Classify the following as acid-base reactions or oxidation-reduction reactions: (a) \(\mathrm{Na}_{2} \mathrm{S}(a q)+2 \mathrm{HCl}(a q) \longrightarrow 2 \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{S}(g)\) (b) \(2 \mathrm{Na}(s)+2 \mathrm{HCl}(a q) \longrightarrow 2 \mathrm{NaCl}(a q)+\mathrm{H}_{2}(g)\) (c) \(\mathrm{Mg}(s)+\mathrm{Cl}_{2}(g) \longrightarrow \mathrm{MgCl}_{2}(s)\) (d) \(\mathrm{MgO}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{MgCl}_{2}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) (e) \(\mathrm{K}_{3} \mathrm{P}(s)+2 \mathrm{O}_{2}(g) \longrightarrow \mathrm{K}_{3} \mathrm{PO}_{4}(s)\) (f) \(3 \mathrm{KOH}(a q)+\mathrm{H}_{3} \mathrm{PO}_{4}(a q) \rightarrow \mathrm{K}_{3} \mathrm{PO}_{4}(a q)+3 \mathrm{H}_{2} \mathrm{O}(l)\)

The concept of proton exchange is fundamental to understanding acid-base reactions. Protons, or hydrogen ions (\textbf{H}\(^+\)), are exchanged between substances. Acids are proton donors, while bases are proton acceptors. In the context of our textbook exercise, when looking at reaction (a), \textbf{Na}\(_{2}\)S reacts with \textbf{HCl} to form \textbf{NaCl} and \textbf{H}\(_{2}\)S. In this process, the \textbf{S} in \textbf{Na}\(_{2}\)S accepts protons from the \textbf{HCl}, releasing \textbf{H}\(_{2}\)S gas and leaving behind \textbf{NaCl} in solution.

Similarly, reaction (d) involves \textbf{MgO} reacting with \textbf{HCl}, producing \textbf{MgCl}\(_{2}\) and \textbf{H}\(_{2}\)O. \textbf{MgO} acts as a base accepting protons from \textbf{HCl}, an acid, leading to the formation of water, a classic example of neutralization. An easy way to recognize these reactions is by checking if there's formation or consumption of water, a common product in acid-base reactions, where water is formed by the combination of hydrogen ions (protons) from the acid and hydroxide ions from the base.

In oxidation-reduction (redox) reactions, the change in oxidation states is a critical indicator to observe. Oxidation states refer to the hypothetical charge that an atom would have if all bonds to atoms of different elements were completely ionic. A rise in oxidation state signifies oxidation, while a reduction indicates, appropriately enough, reduction. In reaction (b) of the exercise, we start with solid sodium (\textbf{Na}) which has an oxidation state of 0, since it's an uncombined element. After it reacts with \textbf{HCl}, \textbf{Na} becomes \textbf{NaCl}, and its oxidation state changes from 0 to +1. Concurrently, hydrogen (\textbf{H}), originally +1 in \textbf{HCl}, ends up as \textbf{H}\(_{2}\) gas, with an oxidation state of 0. Here, \textbf{Na} is oxidized and \textbf{H} is reduced, showcasing the classic electron transfer of a redox reaction.

Another instance is reaction (c), where magnesium (\textbf{Mg}) changes from an oxidation state of 0 to +2 in \textbf{MgCl}\(_{2}\), and chlorine (\textbf{Cl}\(_{2}\)) changes from 0 to -1. This change in oxidation state across different substances within the same reaction is the hallmark of a redox process.

Electron transfer is the movement of electrons from one atom or molecule to another. It is the physical basis underlying oxidation-reduction reactions. During these reactions, the losing electron species undergoes oxidation, and the gaining electron species undergoes reduction. In our textbook example, reaction (b) demonstrates this as \textbf{Na} loses electrons to become \textbf{Na}\(^+\), and \textbf{H}\(^+\) gains those electrons to become \textbf{H}\(_{2}\) gas.

Moreover, reaction (e) also showcases electron transfer. When \textbf{K}\(_{3}\)P is exposed to \textbf{O}\(_{2}\), the phosphorus (P) goes from an oxidation state of 0 to +5 in \textbf{K}\(_{3}\)\textbf{PO}\(_{4}\), and oxygen (O) is reduced from 0 to -2. Electron transfer reactions are essential in the world of chemistry as they are involved in processes such as metabolism, photosynthesis, and battery operation. Recognizing these reactions requires identifying shifts in oxidation states, often accompanied by energy changes, such as heat, light, or electrical energy.