Question

Whenever a reaction between an oxidising agent and a reducing agent is carried out, a compound of lower oxidation state is formed if the reducing agent is in excess and a compound of higher oxidation state is formed if the oxidising agent is in excess. Justify this statement giving three illustrations.

Short answer

Excess reactant affects products: reducing agent forms lower oxidation state, oxidizing forms higher.

Step-by-step solution

Understanding Oxidizing and Reducing Agents

Oxidizing agents are substances that gain electrons and get reduced, while reducing agents are substances that lose electrons and get oxidized in a chemical reaction. In a reaction where both are involved, the oxidizing agent will decrease in oxidation state, whereas the reducing agent will increase.

Analyzing the First Illustration

Consider the reaction between chlorine (\(\text{Cl}_2\)) as an oxidizing agent and sodium (\(\text{Na}\)) as a reducing agent. If sodium is in excess, chlorine is reduced from 0 to -1, forming sodium chloride, which has the chlorine in a lower oxidation state (Cl^-).

Analyzing the Second Illustration

Now, consider a reaction involving potassium permanganate (KMnO_4) as an oxidizing agent and oxalic acid (C_2H_2O_4) as the reducing agent. If oxalic acid is in excess, MnO_4^- reduces to Mn^{2+}, decreasing its oxidation state from +7 to +2, forming manganese(II) ions.

Analyzing the Third Illustration

In the reaction between hydrogen peroxide (H_2O_2) and iodide ions (I^-), H_2O_2 acts as the oxidizing agent and iodide as the reducing agent. If H_2O_2 is in excess, iodide ions are oxidized to iodine (I_2), increasing their oxidation state from -1 to 0, while H_2O_2 forms water.

Key concepts

Oxidizing Agents

In redox reactions, oxidizing agents are essential players that help drive the chemical processes. They are substances that gain electrons in a reaction and, through this process, become reduced. This might seem a bit counterintuitive because, despite being called oxidizing agents, they themselves undergo reduction. This is because they facilitate the oxidation of another substance by accepting electrons from it.
For example, when chlorine (\( \text{Cl}_2 \)) acts as an oxidizing agent in the reaction with sodium (\( \text{Na} \)), it accepts electrons from sodium atoms. Here, each chlorine molecule gains two electrons to form two chloride ions (\( \text{Cl}^- \)), thereby reducing its oxidation state from 0 to -1.

• Oxidizing agents receive electrons and get reduced.
• They cause other substances to lose electrons.
• Help in increasing the oxidation state of other reactants.

Understanding how oxidizing agents work is vital as they play a crucial role in many chemical and biological processes, ranging from industry applications to our own metabolism.

Reducing Agents

Reducing agents are the counterparts to oxidizing agents in redox reactions. While oxidizing agents gain electrons, reducing agents lose electrons, thereby undergoing oxidation themselves. A good way to remember this is that reducing agents "reduce" their own electron count. They donate electrons to the oxidizing agents, which results in the reduction of the oxidizing agents.
In the reaction between potassium permanganate (\( \text{KMnO}_4 \)) and oxalic acid (\( \text{C}_2\text{H}_2\text{O}_4 \)), oxalic acid serves as the reducing agent. As it donates electrons to potassium permanganate, its own oxidation state increases, while the oxidation state of manganese in \( \text{KMnO}_4 \) decreases from +7 to +2, forming manganese(II) ions.

• Reducing agents donate electrons and undergo oxidation.
• They cause other substances to gain electrons.
• Assist in decreasing the oxidation state of oxidizing substances.

This ability to shift electrons makes reducing agents vital in numerous applications, from industrial processes to biological systems.

Oxidation State

The oxidation state is a key concept that helps to keep track of electron transfer in redox reactions. It is a number that indicates the degree of oxidation or reduction of an atom in a compound. You can think of it as a bookkeeping tool to see how many electrons have been gained or lost. The general rule of thumb is:

• An increase in oxidation state means oxidation.
• A decrease in oxidation state indicates reduction.

For example, in the reaction involving hydrogen peroxide (\( \text{H}_2\text{O}_2 \)) and iodide ions (\( \text{I}^- \)), hydrogen peroxide acts as the oxidizing agent. It helps oxidize iodide ions, increasing their oxidation state from -1 to 0, turning them into iodine (\( \text{I}_2 \)), while it reduces to form water.
This change in oxidation states maintains a balance in the chemical equation and in the transfer of electrons, which is crucial for understanding how reactions proceed and predicting their outcomes. Knowing the oxidation state allows chemists to predict \( \text{and} \) balance complex redox equations accurately.