Write all the species (except water) that are present in a phosphoric acid solution. Indicate which species can act as a Bronsted acid, which as a Bronsted base, and which as both a Bronsted acid and a Bronsted base.

In Bronsted-Lowry theory, an acid is defined as a substance that donates a proton (hydrogen ion, represented as \(H^+\)), while a base is a substance that accepts a proton. This concept is crucial to understanding many reactions in chemistry, especially those occurring in aqueous solutions.

For example, in the case of phosphoric acid (\(H_3PO_4\)), when it is dissolved in water, it donates a proton to the water molecule, which is acting as a base. This first step produces the dihydrogen phosphate ion \(H_2PO_4^-\), which is also amphoteric. An amphoteric species can behave either as a Bronsted acid or as a Bronsted base. It can further lose a proton forming \(HPO_4^{2-}\), and finally, it yields \(PO_4^{3-}\) upon losing its third proton. These reactions showcase the substance acting once as a Bronsted acid and then potentially as a Bronsted base, displaying the dynamic nature of proton donation and acceptance in aqueous solutions.

A dissociation reaction in chemistry involves the splitting of a compound into two or more particles, often ions. This can occur physically, as when a solid ionic compound dissolves in water, or it can be a chemical reaction wherein covalent bonds are broken, and the compound splits into ions or smaller molecules.

When discussing phosphoric acid, this concept comes into play in its ability to dissociate in water. As we've seen, phosphoric acid undergoes a stepwise dissociation, releasing protons one at a time and forming different ions at each step. This results in a mixture of species, depending on the solution's pH: primarily \(H_3PO_4\) in very acidic conditions, the ions \(H_2PO_4^-\), \(HPO_4^{2-}\) in intermediate pH, and \(PO_4^{3-}\) at very basic conditions. Understanding these reactions is fundamental for students when analyzing the behavior of phosphoric acid in different environments.

An amphoteric species is one that can act as both an acid and a base, depending on its chemical environment. This flexibility is due to the presence of both a site that can donate a proton and another that can accept a proton.

With phosphoric acid, for instance, we encounter the amphoteric species \(H_2PO_4^-\) and \(HPO_4^{2-}\). These ions can act as acids when they donate a proton, converting to \(HPO_4^{2-}\) and \(PO_4^{3-}\), respectively. Conversely, they can accept a proton when acting as bases, moving back to \(H_3PO_4\) and \(H_2PO_4^-\). Understanding amphoteric behavior is essential because it dictates how substances may behave in a given pH range and allows for better prediction of reactions in biological systems, industrial applications, and environmental chemistry.