Why does the Lewis definition of acids specify a PAIR of electrons and not a single electron?

The Lewis definition of an acid is: "a compound or ionic species which can accept an electron pair from a donor compound."

Why does it specify a pair of electrons and not just a single electron?

I initially thought maybe covalent bonds require a pair of electrons but it says that is not the case here : https://en.wikipedia.org/wiki/Covalent_bond#One-_and_three-electron_bonds

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For reference:

Question from textbook : Which of the following species can act as a Lewis acid?

$\ce{NH3}$
$\ce{F-}$
$\ce{H2 O}$
$\ce{NH4+}$
$\ce{BF3}$ (answer)

I thought the answer was $\ce{NH4+}$ as it could accept an electron, but the Lewis definition specifies accepting a pair of electrons.

Thanks.

Answer

It partially has something to do with the history of acid-base definitions. We started with the Arrhenius definition which was based on the generation of $\ce{H+}$ or $\ce{OH-}$ in aqueous solution, and then to the Bronsted-Lowry definition which moved towards acceptance or donation of $\ce{H+}$. Finally came the Lewis Definition which removed $\ce{H+}$ and left us only with electron pairs, to be consistent with the past two definitions. We couldn't just say electrons because that would have included a large class of reactions called redox reactions which are clearly not the same as acid-base.

Now, with respect to your answer to the multiple-choice question:

I thought the answer was $\ce{NH4+}$ as it could accept an electron, but the Lewis definition specifies accepting a pair of electrons.

There are two points that I would like to draw your attention to:

1. You're right in that $\ce{NH4+}$ is an acid (more specifically we call it a conjugate acid), but by the two older definitions.


2. However, $\ce{NH4+}$ cannot accept even one electron in this case. If you draw the Lewis structure, you'll find that although there is a net positive charge, all hydrogens have 2 electrons and the central Nitrogen has eight, fulfilling the octet rule.