The hypochlorite ion, \(\mathrm{ClO}^{-},\) acts as a weak base. (a) Is \(\mathrm{ClO}^{-}\) a stronger or weaker base than hydroxylamine? (b) When ClO \(^{-}\) acts as a base, which atom, Cl or \(\mathrm{O}\), acts as the proton acceptor? (c) Can you use formal charges to rationalize your answer to part (b)?

To compare the basicity of hypochlorite ion (\(\mathrm{ClO}^{-}\)) and hydroxylamine, we should look at their conjugate acids (\(\mathrm{HClO}\) and \(\mathrm{NH_3OH^{+}}\)) and compare their acidity. The acidity can be compared by looking at the stability of the conjugate base. A more stable conjugate base means a stronger acid, and thus a weaker base. In general, more electronegative atoms can stabilize a negative charge better. For hydroxylamine, the conjugate base is \(\mathrm{NH_2OH}^{-}\) where the negative charge is on the nitrogen atom. In the case of hypochlorite ion, the conjugate base is \(\mathrm{HClO}\), and when losing a proton, it becomes \(\mathrm{ClO}^{-}\) with a negative charge on the oxygen atom. Oxygen is more electronegative than nitrogen. Therefore, \(\mathrm{ClO}^{-}\) is a weaker base compared to hydroxylamine, as its conjugate acid (\(\mathrm{HClO}\)) is a stronger acid.

To identify which atom (Cl or O) in \(\mathrm{ClO}^{-}\) acts as the proton acceptor, we must consider the Lewis structure and resonance structures. The oxygen atom has one lone pair of electrons making it the most likely candidate for accepting a proton. Moreover, the oxygen atom is more electronegative than the chlorine atom, which means it is more likely to accept a proton as a Lewis base. So, the oxygen atom in \(\mathrm{ClO}^{-}\) acts as the proton acceptor.

To rationalize our choice of the oxygen atom as the proton acceptor using formal charges, we can consider the formal charges of both atoms in \(\mathrm{ClO}^{-}\). The formal charge of an atom is calculated based on the valence electrons, the number of electrons in lone pairs, and half the number of shared electrons in bonds. The formal charge of oxygen in \(\mathrm{ClO}^{-}\) is: Formal charge (O) = Valence electrons (O) - Lone pair electrons (O) - (1/2) × Shared electrons (O) Formal charge (O) = 6 - 4 - (1/2) × 2 = -1 The formal charge of chlorine in \(\mathrm{ClO}^{-}\) is: Formal charge (Cl) = Valence electrons (Cl) - Lone pair electrons (Cl) - (1/2) × Shared electrons (Cl) Formal charge (Cl) = 7 - 6 - (1/2) × 2 = 0 Since the oxygen atom has a negative formal charge, it makes more sense for the oxygen atom to act as the proton acceptor rather than the chlorine atom, which has a neutral formal charge. Thus, the formal charges support our choice of the oxygen atom as the proton acceptor in \(\mathrm{ClO}^{-}\).