Which amongst the following compounds/species is least basic?

Solution & Explanation

\textbf{Answer:} (d) $(\mathrm{H}{}_2\mathrm{N}){}_2\mathrm{C}{}^{+}-\mathrm{OH}$ \textbf{Solution:} Basicity depends on the availability of a lone pair on nitrogen. In guanidine (a), resonance delocalisation makes its conjugate acid (guanidinium) highly stabilised, so guanidine is very strongly basic. Urea (c) is weakly basic because the C=O group withdraws electron density. The cationic species (b) and (d) are already protonated, and a positively charged species is the poorest base. In (d) the positive charge sits on a carbon flanked by an $-\mathrm{OH}$ group, whose electron-withdrawing $-I$ effect pulls electron density further away from the nitrogens, leaving essentially no available lone pair to donate. Hence (d) is the least basic of the set.