91Ó°ÊÓ

Redox reactions, short for reduction-oxidation reactions, are chemical processes where the oxidation states of atoms change. These reactions involve the transfer of electrons between substances. Understanding these electron movements is critical for analyzing redox processes.
The first step in redox reaction analysis is to assign oxidation states to each element in both the reactants and the products. This allows us to identify which elements are oxidized and which are reduced.
Consider reaction (a): \(\mathrm{Zn(s)}+2 \mathrm{NO}_{3}^{-}(\mathrm{aq})+4\mathrm{H}^{+}(\mathrm{aq}) \longrightarrow \mathrm{Zn}^{2+}(\mathrm{aq})+2 \mathrm{NO}_{2}(\mathrm{g})+2 \mathrm{H}_{2}\mathrm{O}(\ell)\). Here, zinc (Zn) is oxidized, changing from an oxidation state of 0 to +2. The nitrogen in \(\mathrm{NO}_3^-\) is reduced from +5 to +4. Both oxidation and reduction occur, thus classifying it as a redox reaction.
Similarly, in reaction (c), calcium (Ca) undergoes oxidation and the hydrogen from water is reduced. Hence, reaction (c) is also a redox reaction.

Oxidation states are hypothetical charges that atoms would have if all bonds were ionic. They are utilized in determining how electrons are transferred in reactions. Learning to calculate these states is essential for understanding many chemical reactions.

• The oxidation state of an atom in its elemental form is always 0. For instance, in \(\mathrm{Ca(s)}\) and \(\mathrm{Zn(s)}\), both have an oxidation state of 0.
• For monoatomic ions, the oxidation state is equal to the charge of the ion. Thus, \(\mathrm{Zn}^{2+}\) has an oxidation state of +2.
• In compounds, the oxidation states must sum up to the compound's net charge. For example, \(\mathrm{H}_2\mathrm{O}\) has a net oxidation state of 0, balancing +1 of H and -2 of O.

Assigning oxidation states helps determine which atoms are gaining or losing electrons in a reaction. These changes define whether a reaction is a redox process.

Unlike redox reactions, acid-base reactions involve the transfer of protons (\(\mathrm{H}^+\)). These occur without changing the oxidation states of the elements involved. An acid-base reaction typically features an acid donating a proton to a base.
In reaction (b), \(\operatorname{Zn(OH)}_{2}(\mathrm{s})+\mathrm{H}_{2}\mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow \mathrm{ZnSO}_{4}(\mathrm{aq})+2 \mathrm{H}_{2}\mathrm{O}(\ell)\), no change in oxidation states occurs. Zinc remains in the +2 state throughout. Instead, this reaction exemplifies an acid-base process, where sulfuric acid (\(\mathrm{H}_2\mathrm{SO}_4\)) donates \(\mathrm{H}^+\) to form water with \(\mathrm{OH}^-\) from zinc hydroxide.
This type of reaction highlights the basic interplay of acids and bases in chemistry, which contrasts with the electron transfer characterizing redox reactions.