Use Table \(14.3\) to help order the following bases from strongest to weakest. $$ \mathrm{NO}_{3}^{-}, \quad \mathrm{H}_{2} \mathrm{O}, \quad \mathrm{NH}_{3}, \quad \mathrm{C}_{5} \mathrm{H}_{5} \mathrm{~N} $$

Understanding the concept of base strength is essential in chemistry, particularly when studying acid-base reactions. Base strength refers to a substance’s ability to accept a proton (H+) during a chemical interaction. It indicates how readily the base will pull away a proton from an acid in an acid-base reaction.

Strong bases, such as alkali metal hydroxides, are characterized by their complete dissociation in water. This results in the release of hydroxide ions (OH-), which readily combine with protons to form water. Conversely, weak bases, like ammonia (NH3), only partially dissociate, producing fewer hydroxide ions. The relative dissociation in water is a good measure of a substance's basicity; the more complete the dissociation, the stronger the base.

Proton acceptance is a fundamental behavior of bases. It’s the process where a base accepts a proton from an acid during an acid-base reaction. This property of bases is due to the presence of a lone pair of electrons that can be used to bond to a proton.

A base's tendency to accept a proton can be influenced by several factors, including the electronegativity of the atoms near the lone pair, the steric accessibility of the lone pair, and the solvent in which the reaction takes place. For example, nitrogen-containing species like ammonia have a lone pair on the nitrogen atom, making them competent proton acceptors.

The ability to donate electrons is another crucial aspect of bases. In many reactions, bases act as electron pair donors, which is a central theme of Lewis base theory.

Electron donating ability can be influenced by the base's molecular structure; for instance, bases with resonance structures can delocalize the negative charge, increasing their ability to donate electrons. Also, the presence of electron-donating groups attached to the base can enhance its electron-donating power. Strong bases are often excellent electron donors due to these attributes, which enable them to effectively stabilize the extra charge that results from taking on a proton.

When comparing the basicity of different substances, pKa and pKb values are invaluable. The pKa value measures the acidity of a compound, with lower values indicating stronger acids. Conversely, the pKb value refers to the basicity of a compound, whereby lower pKb values signify weaker bases.

Since we are concerned with bases, focusing on pKb is more relevant. The higher the pKb value, the stronger the base, as it implies a less stable conjugate acid, which readily donates its proton. To rank bases from strongest to weakest, compare their pKb values; those with higher values are stronger. In cases where only pKa values are provided, you can still determine basic strength, as the base's strength is inversely related to the pKa of its conjugate acid.