H₂O. The following examples illustrate these points for some other proton-transfer reactions that you should already be familiar with.

Another example, showing the autoprotolysis of water. Note that the conjugate base is also the adduct.

Ammonia is both a Brønsted and a Lewis base, owing to the unshared electron pair on the nitrogen. The reverse of this reaction represents the hydrolysis of the ammonium ion.

Because HF is a weak acid, fluoride salts behave as bases in aqueous solution. As a Lewis base, F^– accepts a proton from water, which is transformed into a hydroxide ion.

The bisulfite ion is amphiprotic and can act as an electron donor or acceptor.

Note

All Brønsted–Lowry bases (proton acceptors), such as OH⁻, H₂O, and NH₃, are also electron-pair donors. Thus the Lewis definition of acids and bases does not contradict the Brønsted–Lowry definition. Rather, it expands the definition of acids to include substances other than the H⁺ ion.

Lewis Acid-Base Neutralization without Transferring Protons

Electron-deficient molecules, such as BCl₃, contain less than an octet of electrons around one atom and have a strong tendency to gain an additional pair of electrons by reacting with substances that possess a lone pair of electrons. Lewis’s definition, which is less restrictive than either the Brønsted–Lowry or the Arrhenius definition, grew out of his observation of this tendency. A general Brønsted–Lowry acid–base reaction can be depicted in Lewis electron symbols as follows:

The proton (H⁺), which has no valence electrons, is a Lewis acid because it accepts a lone pair of electrons on the base to form a bond. The proton, however, is just one of many electron-deficient species that are known to react with bases. For example, neutral compounds of boron, aluminum, and the other Group 13 elements, which possess only six valence electrons, have a very strong tendency to gain an additional electron pair. Such compounds are therefore potent Lewis acids that react with an electron-pair donor such as ammonia to form an acid–base adduct, a new covalent bond, as shown here for boron trifluoride (BF₃):

The bond formed between a Lewis acid and a Lewis base is a coordinate covalent bond because both electrons are provided by only one of the atoms (N, in the case of F₃B:NH₃). After it is formed, however, a coordinate covalent bond behaves like any other covalent single bond.

Species that are very weak Brønsted–Lowry bases can be relatively strong Lewis bases. For example, many of the group 13 trihalides are highly soluble in ethers (R–O–R′) because the oxygen atom in the ether contains two lone pairs of electrons, just as in H₂O. Hence the predominant species in solutions of electron-deficient trihalides in ether solvents is a Lewis acid–base adduct. A reaction of this type is shown in Figure \(\PageIndex{1}\) for boron trichloride and diethyl ether:

\(\textbf{Figure } \PageIndex{1}\): Lewis Acid/Base reaction of boron trichloride and diethyl ether reaction

Note

• Electron-deficient molecules (those with less than an octet of electrons) are Lewis acids.
• The acid-base behavior of many compounds can be explained by their Lewis electron structures.

Many molecules with multiple bonds can act as Lewis acids. In these cases, the Lewis base typically donates a pair of electrons to form a bond to the central atom of the molecule, while a pair of electrons displaced from the multiple bond becomes a lone pair on a terminal atom.

Figure \(\PageIndex{2}\). The highly electronegative oxygen atoms pull electron density away from carbon, so the carbon atom acts as a Lewis acid. Arrows indicate the direction of electron flow.

\(\textbf{Figure } \PageIndex{2}\): Lewis Acid/Base reaction of the hydroxide ion with carbon dioxide

Example \(\PageIndex{1}\)

Identify the acid and the base in the Lewis acid–base reaction.

BH₃ + (CH₃)₂S → H₃B:S(CH₃)₂

Solution

In each equation, identify the reactant that is electron deficient and the reactant that is an electron-pair donor. The electron-deficient compound is the Lewis acid, whereas the other is the Lewis base.

In BH₃, boron has only six valence electrons. It is therefore electron deficient and can accept a lone pair. Like oxygen, the sulfur atom in (CH₃)₂S has two lone pairs. Thus (CH₃)₂S donates an electron pair on sulfur to the boron atom of BH₃. The Lewis base is (CH₃)₂S, and the Lewis acid is BH₃

Example \(\PageIndex{2}\)

Identify the acid and the base in the Lewis acid–base reaction.

CaO + CO₂ → CaCO₃

Solution

In each equation, identify the reactant that is electron deficient and the reactant that is an electron-pair donor. The electron-deficient compound is the Lewis acid, whereas the other is the Lewis base.

CO₂ accepts a pair of electrons from the O²⁻ ion in CaO to form the carbonate ion. The oxygen in CaO is an electron-pair donor, so CaO is the Lewis base. Carbon accepts a pair of electrons, so CO₂ is the Lewis acid.

Problems

Exercise \(\PageIndex{1}\)

Identify the acid and the base in the Lewis acid–base reaction.

(CH₃)₂O + BF₃ → (CH₃)₂O:BF₃

Answer

Lewis base: (CH₃)₂O; Lewis acid: BF₃

Exercise \(\PageIndex{2}\)

Identify the acid and the base in the Lewis acid–base reaction.

H₂O + SO₃ → H₂SO₄

Answer

Lewis base: H₂O; Lewis acid: SO₃

Exercise \(\PageIndex{3}\)

Identify the acid and the base in the Lewis acid–base reaction.

BeCl₂ + 2 Cl⁻ → BeCl₄²⁻

Answer

The chloride ion contains four lone pairs. In this reaction, each chloride ion donates one lone pair to BeCl₂, which has only four electrons around Be. Thus the chloride ions are Lewis bases, and BeCl₂ is the Lewis acid

Exercise \(\PageIndex{4}\)

Here are several more examples of Lewis acid-base reactions that cannot be accommodated within the Brønsted or Arrhenius models. Identify the Lewis acid and Lewis base in each reaction.

1. \(Al(OH)_3 + OH^– \rightarrow Al(OH)_4^–\)
2. \(SnS_2 + S^{2–} \rightarrow SnS_3^{2–}\)
3. \(Cd(CN)_2 + 2 CN^– \rightarrow Cd(CN)_4^{2+}\)
4. \(AgCl + 2 NH_3 \rightarrow Ag(NH_3)_2^+ + Cl^–\)
5. \(Fe^{2+} + NO \rightarrow Fe(NO)^{2+}\)
6. \([Ni^{2+} + 6 NH_3 \rightarrow Ni(NH_3)_5^{2+}\)

Answer a

Lewis base: OH^-; Lewis acid: Al(OH)₃

 

Answer b

Lewis base: S^2-; Lewis acid: SnS₂

 

Answer c

Lewis base: CN^-; Lewis acid: Cd(CN)₂

 

Answer d

Lewis base: NH₃; Lewis acid: AgCl

 

Answer e

Lewis base: NO; Lewis acid: Fe²⁺

 

Answer f

Lewis base: NH₃; Lewis acid: Ni²⁺

Contributors and Attributions

• Stephen Lower, Professor Emeritus (Simon Fraser U.) Chem1 Virtual Textbook

• Wikipedia