Question

What happens when an alkaline solution of potassium ferricyanide is reacted with \(\mathrm{H}_{2} \mathrm{O}_{2} ?\) (a) Potassium ferricyanide is oxidised to potassium ferrocyanide and \(\mathrm{H}_{2} \mathrm{O}_{2}\) is oxidised. (b) Potassium ferricyanide becomes colourless and \(\mathrm{H}_{2} \mathrm{O}_{2}\) is oxidised to \(\mathrm{O}_{2}\). (c) Potassium ferricyanide is reduced to ferric hydroxide and \(\mathrm{H}_{2} \mathrm{O}_{2}\) is oxidised to \(\mathrm{H}_{2} \mathrm{O}\) (d) Potassium ferricyanide is reduced to potassium ferrocyanide and \(\mathrm{H}_{2} \mathrm{O}_{2}\) is oxidised to \(\mathrm{O}_{2}\).

Short answer

When an alkaline solution of potassium ferricyanide reacts with H2O2, the potassium ferricyanide is reduced to potassium ferrocyanide, and H2O2 is oxidized to O2. The correct answer is (d).

Step-by-step solution

Understanding the Reaction

To solve this exercise, we need to understand what happens when an alkaline solution containing potassium ferricyanide (K3[Fe(CN)6]) reacts with hydrogen peroxide (H2O2). During this reaction, the ferricyanide ion [Fe(CN)6]^{3-} is expected to change its oxidation state. Hydrogen peroxide can act as an oxidizing or a reducing agent under different conditions, but in alkaline solutions, it commonly acts as an oxidizing agent.

Determining the Oxidation States

The oxidation state of iron in potassium ferricyanide is +3. When it reacts with an oxidizing agent like H2O2, it will get reduced to a lower oxidation state, which is +2 in potassium ferrocyanide (K4[Fe(CN)6]). Potassium ferricyanide itself cannot be further oxidized as it is already in a high oxidation state. Thus, option (a) is incorrect.

Analyzing the Products

After the reaction, we can eliminate the possibilities where potassium ferricyanide is said to be oxidized or produces a product that does not involve a change in the oxidation state of iron, such as ferric hydroxide. The only feasible option is that potassium ferricyanide (an oxidizing agent) is reduced to potassium ferrocyanide, while the H2O2 (acting as a reducing agent in this reaction) is oxidized to O2. So, option (d) describes what would happen adequately.

Key concepts

Understanding Oxidation States

Oxidation states, also known as oxidation numbers, are a fundamental concept in chemistry that denote the degree of oxidation of an atom within a compound. They are essential for identifying how electrons are transferred between atoms in a chemical reaction. This transfer of electrons is what characterizes redox reactions.

Oxidation occurs when an atom loses electrons, increasing its oxidation state, while reduction involves an atom gaining electrons, decreasing its oxidation state. The summation of oxidation states for all atoms in a neutral molecule must equal zero, whereas for an ion, it must equal the ionic charge.

In our exercise, the potassium ferricyanide \((K_3[Fe(CN)_6])\) contains iron with an oxidation state of +3. As the exercise reveals, when this complex reacts with hydrogen peroxide in an alkaline solution, the oxidation state of iron is reduced to +2, as found in potassium ferrocycanide \((K_4[Fe(CN)_6])\). Understanding these state changes is crucial for interpreting the outcome of the reaction correctly.

The Nature of Redox Reactions in Chemistry

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two substances. They are a backbone of many chemical processes, from industrial syntheses to biological metabolism. For a redox reaction to occur, there must be at least two species - one that gets oxidized, losing electrons, and one that gets reduced, gaining electrons.

In the context of the exercise, we have a redox reaction where potassium ferricyanide and hydrogen peroxide react together. Through this reaction, we observe the change in the oxidation state of iron from +3 in ferricyanide to +2 in ferrocyanide, indicating a gain of electrons, hence reduction. At the same time, hydrogen peroxide acts as the reducing agent, being itself oxidized to dioxygen (\(O_2\)).

To solve problems involving redox reactions, it is essential to identify the substances getting oxidized and reduced. This facilitates the correct prediction of reaction products and an understanding of the overall chemical process.

Hydrogen Peroxide Reactions

Hydrogen peroxide \((H_2O_2)\) is known for its versatility as a chemical, being able to act as either an oxidizing agent or a reducing agent depending on the reaction environment. When reacting in an alkaline solution, \((H_2O_2)\) often acts as an oxidizing agent as seen in the exercise with potassium ferricyanide.

This reaction's pathway highlights the ability of \((H_2O_2)\) to undergo disproportionation – a type of redox reaction in which a substance is simultaneously oxidized and reduced, producing two different products. In the reaction we are discussing, hydrogen peroxide is oxidized to form oxygen gas \((O_2)\) as one of the products, showcasing its role as an oxidant.

Understanding the behavior of \((H_2O_2)\) in different environments is pivotal for accurately predicting its reactions. Its concentration and the pH of the solution are key factors that can influence whether \((H_2O_2)\) will act as an oxidizing or reducing agent, which makes it an interesting subject of study in chemistry, especially in redox chemistry.