Question

Question: Show by suitable net ionic equations that each of the following species can act as a Bronsted-Lowry base:

\({\rm{a) H}}{{\rm{S}}^ - }\)

\({\rm{b)\;PO}}_4^{3 - }\)

\({\rm{c) NH}}_2^ - \)

\({\rm{d)}}{{\rm{C}}_2}{{\rm{H}}_5}{\rm{OH}}\)

\({\rm{e)}}{{\rm{O}}^{2 - }}\)

\({\rm{f) }}{{\rm{H}}_2}{\rm{PO}}_4^ - \)

Short answer

The net ionic equations are as follows

Step-by-step solution

Bronsted-Lowry's Concept

The Bronsted-Lowry's acid is proton \(\left( {{H^ + }} \right)\) donor and its base is proton acceptor.

Suitable net ionic equations

If a base is added to water, protons moved from water molecules to base molecules which is known as a base ionization reaction. As per the nature of solute dissolved in it, water can act as either an acid or a base.