The gas \(\mathrm{O}_{3}\) (ozone) can not oxidise (a) \(\mathrm{KI}\) (b) \(\mathrm{FeSO}_{4}\) (c) \(\mathrm{KMnO}_{4}\) (d) \(\mathrm{K}_{2} \mathrm{MnO}_{4}\)

Chemical oxidation involves a process where a substance loses electrons and increases its oxidation state. This process is crucial in many chemical reactions, especially in the context of energy release. For instance, when a metal corrodes or an organic substance combusts, oxidation is taking place.

To understand this using the exercise as an example, ozone \textbf{(O\(_3\))} acts as an oxidizer for substances like \textbf{KI} and \textbf{FeSO\(_4\)} by accepting electrons from iodine and iron respectively. As these elements lose electrons, their oxidation states increase, which reflects their oxidation. In contrast, manganese in \textbf{KMnO\(_4\)} is already in its highest oxidation state (+7) and cannot be oxidized any further, making it resistant to oxidation by ozone.

Oxidizing agents, also known as oxidants, are substances that gain electrons during a chemical reaction and in the process oxidize another substance. They have the ability to take electrons away from other substances, thereby increasing the oxidation state of those substances.

Ozone, mentioned in the exercise, is a powerful oxidizing agent due to its ability to accept electrons readily from other substances like potassium iodide (KI) and iron sulfate (FeSO\(_4\)). The strength of an oxidizing agent can be judged by how readily it accepts electrons, which is crucial in many industrial and biological processes.

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two substances. In every redox process, one substance is oxidized (loses electrons) and another is reduced (gains electrons). These reactions are fundamental to numerous processes including cellular respiration, battery operations, and metallurgical procedures.

In the context of our exercise, when ozone acts on KI or FeSO\(_4\), it is involved in a redox reaction where iodine and iron are oxidized while ozone is reduced. However, with KMnO\(_4\), a redox reaction with ozone wouldn't result in any change as manganese is already maximally oxidized, exhibiting the importance of identifying the capable redox partners in chemical reactions.

Oxidation numbers, also known as oxidation states, are used to keep track of electron transfer in redox reactions. They represent the theoretical charge an atom would have if all its bonds were ionic. Oxidation numbers are assigned based on a set of rules that consider electronegativity and the molecule's structure.

In the exercise, we assigned oxidation numbers to iodine, iron, and manganese to determine their susceptibility to further oxidation. For example, the transition of iron from an oxidation state of +2 in FeSO\(_4\) to a higher state indicates its ability to be oxidized. The concept of oxidation numbers allows chemists to predict reaction outcomes and balance redox equations with precision.