Using your results from Exercise 133, place the species in...

Using your results from Exercise $133,$ place the species in each of the following groups in order of increasing base strength.
a. $\mathrm{OH}^{-}, \mathrm{SH}^{-}, \mathrm{SeH}^{-}$
b. $\mathrm{NH}_{3}, \mathrm{PH}_{3}$
c. $\mathrm{NH}_{3}, \mathrm{HONH}_{2}$

Key Concepts

Basicity

Basicity refers to the ability of a species to accept a proton (H?). A stronger base has a greater tendency to donate its electron pair to bind with a proton. The concept is central to understanding reaction mechanisms in acid-base chemistry and is evaluated by considering both the inherent electronic properties of the base and the stability of its conjugate acid.

Conjugate Base Stability

The strength of a base is inversely related to the stability of its conjugate acid: a more stable conjugate acid corresponds to a weaker base, because the proton is held more tightly. Evaluating conjugate base stability involves considering factors such as charge distribution, resonance stabilization, and electron density on the atom that accepts the proton.

Periodic Trends

Periodic trends, including changes in electronegativity, atomic size, and polarizability, play a crucial role in the variation of basicity across different elements. Generally, as you move down a group, the larger atomic radius and increased polarizability can make bases less strong because the electron density is more diffuse, affecting the molecule’s ability to attract protons.

Electronegativity and Atomic Characteristics

Electronegativity reflects an atom’s ability to attract electrons in a bond. A highly electronegative atom holds its electrons more tightly, which can reduce the availability of the electron pair for protonation, thereby decreasing basic strength. In contrast, atoms with lower electronegativity and greater size often exhibit increased basicity due to more diffuse electron density and higher polarizability.

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