# Why can water be a Lewis base?

Jul 15, 2016

Because it follows the definition.

A Lewis base can donate an electron pair to a Lewis acid (which accepts electron pairs, being the opposite of a Lewis base), and water has two electron pairs.

Therefore, it is entirely capable of being a Lewis base in the presence of a strong enough Lewis acid. It just needs the right situation.

Simple examples of suitable Lewis acids are ${\text{H"_2"SO}}_{4}$ ($\text{pKa} \approx - 3$), ${\text{NH}}_{4}^{+}$ ($\text{pKa} \approx 9.26$), and other acids with a $\text{pKa}$ less than that of water ($\text{pKa} \approx 15.7$), since:

• An acid with a lower $\text{pKa}$ is a stronger Bronsted acid.
• A Lewis base in the presence of a Bronsted acid may grab its proton, as Bronsted acids are proton donors.
• A stronger Bronsted acid has weaker bonds with its protons than a weaker Bronsted acid would.
• Weaker bonds with their protons allow them to donate their protons more easily.
• Protons being donated more easily to water means water's electron pair donation is more favored.

Lastly, acidity is relative - you don't know for sure how strong or weak an acid is unless you are in context, or if it's so strong that it's known to be one of the strongest acids (like ${\text{H"_2"SO}}_{4}$, ${\text{HNO}}_{3}$, $\text{HCl}$, $\text{HBr}$, $\text{HI}$, ${\text{HClO}}_{3}$, and ${\text{HClO}}_{4}$, the typically-labeled "strong acids").