The 1920s were a boom time for the study of
acids and
bases in
chemistry. Brønsted and Lowry had both independently confirmed a
new theory of the nature of those chemicals to refine the older
Arrhenius Theory. Allowing the study of odd acid and base types, like
ternary acids and
polyprotic acids, that could not be explained or analyzed according to Arrhenius Theory. At the same time in stepped
Gilbert Lewis, whose own comprehensive study of
electrons yielded one final interpretation of the nature of
acids and
bases, the most accurate yet. Approaching from a different direction, this theory was the last to fully explain the existence of all chemicals exhibiting
acidic and
basic properties.
According to Lewis Theory, an acid is any molecule that accepts a pair of electrons. On the other hand, a base is any molecule that donates an electron pair. Reversing the positions of donation and acceptance from Brønsted-Lowry Theory, Lewis Theory gets at the very root of the acid/base effect, the sharing of electrons. It does not require that an electron pair must transfer from one atom to another, as Bronsted-Lowry specifies, rather it is enough that an electron pair originally on one atom must be shared. Under this circumstance, the bond is called a coordinate covalent bond.
Lewis acids and bases, somewhat uncommon, look strikingly different from standard expectations. They do not contain hydrogen or hydroxide as Arrhenius Theory mandates, they do not even have hydrogen protons to share as required by Bronsted-Lowry Theory. Instead, all that's necessary is an atom capable of accepting electrons and one willing to share them, usually covalent compounds of the group III elements (best explained under periodic table). An example of a typical Lewis neutralization reaction between the acid Boron trifloride (BF3) and ammonia (NH3).
..
:F: H
.. | |
:F--B + :N--H
" | |
:F: H
"
Results in...
..
:F: H
.. | |
:F--B--N--H
" | |
:F: H
"