Problem 60 - Chemistry and Chemical Reactivity

Question

To prepare chlorine from chloride ion a strong oxidizing agent is required. The dichromate ion, \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-},\) is one example (Figure 21.32 ). Consult the table of standard reduction potentials (Appendix \(\mathbf{M}\) ), and identify several other oxidizing agents that may be suitable. Write balanced equations for the reactions of these substances with chloride ion.

Step-by-Step Solution

Answer

\( \mathrm{MnO}_{4}^{-} \), \( \mathrm{O}_{2} \), and \( \mathrm{Ag}^{+} \) can oxidize chloride to chlorine, with balanced equations provided.

Step 1: Identify Oxidizing Agents

First, consult the table of standard reduction potentials. Look for species with a higher reduction potential than chlorine, as they can oxidize chloride ions to chlorine gas.

Step 2: List Potential Oxidizing Agents

Examples from the table include: \( \mathrm{MnO}_{4}^{-} \), \( \mathrm{O}_{2} \), and \( \mathrm{Ag}^{+} \). These have high standard reduction potentials: \( +1.51 \text{ V} \) for \( \mathrm{MnO}_{4}^{-} \), \( +1.23 \text{ V} \) for \( \mathrm{O}_{2} \), and \( +0.80 \text{ V} \) for \( \mathrm{Ag}^{+} \).

Step 3: Write Balanced Equations

Using the identified agents, write balanced reactions for each:1. **\( \mathrm{MnO}_{4}^{-} \) reaction:** \[ \mathrm{MnO}_{4}^{-} + 8\mathrm{H}^{+} + 5\mathrm{Cl}^{-} \rightarrow \mathrm{Mn}^{2+} + 4\mathrm{H}_{2}\mathrm{O} + \frac{5}{2}\mathrm{Cl}_{2} \]2. **\( \mathrm{O}_{2} \) reaction:** \[ \mathrm{O}_{2} + 4\mathrm{H}^{+} + 4\mathrm{Cl}^{-} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + 2\mathrm{Cl}_{2} \]3. **\( \mathrm{Ag}^{+} \) reaction:** \[ 2\mathrm{Ag}^{+} + 2\mathrm{Cl}^{-} \rightarrow 2\mathrm{Ag} + \mathrm{Cl}_{2} \]

Step 4: Verify Reaction Conditions

Ensure that reaction conditions (e.g., acidic or neutral medium) match those in which these agents are effective. For some reactions, such as with \( \mathrm{MnO}_{4}^{-} \) or \( \mathrm{O}_{2} \), an acidic medium is necessary.

Key Concepts

Chloride Ion

Chloride ions (\( \mathrm{Cl}^{-} \)) are derived from chlorine, an essential element in chemistry with a variety of uses. Chloride ions are formed when chlorine atoms gain an electron. This process stabilizes the chlorine into a negatively charged ion, commonly found in various salts like sodium chloride (table salt).
These ions are pivotal in many chemical reactions, especially those involving oxidation and reduction (redox) processes. Chloride ions are not only common in everyday life but also play critical roles in industrial applications, such as water treatment and the production of disinfectants.
A fascinating aspect of chloride ions is their role as a reducing agent, meaning they can donate electrons to other substances. In reactions where we aim to produce chlorine gas from chloride ions, the chloride ions lose electrons, thus getting oxidized themselves. This transformation requires a strong oxidizing agent, which can seize the extra electrons from the chloride ions to facilitate the change from \( \mathrm{Cl}^{-} \) back to \( \mathrm{Cl}_{2} \) gas.

Standard Reduction Potentials

Understanding standard reduction potentials is crucial in determining which substances can act as oxidizing agents. Standard reduction potentials, measured in volts (V), indicate a substance's readiness to gain electrons and, thus, its capability as an oxidizing agent.
These potentials are listed in tables order by their strength, with higher values representing a greater ability to gain electrons. For example, when looking to oxidize chloride ions into chlorine gas, substances with a higher standard reduction potential than chlorine itself are required.
The values for some strong oxidizing agents include:

• \( \mathrm{MnO}_{4}^{-} \) with \( +1.51 \text{ V} \)
• \( \mathrm{O}_{2} \) with \( +1.23 \text{ V} \)
• \( \mathrm{Ag}^{+} \) with \( +0.80 \text{ V} \)

By consulting such tables, chemists can choose suitable oxidizing agents to effectively carry out desired reactions. Selecting an agent with a higher potential ensures a more spontaneous and efficient reaction, specifically in oxidizing chloride ions to chlorine gas.

Balanced Equations

In chemical reactions, balanced equations represent a fundamental principle: the conservation of mass. Each side of the reaction must have an equal amount of each element and charge. This ensures that no atoms are lost or gained in the process.
Let's look at the reactions used to oxidize chloride ions with different agents.

• Firstly, consider the reaction with \( \mathrm{MnO}_{4}^{-} \): \[ \mathrm{MnO}_{4}^{-} + 8\mathrm{H}^{+} + 5\mathrm{Cl}^{-} \rightarrow \mathrm{Mn}^{2+} + 4\mathrm{H}_{2}\mathrm{O} + \frac{5}{2}\mathrm{Cl}_{2} \]. Here, the reaction balances both the atoms and the charges perfectly, considering it occurs in an acidic medium.


• Secondly, check the reaction involving oxygen, \( \mathrm{O}_{2} \): \[ \mathrm{O}_{2} + 4\mathrm{H}^{+} + 4\mathrm{Cl}^{-} \rightarrow 2\mathrm{H}_{2}\mathrm{O} + 2\mathrm{Cl}_{2} \]. Again, both reactants and products are balanced in terms of atoms and charge.


• Lastly, the reaction with silver ions, \( \mathrm{Ag}^{+} \): \[ 2\mathrm{Ag}^{+} + 2\mathrm{Cl}^{-} \rightarrow 2\mathrm{Ag} + \mathrm{Cl}_{2} \]. In this reaction, the number of atoms and charges on each side align precisely.

Balancing chemical equations is essential for accurately representing reactions. It provides clarity about the stoichiometry, ensuring that calculations for reactants or products are correct and maintain the reaction's integrity.