Chem1 General Chemistry Virtual Textbook →Lewis acids and bases
Lewis acids and bases
Electron-pair donors and acceptors
The Brønsted-Lowry proton donor-acceptor concept has been one of the most successful theories of Chemistry. But as with any such theory, it is fair to ask if this is not just a special case of a more general theory that could encompass an even broader range of chemical science. In 1916, G.N. Lewis of the University of California proposed that the electron pair is the dominant actor in acid-base chemistry. The Lewis theory did not become very well known until about 1923 (the same year that Brønsted and Lowry published their work), but since then it has been recognized as a very powerful tool for describing chemical reactions of widely different kinds and is widely used in organic and inorganic chemistry.
In the last section of this unit, we will change the subject slightly and take a brief look at protonic acid-base systems that involve solvents other than water.
According to Lewis,
In modern chemistry, electron donors are often referred to as nucleophiles, while acceptors are electrophiles.
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 don't 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 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 don't 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, being amphiprotic, can act as an electron donor or acceptor.|
Linda Sweeting of Towson State U. has written a tutorial article
on aqueous bases as nucleophiles in organic chemistry.
The major utility of the Lewis definition is that it extends the concept of acids and bases beyond the realm of proton transfer reactions.
The classic example, given in every textbook, is the reaction of boron trifluoride with ammonia to form an adduct:
BF3 + NH3 → F3B-NH3
If you are already familiar with the concept of hybrid orbitals, this diagram (from a Purdue University site which is well worth looking at) might help you visualize this process. The lone pair electrons on the nitrogen atom can be acommodated in the empty 2pz orbital of BF3, creating a shared-electron B–N bond.
One of the most commonly-encountered kinds of Lewis acid-base reactions occurs when electron-donating ligands form coordination complexes with transition-metal ions.
In this example, the tin atom in SnCl4 can expand its valence shell by utilizing a pair of d-orbitals, changing its hybridization from sp3 to sp3d2.
Here are several more examples of Lewis acid-base reactions that cannot be accommodated within the Brønsted or Arrhenius models; be sure you are able to identify the Lewis acid and Lewis base in each reaction.
Al(OH)3 + OH– → Al(OH)4–
SnS2 + S2–→ SnS32–
Cd(CN)2 + 2 CN–→ Cd(CN)42+
AgCl + 2 NH3→ Ag(NH3)2+ + Cl–
Fe2+ + NO → Fe(NO)2+
Ni2+ + 6 NH3→ Ni(NH3)52+
A much more extensive discussion of Lewis acids and bases and extension of the concept to HOMO/LUMO theory and "hard" and "soft" acids (for more advanced students), can be found at Mark Leach's ChemoGenesis site.
Although organic chemistry is beyond the scope of these lessons, it is instructive to see how electron donors and acceptors play a role in chemical reactions. The following two diagrams (from this site by William Reusch of Michigan State U.) show the mechanisms of two common types of reactions initiated by simple inorganic Lewis acids:
In each case, the species labeled "Complex" is an intermediate that decomposes into the products, which are conjugates of the original acid and base pairs. The electric charges indicated in the complexes are formal charges, but those in the products are "real".
In (1), the incomplete octet of the aluminum atom in AlCl3 serves as a better electron acceptor to the chlorine atom than does the isobutyl part of the base. In (2), the pair of non-bonding electrons on the dimethyl ether coordinates with the electron-deficient boron atom, leading to a complex that breaks down by releasing a bromide ion.
We ordinarily think of Brønsted-Lowry acid-base reactions as taking place in aqueous solutions, but this need not always be the case. A more general view encompasses a variety of acid-base solvent systems, of which the water system is only one. Each of these has as its basis an amphiprotic solvent (one capable of undergoing autoprotolysis), in parallel with the familiar case of water.
This is the one you are most likely to encounter if you do advanced work in Chemistry. Liquid ammonia boils at –33° C, and can conveniently be maintained as a liquid by cooling with dry ice (–77° C.) It is a good solvent for substances that also dissolve in water, such as ionic salts and organic compounds capable of forming hydrogen bonds.
Ammonia-based life - an interesting article
Many other familiar substances can serve as the basis of protonic solvent systems, a few of which are listed here:
|water||2 H2O → H3O+ + OH–||14|
|ammonia||2 NH3 → NH4+ + NH2–||33|
|acetic acid||2 CH3COOH → CH3COOH2+ + CH3COO–||13|
|ethanol||2 C2H5OH → C2H5OH2+ + C2H5O–||19|
|hydrogen peroxide||2 HO-OH → HO-OH2+ + HO-O–||13|
|hydrofluoric acid||2 HF → H2F+ + F–||10|
|sulfuric acid||2 H2SO4 → H3SO4+ + HSO4–||3.5|
One use of nonaqueous acid-base systems is to examine the relative strengths of the strong acids and bases, whose strengths are "leveled" by the fact that they are all totally converted into H3O+ or OH– ions in water. By studying them in appropriate non-aqueous solvents which are poorer acceptors or donors of protons, their relative strengths can be determined.
Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic.
Page last modified: 17.02.2007