Dichromate ion \(\left(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\right)\) in acidic solution is a good oxidizing agent. Which of the following oxidations can be accomplished with dichromate ion in acidic solution? Explain. (a) \(\operatorname{sn}^{2+}\left(\text { aq) to } \operatorname{Sn}^{4+}(\text { aq })\right.\) (b) \(\mathrm{I}_{2}(\mathrm{s})\) to \(\mathrm{IO}_{3}^{-}(\mathrm{aq})\) (c) \(\mathrm{Mn}^{2+}(\mathrm{aq})\) to \(\mathrm{MnO}_{4}^{-}(\mathrm{aq})\)

Understanding the oxidation state, often referred to as oxidation number, is crucial when studying redox reactions. It's a number that represents the total number of electrons an atom gains or loses when forming a compound. Essentially, it's a bookkeeping method to keep track of electron movement in chemical reactions.

Oxidation states are straightforward for pure elements, which have an oxidation state of zero because they have not gained or lost electrons. In compounds or ions, the oxidation state is more complex, often depending on the rules of electronegativity and the type of bond between atoms. For instance, in sulfate ion \(\mathrm{SO}_{4}^{2-}\), sulfur has an oxidation state of +6, while each oxygen has an oxidation state of -2. This concept explains why on oxidizing \(\mathrm{Sn}^{2+}\) to \(\mathrm{Sn}^{4+}\), the oxidation state increases from +2 to +4, signifying a loss of electrons.

Redox reactions are a family of reactions where oxidation and reduction occur. An easy way to remember this is 'OIL RIG' which stands for Oxidation Is Loss, Reduction Is Gain—referring to electrons. When an atom loses electrons, its oxidation state increases, showing oxidation. Conversely, when it gains electrons, its oxidation state decreases, indicating reduction.

In academic settings, a redox reaction is often broken down for clarity. For instance, the change from \(\mathrm{Sn}^{2+}\) to \(\mathrm{Sn}^{4+}\) is an oxidation process because tin loses electrons. Redox reactions can also involve a change in the oxidation state without a noticeable electron transfer, because they are often not free but within compounds.

In chemistry, oxidizing agents, or oxidants, are substances that can accept electrons from other species. Because they take electrons away, they are said to oxidize the other substance. The dichromate ion (\(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\)) is an excellent example of a powerful chemical oxidant you might encounter. It can accept a fair number of electrons, changing its own oxidation state in the process.

Oxidants vary in strength. Some can oxidize elements to a high oxidation state, whereas others have a lower capacity. For example, the dichromate ion can reliably oxidize elements to an oxidation state of up to +6. This means it can easily oxidize tin from \(\mathrm{Sn}^{2+}\) to \(\mathrm{Sn}^{4+}\), but it would not be effective for manganese going from \(\mathrm{Mn}^{2+}\) to \(\mathrm{MnO}_{4}^{-}\), where the oxidation state is +7.

The behaviour of oxidizing agents like dichromate differs greatly depending on the pH of the solution. In acidic conditions, dichromate ions have a high oxidizing power and facilitate various chemical transformations. The acidity provides additional protons (\(\mathrm{H}^+\) ions), which can work with the oxidizing agent to facilitate the transfer of electrons.

For example, in an acidic solution, the dichromate ion can oxidize iodine from its elemental state to \(\mathrm{IO}_{3}^{-}\) with an oxidation state of +5. The \(\mathrm{H}^+\) ions from the acid donate protons that combine with the electrons from the iodine atoms, thereby making the reaction possible. Understanding the role of acidity can be crucial for predicting and explaining the outcomes of redox reactions involving oxidizing agents like dichromate ions.