For each pair of molecules or ions, select the stronger base, and write its Lewis structure. (a) \(\mathrm{CH}_{3} \mathrm{~S}^{-}\)or \(\mathrm{CH}_{3} \mathrm{O}^{-}\) (b) \(\mathrm{CH}_{3} \mathrm{NH}^{-}\)or \(\mathrm{CH}_{3} \mathrm{O}^{-}\) (c) \(\mathrm{CH}_{3} \mathrm{COO}^{-}\)or \(\mathrm{OH}^{-}\) (d) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{O}^{-}\)or \(\mathrm{H}^{-}\) (e) \(\mathrm{NH}_{3}\) or \(\mathrm{OH}^{-}\) (f) \(\mathrm{NH}_{3}\) or \(\mathrm{H}_{2} \mathrm{O}\) (g) \(\mathrm{CH}_{3} \mathrm{COO}^{-}\)or \(\mathrm{HCO}_{3}^{-}\) (h) \(\mathrm{HSO}_{4}^{-}\)or \(\mathrm{OH}^{-}\) (i) \(\mathrm{OH}^{-}\)or \(\mathrm{Br}^{-}\)

Understanding the comparative basicity of various molecules and ions is crucial when predicting their chemical behavior. Basicity refers to the ability of a species to donate an electron pair. In general, a stronger base has a greater tendency to donate its electron pair to a proton (H⁺).

When comparing the basicity of two species, such as \textbf{CH}\(_3\)S^- and \textbf{CH}\(_3\)O^-, we look at factors like atom size and electronegativity. Since sulfur is larger than oxygen, \textbf{CH}\(_3\)S^- can better accommodate the negative charge and is thus a stronger base compared to \textbf{CH}\(_3\)O^-. This concept is similarly applied to other comparisons in the exercise.

For example, in comparing hydride (H^-) with an ethoxide ion (\textbf{CH}\(_3\)\textbf{CH}\(_2\)O^-), we consider that the hydride ion, being smaller and less electronegative, is more willing to donate its electron pair and thus acts as a stronger base.

In each case, drawing the Lewis structure can visually highlight the presence and disposition of lone electron pairs, making the comparison of basicities more understandable.

Resonance stabilization significantly influences the basicity of molecules. A molecule or ion is resonance-stabilized when it has a structure that allows for the delocalization of electrons across multiple atoms. This delocalization leads to a stabilization of the molecule or ion, making it less likely to donate its electron pair.

For instance, the carboxylate ion (\textbf{CH}\(_3\)\textbf{COO}^-) has resonance structures that delocalize the negative charge between two oxygen atoms. This dispersal of charge over more than one atom makes the ion less reactive as a base, compared to \textbf{OH}^-, which doesn't have such delocalization.

Considering another example, bicarbonate (HCO₃^-) versus carboxylate (\textbf{CH}\(_3\)\textbf{COO}^-), the negative charge on bicarbonate is not as delocalized as in the carboxylate ion, which has two oxygen atoms to share the negative charge, making bicarbonate a slightly stronger base due to less resonance stabilization.

Electronegativity plays a pivotal role in determining basicity. It is the measure of an atom's ability to attract electrons towards itself. In a basicity comparison, atoms with lower electronegativity are typically better electron pair donors, hence stronger bases.

When comparing oxygen and nitrogen, such as in the case of \textbf{CH}\(_3\)\textbf{NH}^- and \textbf{CH}\(_3\)\textbf{O}^-, oxygen, being more electronegative, holds onto its electrons more tightly than nitrogen. This makes \textbf{CH}\(_3\)\textbf{O}^- a stronger base than \textbf{CH}\(_3\)\textbf{NH}^- because it better utilizes its lone pair for bonding with a proton.

Additionally, the size of the atom is also a factor; larger atoms with the same charge as smaller atoms are typically less electronegative. For example, this is why bromide (Br^-) is a weaker base than hydroxide (\textbf{OH}^-). Despite having a full negative charge, the larger size and lower electronegativity of bromine compared to oxygen means that bromide is less efficient in donating its lone pair.

The tendency of an atom within a molecule or ion to donate a lone electron pair is the cornerstone of basicity. This property is influenced by various factors, including the atom's electronegativity, the size of the atom, its resonance effects, and the overall molecular structure.

Base strength increases with the availability of the lone pair for donation. For example, in ammonia (\textbf{NH}\(_3\)), the lone pair on the nitrogen is readily available, making it a stronger base than water (\textbf{H}\(_2\)\textbf{O}), where the lone pairs are more involved in hydrogen bonding, thus less available for donation.

Furthermore, in a species like hydroxide (\textbf{OH}^-), the negatively charged oxygen atom has a high electron density which makes it more inclined to donate its lone pair compared to a neutral atom like in water. Properly visualizing the Lewis structures of these species can help students understand the location and availability of lone pairs which translates to the molecule's or ion's basic potential.