Problem 75 - Chemistry and Chemical Reactivity

Question

Bromine is obtained from sea water by the following redox reaction: $$\mathrm{Cl}_{2}(\mathrm{g})+2 \mathrm{NaBr}(\mathrm{aq}) \rightarrow 2 \mathrm{NaCl}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)$$ (a) What has been oxidized? What has been reduced? (b) Identify the oxidizing and reducing agents.

Step-by-Step Solution

Answer

Bromine is oxidized and chlorine is reduced; \( \text{Cl}_2 \) is the oxidizing agent, and \( \text{Br}^- \) is the reducing agent.

Step 1: Determine Oxidation States

To find what is oxidized and reduced, determine the oxidation states of Br and Cl before and after the reaction. Chlorine starts in \( \text{Cl}_2 \) with an oxidation state of 0 and ends in \( \text{NaCl} \) with an oxidation state of -1. Bromine starts in \( \text{NaBr} \) with an oxidation state of -1 and ends in \( \text{Br}_2 \) with an oxidation state of 0.

Step 2: Identify Changes in Oxidation States

Reduction is a decrease in oxidation state and oxidation is an increase. Chlorine goes from 0 to -1, indicating a reduction. Bromine goes from -1 to 0, indicating an oxidation.

Step 3: Determine Oxidizing and Reducing Agents

The substance being reduced acts as the oxidizing agent, and the substance being oxidized acts as the reducing agent. Since \( \text{Cl}_2 \) is reduced, it is the oxidizing agent. Since \( \text{Br}^- \) (from \( \text{NaBr} \)) is oxidized, it is the reducing agent.

Key Concepts

Oxidation States

An oxidation state, often called an oxidation number, represents the degree of oxidation or reduction of an atom in a chemical compound. By identifying oxidation states, we can determine how electrons are distributed in a reaction. This helps us to understand which elements lose or gain electrons, thereby identifying what has been oxidized and reduced.

• In the reaction \( \text{Cl}_2(\text{g}) + 2 \text{NaBr}(\text{aq}) \rightarrow 2 \text{NaCl}(\text{aq}) + \text{Br}_2(\ell) \), chlorine begins with an oxidation state of 0 and ends with -1.
• Bromine goes from an oxidation state of -1 in \( \text{NaBr} \) to 0 in \( \text{Br}_2 \).

The change for chlorine shows a reduction since it gains electrons, whereas bromine's increase indicates it is oxidized, as it loses electrons. Understanding these shifts in oxidation states allows us to identify the key components in redox reactions.

Oxidizing Agents

The oxidizing agent in a redox reaction is the substance that gains electrons. It causes another substance to lose electrons, thereby undergoing reduction itself.
In our example reaction, chlorine is the oxidizing agent.

• Chlorine goes from an oxidation state of 0 to -1, meaning it gains electrons.
• As it gains electrons from bromine (Br in \( \text{NaBr} \)), it facilitates bromine's oxidation.

By accepting electrons, the oxidizing agent plays a crucial role in driving the redox process forward.

Reducing Agents

A reducing agent is responsible for donating electrons to another substance, causing that substance to be reduced. During this process, the reducing agent itself is oxidized.

• In the given reaction, bromine in \( \text{NaBr} \) acts as the reducing agent.
• It begins with an oxidation state of -1 and increases to 0, indicating it loses electrons to chlorine.

As a reducing agent, bromine's electron donation is essential for chlorine's reduction. Recognizing the reducing agent helps us understand how electron transfers shape chemical reactions.