Which of the following is true about the given redox reaction? $$ \mathrm{SnCl}_{2}+2 \mathrm{FeCl}_{3} \rightarrow \mathrm{SnCl}_{4}+2 \mathrm{FeCl}_{2} $$ (a) \(\mathrm{SnCl}_{2}\) is oxidised and \(\mathrm{FeCl}_{3}\) acts as oxidising agent. (b) \(\mathrm{FeCl}_{3}\) is oxidised and acts as oxidising agent. (c) \(\mathrm{SnCl}_{2}\) is reduced and acts as oxidising agent. (d) \(\mathrm{FeCl}_{3}\) is oxidised and \(\mathrm{SnCl}_{2}\) acts as a oxidising agent.

Understanding oxidation states is crucial in the analysis of redox reactions, which involve the transfer of electrons between substances. Oxidation states, sometimes referred to as oxidation numbers, are a bookkeeping system that helps to keep track of the electrons in a compound or chemical species. They represent the charge that an atom would have if all bonds to atoms of different elements were 100% ionic.

When identifying oxidation states, some rules are widely accepted: elements in their elemental form have an oxidation state of zero, and the sum of oxidation states for a neutral molecule must be zero. For ions, the sum must equal the charge of the ion. For example, in a compound like \(\mathrm{SnCl}_{2}\), tin (Sn) assumes an oxidation state that balances out the negative charges of the chloride ions.

In the reaction presented, determining the oxidation states of metal atoms allows us to see the electron transfer process: tin (Sn) increases from +2 to +4, signifying a loss of electrons, and iron (Fe) decreases from +3 to +2, indicating a gain of electrons. Being able to calculate these values is essential for the next steps of analyzing the redox behavior of substances involved in chemical reactions.

An oxidising agent, or oxidant, plays a pivotal role in redox reactions. It is the chemical that takes electrons away from another substance, causing that substance to be oxidised. This means the oxidising agent itself undergoes reduction, which is the gain of electrons, resulting in a decrease in its oxidation state.

In our example reaction, \(\mathrm{FeCl}_{3}\) is an oxidising agent because it is responsible for oxidising \(\mathrm{SnCl}_{2}\) by accepting electrons from the tin (Sn). This process transforms iron (Fe) from a +3 oxidation state to +2. Therefore, \(\mathrm{FeCl}_{3}\) decreases in oxidation state, showcasing its role as the oxidising agent.

An effective way to remember this concept is to think of an oxidising agent as facilitating oxidation in another species and thus, 'gets reduced' in the process. Therefore, agents that have a relatively high affinity for electrons or have high oxidation states to begin with, often act as oxidising agents in chemical reactions.

Conversely, a reducing agent, or reductant, is the substance in a redox reaction that donates electrons to another chemical species, causing itself to be oxidised. The reducing agent sees an increase in its oxidation state as it loses electrons during the reaction.

For the example of the reaction between \(\mathrm{SnCl}_{2}\) and \(\mathrm{FeCl}_{3}\), \(\mathrm{SnCl}_{2}\) serves as the reducing agent because it surrenders electrons to \(\mathrm{FeCl}_{3}\), which is reduced. The oxidation state of tin (Sn) in \(\mathrm{SnCl}_{2}\) goes from +2 to +4, which confirms that it loses electrons and is therefore the reducing agent in the reaction.

Remembering that a reducing agent 'reduces another species' while 'being oxidised itself' can be helpful. Substances that are effective as reducing agents typically have a low affinity for electrons, or potentially have additional electrons that can be readily donated to other substances in a chemical reaction.