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History
G.N. Lewis had suggested in 1916 that two atoms are held together in a chemical bond by sharing a pair of electrons. When each atom contributed one electron to the bond is was called a covalent bond. When both electrons come from one of the atoms is was called a dative covalent bond or coordinate bond. The distinction is not clear-cut as the diagram at the right shows; although the ammonia molecule donates a pair of electrons to the hydrogen ion, the identity of the electrons is lost in the ammonium ion that is formed. Nevertheless, Lewis suggested that an electron-pair donor be classified as a base and an acid and an electron-pair acceptor be classified as acid.
The modern definition of a Lewis acid is an atomic or molecular species that has an empty atomic or molecular orbital of low energy (LUMO) that can accomodate a pair of electrons, as illustrated in the molecular orital diagram at the right.
Comparison with Brønsted-Lowry theory
A Lewis base is usually a Brønsted-Lowry base as it can donate a pair of electrons to a proton; the proton is a Lewis acid as it can accept a pair of electrons. The conjugate base of a Brønsted-Lowry acid is is also a Lewis base as loss of a proton from the acid leaves those electrons which were used for the A—H bond as a lone pair on the conjugate base. However, a Lewis base can be very difficult to protonate, yet still react with a Lewis acid. For example, carbon monoxide is a very weak Brønsted-Lowry base but it forms a strong adduct with BF3.
In another comparison of Lewis and Brønsted-Lowry acidity by Brown and Kanner[2], 2,6-di-t-butylpyridine reacts to form the hydrochloride salt with HCl but does not react with BF3. This example demonstrates that for pyridine bases, HCl (typically thought of as a is a Brønsted-Lowry acid) is a "stronger" acid than BF3 (a Lewis acid).
A Brønsted-Lowry acid is a proton donator, not an electron-pair acceptor.
Lewis acids
Acceptor orbitals of a Lewis acid are as in the following acid + base reactions.
1s orbital
- H+ + NH3: → NH4+
p orbitals: elements in groups 1—3
- B2H6 + 2H- → 2BH4-
- BF3 + F- → BF4-
- Al2Cl6 + 2Cl- → 2AlCl4-
d orbitals: elements in the second and lower rows of the periodic table
- AlF3 + 3F- → AlF63-
- SiF4 + 2F- → SiF62-
- PCl5 + Cl- → PCl6-
- SF4 + F- → SF5-
- Metal ions forming solvates, such as [Mg(H2O)62+, [Al(H2O)63+, etc. where the solvent is a Lewis base.
A typical example of a Lewis acid in action is in the Friedel-Crafts alkylation reaction. The key step is the acceptance by AlCl3 of a chloride ion lone-pair, forming AlCl4- and creating the strongly acidic, that is, electrophilic, carbonium ion.
- RCl +AlCl3 → R+ + AlCl4-f BF3+NH3=BF3NH3
Lewis Bases
A Lewis base is an atomic or molecular species that has an lone pair of electrons in the HOMO. Typical examples are
- compounds of N, P, As, Sb and Bi in oxidation state 3
- compounds of O, S, Se and Te in oxidation state 2, including water, ethers, ketones, sulphoxides
- molecules like carbon monoxide
Hard and soft classification
Considerations concerning the strength of acid base adducts lead R.G. Pearson to propose, in 1963, the classification of both acids and bases into hard and soft. Within each category he established an order of binding strengths such as
- hard acids: R3P << R3N, R2S << R2O
- soft acids: R2O << R3N, R2S << R3P
For example, an amine will displace a phosphine from the adduct with the acid BF3. In the same way, bases could be classified. For example, bases donating a lone pair from an oxygen atom are harder than bases donating through a nitrogen atom. Although the classification was never quantified it proved to be very useful in predicting the strength of adduct formation, using the key concepts
- hard acid — hard base interactions are stronger than hard acid — soft base or soft acid — hard base interactions.
- soft acid — soft base interactions are stronger than soft acid — hard base or hard acid — soft base interactions.
Later investigation of the thermodynamics of the interaction suggested that hard—hard interactions are enthalpy favoured, whereas soft—soft are entropy favoured. If the interaction between acid and base in solution results in an equilibrium mixture the strength of the interaction can be quantified in terms of an equilibrium constant. An alternative quantitative measure is the standard heat (enthalpy) of formation of the adduct in a non-coordinating solvent. Drago and Wayland proposed a two-parameter equation which predicts the formation of a very large number of adducts quite accurately.
- –ΔH
O(A—B) = EAEB + CACB
Value of the E and C parameters can be found in Drago et. al.[3]
Another quantitative system as been proposed, in which Lewis acid strength is based on gas-phase affinity for fluoride. [4]
References
- ^ Miessler, L. M., Tar, D. A., (1991) p166 - Table of discoveries attributes the date of publication/release for the Lewis theory as 1923.
- ^ Brown HC and Kanner B. "Preparation and Reactions of 2,6-Di-t-butylpyridine and Related hindered Bases. A case of Steric Hindrance twoard the Proton." J. Am. Chem. Soc. 88, 986 (1966).
- ^ Drago, R.S; Wong, N.; Bilgrien, C.; Vogel, C.; E and C parameters from Hammett substituent constants and use of E and C to understand cobalt-carbon bond energies Inorg. Chem., (1987) 26, 9–14.
- ^ Christe, K.O.; Dixon, D.A.; McLemore, D.; Wilson, W.W.; Sheehy, J.A.; and Boatz, J.A. (2000). "On a quantitative scale for Lewis acidity and recent progress in polynitrogen chemistry". Journal of Fluorine Chemistry 101 (2): 101, 151–153. ISSN 0022-1139.
Further reading
Jensen, W.B. (1980). The Lewis acid-base concepts : an overview. New York: Wiley. ISBN 0471039020.
Yamamoto, Hisashi (1999). Lewis acid reagents : a practical approach. New York: Oxford University Press. ISBN 0198500998.
See also
