Lewis theory sees the acid-base equilibrium as the movement of the electron pair from one compound to another. The hallmark of the Lewis acid-base equilibria is making of a new donor-acceptor bond between an element that was a donor of electrons and an elements that accepts those electrons.
Thus, the acid in the Lewis theory is defined as an acceptor of the electrons. The base is a donor of electrons. It’s very important to remember this difference between the Lewis theory and the Bronsted-Lowry! Otherwise, it’s really easy to get all confused on the test! 😉
So, let’s look at a couple of examples of the Lewis acid-base equilibria:
In the examples above, I have the same base but two different acids. The amine (nitrogen-containing base) donates its spare electron pair on the nitrogen to make a new bond with boron or carbon. Notice, that donating electrons to another atom will decrease the electron density on the donor making it neutral if it was negatively charged or positively charged if it was neutral. Likewise, an acceptor of electron density will change its charge from neutral to negative or from positive to neutral.
Generally, we’re not going to classify the Lewis acids as strong or weak as it’s quite difficult to do. So, for the most part, we’re going to assume that all Lewis acids we work with are strong enough for our purposes. We’ll also see an very limited number of those, so you don’t really need to be concerned with their strength. The main focus will be always to identify one.