Question

Reduction of the metal centre in aqueous permanganate ion involves (A) 3 electrons in neutral medium (B) 5 electrons in neutral medium (C) 3 electrons in alkaline medium (D) 5 electrons in acidic medium

Short answer

The reduction of the metal center in aqueous permanganate ion involves 3 electrons in neutral medium (A) and 5 electrons in acidic medium (D).

Step-by-step solution

Identify the Chemical Reaction for Neutral Medium

The reduction of permanganate in a neutral medium can be represented by the half-reaction: \( \mathrm{MnO_4^-} + 2H_2O + 3e^- \rightarrow \mathrm{MnO_2} + 4OH^- \). In this reaction, 3 electrons are involved in reducing the permanganate ion to manganese dioxide (MnO2).

Identify the Chemical Reaction for Acidic Medium

The reduction of permanganate in an acidic medium is different and is represented by the half-reaction: \( \mathrm{MnO_4^-} + 8H^+ + 5e^- \rightarrow \mathrm{Mn^{2+}} + 4H_2O \). Here, 5 electrons are required to reduce the permanganate ion to Mn(2+).

Analyze the Options Provided

Based on the above half-reactions, we can conclude that option (A) corresponds to the number of electrons transferred in the neutral medium, and option (D) corresponds to the electron transfer in the acidic medium. Options (B) and (C) are incorrect for the given media.

Choose the Correct Answer

Since the exercise asks for the reduction in neutral and acidic media, we choose (A) for the neutral medium, which involves 3 electrons, and (D) for the acidic medium, which involves 5 electrons.

Key concepts

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are chemical processes where electrons are transferred between atoms, ions, or molecules. This transfer results in changes in the oxidation states of the reactants involved. In every redox reaction, two half-reactions occur simultaneously: one substance is oxidized by losing electrons, while another is reduced by gaining electrons.

An easy-to-understand example is the rusting of iron, where iron (Fe) is oxidized to form rust (iron oxide, Fe2O3), and oxygen is reduced. In the case of the permanganate ion (MnO_4^-), it undergoes reduction, meaning it gains electrons. In a neutral medium, it takes up 3 electrons and is reduced to manganese dioxide (MnO_2), while in an acidic medium it gains 5 electrons to form Mn(2+) ions.

Through redox reactions, we observe the fundamental principle of conservation of charge, as the gain of electrons (reduction) by one species is balanced by the loss of electrons (oxidation) by another. This unity between oxidation and reduction is central to understanding chemical reactivity and is widely applied in fields like energy storage (batteries), metallurgy, and biochemistry.

Half-reaction Method

The half-reaction method is a systematic approach to balancing redox reactions. This technique divides the overall chemical equation into two separate half-reactions: one for oxidation and one for reduction. Each half-reaction is then balanced individually, making it easier to follow the electron transfer process. This approach is incredibly useful when dealing with complex reactions, such as those involving permanganate ion reduction.

H4: Break Down Steps
To apply the half-reaction method, follow these general steps:

• Separate the overall redox reaction into two half-reactions.
• Balance all elements in the half-reactions, except for hydrogen and oxygen.
• Balance oxygen atoms by adding water molecules.
• Balance hydrogen atoms by adding hydrogen ions or hydroxide ions, depending on the reaction's pH.
• Balance the charge by adding electrons to the more positive side of each half-reaction.
• Equalize the number of electrons in both half-reactions, if necessary, by multiplying by appropriate factors.
• Combine the two half-reactions to form the complete balanced redox equation.

This method not only helps in balancing complicated redox reactions but also offers insight into the transfer of electrons, which is a cornerstone of understanding redox chemistry.

Oxidation States

Understanding oxidation states is crucial for interpreting redox reactions. An oxidation state, often referred to as oxidation number, is a theoretical charge that an atom would have if all bonds to atoms of different elements were completely ionic. It is a tool for keeping track of electron transfer during redox reactions.

H4: Tracking Electrons
Oxidation state changes indicate whether an atom is oxidized (loss of electrons) or reduced (gain of electrons). For instance, the permanganate ion (MnO_4^-) reduction involves changes in the oxidation state of manganese. In a neutral medium, the oxidation state of manganese changes from +7 in MnO_4^- to +4 in MnO_2. In an acidic medium, it is further reduced to a +2 oxidation state as manganese becomes Mn(2+).

Some rules for assigning oxidation states include:

• The oxidation state of an atom in its elemental form is 0.
• The sum of oxidation states for all atoms in a neutral compound must be zero.
• In ions, the sum of oxidation states must equal the overall charge of the ion.
• Group 1 metals have an oxidation state of +1, and group 2 metals have +2, in their compounds.

Oxidation states serve not only as a bookkeeping tool but also as a way to predict the reactivity and properties of chemical species, and they are essential for solving redox equations using the half-reaction method.