A hydride (a) of group 15 element is distinctly basic and has unexpectably high boiling point. It reacts with \(\mathrm{NaOCl}\) to give another hydride (b) which is a strong reducing agent and is used in organic analysis. A and \(\mathrm{B}\) are (a) \(\mathrm{NH}_{3}, \mathrm{~N}_{2} \mathrm{H}_{4}\) (b) \(\mathrm{PH}_{3}, \mathrm{P}_{2} \mathrm{H}_{4}\) (c) \(\mathrm{NH}_{3}, \mathrm{NH}_{4} \mathrm{Cl}\) (d) \(\mathrm{AsH}_{3}, \mathrm{As}_{2} \mathrm{H}_{4}\)

Basicity refers to the capacity of a compound to donate a pair of electrons, acting as a base. In the context of Group 15 hydrides, such as amides (NH₃), phosphine (PH₃), arsine (AsH₃), stibine (SbH₃), and bismuthine (BiH₃), basicity is an important chemical property that influences their reactivity. Each of these hydrides contains a lone pair of electrons on the central atom, which can be offered to electron acceptors.

The basicity typically decreases as we move down the group, primarily due to the increasing size of the central atom which makes the lone pair less accessible to potential proton donors. Additionally, as the electronegativity decreases down the group, the hydrogen atoms become less positive, making it harder for them to attract donor species. Ammonia (NH₃) stands out in Group 15 with its significantly higher basicity. This is largely due to nitrogen's relatively high electronegativity, which keeps the lone pair electrons closer to the atom, making them readily available to bond with hydrogen ions.

Hydrogen bonding is an attractive force that occurs between a hydrogen atom, which is covalently bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, and another electronegative atom with a lone pair of electrons. This interaction is stronger than a typical van der Waals force but weaker than a true chemical bond.

In the case of ammonia (NH₃), hydrogen bonds form because the nitrogen atom has a lone pair of electrons and the hydrogen atoms have a partial positive charge due to nitrogen's electronegativity. These hydrogen bonds are the reason ammonia has an unexpectedly high boiling point compared to its heavier congeners like phosphine and arsine, where such bonding is absent or much weaker due to the lower electronegativities of phosphorus and arsenic. Hydrogen bonding not only affects the physical properties such as the boiling point of these hydrides but also influences their solubility in water and their chemical reactivity.

Hydrazine (N₂H₄) is known for its role as a reducing agent, meaning it readily donates electrons and becomes oxidized. This property makes it useful in various chemical processes, particularly in organic synthesis and analytical chemistry. The molecule contains two nitrogen atoms that are each bonded to two hydrogen atoms, and these nitrogens share a single bond, giving N₂H₄ a rich supply of electrons that can be donated to other substances.

Hydrazine's reducing power is employed in the synthesis of pharmaceuticals, agricultural chemicals, and polymers. It is also used in rocket fuel formulations due to its ability to rapidly decompose and release large amounts of energy. In industry, hydrazine is often used in its anhydrous form, but in laboratories, it is typically handled as a hydrate to reduce risk, since hydrazine is highly toxic and dangerously unstable in its pure form.