Lewis acid and base theory views the electron as the active species in an acid-base reaction. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. This contrasts with Arrhenius and Bronsted-Lowry acids and bases, which view the reaction from the behavior of the hydrogen ion or proton, respectively. The advantages of Lewis theory is that it expands the list of acids and bases and it works well with oxidation-reduction reactions.
- A Lewis acid accepts an electron pair to form a covalent bond.
- A Lewis base donates an electron pair to form a covalent bond.
History
American physical chemist Gilbert N. Lewis applied his understanding of chemical bonding to his acid-base theory. In 1916, Lewis proposed that a covalent bond forms when each atom contributes one electron to form an electron pair that the atoms share. When both electrons come from one atom, the chemical bond is a coordinate or dative covalent bond. In 1923, Lewis described an acid as as a substance which “can employ an electron lone pair from another molecule in completing the stable group of one of its own atoms.” In 1963, the theory was expanded to classify hard and soft acids and bases (HSAB theory).
How Lewis Acids and Bases Work
A Lewis acid-base reaction involves the transfer of a pair of electrons from a base to an acid. For example, the nitrogen atom in ammonia (NH3) has an electron pair. When ammonia reacts with the hydrogen ion (H+), the electron pair transfers to the hydrogen, forming the ammonium ion (NH4+).
NH3 + H+ → NH4+
So, ammonia is a Lewis base and the hydrogen cation is a Lewis acid. Both Arrhenius and Bronsted-Lowry theory describe this acid-base reaction.
However, Lewis acid and base theory also allows for acids that do not contain hydrogen. For example, boron trifluoride (BF3) is a Lewis acid when it reacts with ammonia (which is once again a Lewis base):
NH3 + BF3 → NH3BF3
The nitrogen donates the electron pair to the boron atom. The two molecules directly combine and form an adduct. The bond that forms between the two species is a coordinate bond or dative covalent bond.
Examples of Lewis Acids and Bases
Lewis bases include the usual bases under other definitions. Examples of Lewis bases include OH–, NH3, CN–, and H2O. Lewis acids include the usual acids, plus species not viewed as acids under other definitions. Examples of Lewis acids include H+, HCl, Cu2+, CO2, SiBr4, AlF3, BF3, H2O.
| Lewis Acids | Lewis Bases |
|---|---|
| lone-pair acceptors | lone-pair donors |
| electrophiles | nucleophiles |
| metal cations (e.g., Ag+, Mg2+) | Bronsted-Lowry bases |
| the proton (H+) | ligands |
| electron-poor π-systems | electron-rich π-systems |
Hard and Soft Lewis Acids and Bases (HSAB Theory)
Lewis acids and bases are classified according to hardness or softness. Hard implies small and not polarizable. Soft applies to larger, polarizable atoms.
- Examples of hard acids are H+, alkali metal cations, alkaline earth metal cations, Zn2+, boranes.
- Examples of soft acids are Ag+, Pt2+, Ni(0), Mo(0).
- Typical hard bases are ammonia, amines, water, fluoride, chloride, and carboxylates.
- Examples of soft bases are carbon monoxide, iodide, thioethers, and organophosphines.
HSAB theory helps when predicting the strength of adduct formation or the products of metathesis reactions. Hard-hard interactions are enthalpy-favored. Soft-soft interactions are entropy-favored.
Amphoteric Species
Some chemical species are amphoteric, meaning they can act as either a Lewis acid or as a Lewis base, depending on the situation. Water (H2O) is a great example.
Water acts as an acid when it reacts with ammonia:
H2O + NH3 → NH4+ + OH−
It acts as a base when it reacts with hydrochloric acid:
H2O + HCl → Cl– + H3O+
Aluminum hydroxide [Al(OH)3] is an example of an amphoteric compound under the Lewis theory. It acts as a Lewis base in the reaction with the hydrogen ion:
Al(OH)3 + 3H+ → Al3+ + 3H2O
It acts as a Lewis acid in the reaction with the hydroxide ion:
Al(OH)3 + OH− → Al(OH)4–
Lewis Acids and Bases vs Bronsted-Lowry Acids and Bases
The Bronsted-Lowry theory of acids and bases was published the same year as the Lewis theory. The two theories predict acids and bases using different criteria, but mostly the list of acids and bases is the same.
All Bronsted-Lowry bases are Lewis bases. All Bronsted-Lowry acids are Lewis acids. Also, the conjugate base of a Bronsted-Lowry acid is a Lewis base. However, there are some Lewis acids which are not Bronsted-Lowry acids. Also, some Lewis bases do not readily protonate, yet they react with Lewis acids. For example, carbon monoxide (CO) is a Lewis base that is very weak Bronsted-Lowry base. Carbon monoxide forms a strong adduct with beryllium fluoride (BF3).
References
- Carey, Francis A. (2003). Organic Chemistry (5th ed.). Boston: McGraw-Hill. ISBN 0-07-242458-3.
- IUPAC (1997). “Lewis acid”. Compendium of Chemical Terminology (2nd ed.) (the “Gold Book”). Blackwell Scientific Publications. doi:10.1351/goldbook.L03508
- Jensen, W.B. (1980). The Lewis Acid-Base Concepts: An Overview. New York: Wiley. ISBN 0-471-03902-0.
- Lepetit, Christine; Maraval, Valérie; Canac, Yves; Chauvin, Remi (2016). “On the nature of the dative bond: Coordination to metals and beyond. The carbon case”. Coordination Chemistry Reviews. 308: 59–75. doi:10.1016/j.ccr.2015.07.018
- Lewis, Gilbert Newton (1923). Valence and the Structure of Atoms and Molecules. American Chemical Society. Monograph series. New York, New York, U.S.A.: Chemical Catalog Company. ISBN 9780598985408.