Problem 57 - A Complete Resource Book in Chemistry for JEE Main

Question

In the reaction, \(3 \mathrm{Br}_{2}+6 \mathrm{CO}_{3}^{2-}+3 \mathrm{H}_{2} \mathrm{O} \longrightarrow 5 \mathrm{Br}^{-}+\mathrm{BrO}_{3}^{-}+6 \mathrm{HCO}_{3}^{-}\) (a) bromine is oxidized and carbonate is reduced (b) bromine is reduced and water is oxidized (c) bromine is neither reduced nor oxidized (d) bromine is both reduced and oxidized

Step-by-Step Solution

Answer

Bromine is both reduced and oxidized in the reaction, so (d) is correct.

Step 1: Identify Changes in Oxidation States

To solve this, first identify the oxidation states of the elements involved in the reaction. Start by analyzing bromine (Br). In \( \text{Br}_2 \), each bromine atom has an oxidation state of 0, while in \( \text{Br}^- \), it is -1, and in \( \text{BrO}_3^- \), it is +5. This suggests a change in oxidation state.

Step 2: Determine Reduction and Oxidation Processes

Oxidation refers to an increase in oxidation state, while reduction refers to a decrease. In this reaction, bromine in \( \text{Br}_2 \) changes in different ways: it decreases from 0 to -1 (reduction) in \( \text{Br}^- \) and increases from 0 to +5 (oxidation) in \( \text{BrO}_3^- \). This confirms that bromine is both oxidized and reduced in this reaction.

Step 3: Assessment of Carbonate and Water

The carbonate ion (\( \text{CO}_3^{2-} \)) is a reactant and turns into \( \text{HCO}_3^- \). Its oxidation state for carbon remains unchanged, so no reduction/oxidation happens here. Water also remains chemically unchanged in terms of oxidation states.

Step 4: Choose the Correct Option from Multiple Choice

Based on the changes identified, option (d), which states that 'bromine is both reduced and oxidized,' is confirmed by the oxidation and reduction of bromine in different parts of the reaction.

Key Concepts

Oxidation States

Understanding oxidation states is crucial for solving redox reactions. An oxidation state indicates the degree of oxidation of an atom in a compound. It helps to track how electrons are transferred during the reaction.
In the reaction provided, the oxidation state of bromine changes depending on the species it is in. For instance:

• In elemental bromine, \( \text{Br}_2 \), the oxidation state of each bromine atom is 0 because they are in pure elemental form.
• In bromide ion, \( \text{Br}^- \), the oxidation state of bromine is -1, representing a gain of an electron.
• In bromate ion, \( \text{BrO}_3^- \), the oxidation state of bromine is +5, indicating a loss of electrons.

These changes demonstrate how bromine undergoes both reduction and oxidation, as it engages in electron exchange, which is the essence of redox reactions.

Bromine

Bromine is a chemical element with symbol \( \text{Br} \) and atomic number 35. It plays an active role in many redox reactions.
In chemistry, bromine is notable for its ability to change oxidation states, which makes it a versatile participant in these processes. In the given redox reaction:

• Bromine starts in a neutral state, \( \text{Br}_2 \), where each atom holds an oxidation state of 0.
• Upon reduction, it becomes \( \text{Br}^- \), demonstrating a decrease in oxidation state due to electron gain.
• Simultaneously, it can undergo oxidation, forming \( \text{BrO}_3^- \), which requires an increase in oxidation state through electron loss.

This dual behavior of bromine—being both reduced and oxidized in the same reaction—illustrates its dynamic nature in redox chemistry. It's essential to identify the various compounds bromine forms to understand its role thoroughly.

Reduction and Oxidation Processes

Redox reactions are characterized by the transfer of electrons, categorized into reduction and oxidation processes. To conceptualize these processes:

• Oxidation: Involves the increase in oxidation state, meaning loss of electrons. In our reaction example, bromine being oxidized to \( \text{BrO}_3^- \) represents oxidation.
• Reduction: Involves the decrease in oxidation state, meaning gain of electrons. Here, the transformation of bromine into \( \text{Br}^- \) signifies reduction.

In any redox process, these two occur simultaneously, where one species is oxidized, losing electrons, and another is reduced, gaining those electrons. The reaction's complexity can vary, with substances like bromine displaying both phenomena simultaneously, a process known as "disproportionation," where a single substance undergoes both oxidation and reduction.