Account for the following: (a) AgCI dissolves in ammonia. (b) White phosphorus is more reactive than red phosphorus. (c) \(\mathrm{NH}_{3}\) does not act as a reducing agent. (d) \(\mathrm{H}_{3} \mathrm{PO}_{3}\) is a stronger acid than \(\mathrm{H}_{3} \mathrm{PO}_{4}\).

Understanding the dissolution of ionic compounds in solution often involves the key concept of complex ion formation. Complex ions are species formed from a central metal ion surrounded by molecules or ions called ligands, which are able to donate a pair of electrons. A classic example is when silver chloride (AgCl) dissolves in ammonia (NH₃). The ammonia molecules act as ligands and bind to the silver ion, forming the diamminesilver complex ion.
This process can be described by the chemical equation:
\( AgCl(s) + 2NH_3(aq) \rightarrow [Ag(NH_3)_2]^+(aq) + Cl^-(aq) \).
The diamminesilver ion \( [Ag(NH_3)_2]^+ \) is much more soluble in water than AgCl, leading to the dissolution of AgCl in ammonia. This concept is crucial in various applications, including analytical chemistry where complex ion formation is used to separate or identify specific metal ions.

Phosphorus displays allotropic forms, with white and red phosphorus being the most commonly discussed ones in inorganic chemistry. White phosphorus exhibits high reactivity, which stems from its molecular arrangement. Molecules of white phosphorus are composed of P₄ tetrahedra. This form is less stable due to the presence of significant strain in the P-P-P angles and weaker van der Waals forces between the molecules. The strain makes the bonds within the molecule easier to break and hence, white phosphorus is more reactive.
In contrast, red phosphorus presents as a polymeric structure formed by interconnected phosphorus atoms. This network makes the molecules more stable and substantially less reactive compared to their white counterpart. This fundamental difference in structure dictates their distinct reactivity and precautions for handling, with white phosphorus being much more hazardous.

Ammonia (NH₃), contrary to many other common substances, does not typically behave as a reducing agent in chemical reactions. A reducing agent usually donates electrons, reducing another reactant and in the process, becomes oxidized. However, ammonia tends to participate in acid-base reactions where it accepts a proton due to its lone pair of electrons on the nitrogen atom. This makes it a good nucleophile or a base rather than a reducing agent. For ammonia to act as a reducing agent, it would have to lose its lone pair which does not occur easily under most conditions. Thus, under normal circumstances such as in aqueous solutions, it generally remains unoxidized.

Acid strength in inorganic chemistry is an important concept, and it is commonly assessed by the acidity constant (Ka) or the readiness with which an acid donates protons (H⁺). Comparing the acid strength of phosphoric acid (H₃PO₄) and phosphorous acid (H₃PO₃), one finds that H₃PO₃ is a stronger acid, a fact that may seem counterintuitive at first glance. The stronger acidity of H₃PO₃ is tied to the number of protons it can actually lose as H⁺, which is two. It's more willing to part with these protons compared to H₃PO₄, which although has three potentional protons, does not easily lose all of them. The comparison reveals insights into acid behavior that is explained by molecular structure, which influences how acids dissociate and contribute protons to a solution.