Question

Hydrochloric acid is not an oxidizing agent in the sense that sulfuric acid and nitric acid are. Explain why the chloride ion is not a strong oxidizing agent like ${\mathrm{SO}}_4^{2-}$ and ${\mathrm{NO}}_3^{-}$.

Short answer

Chloride ions are stable and do not readily gain electrons, unlike ${\mathrm{SO}}_4^{2-}$ and ${\mathrm{NO}}_3^{-}$, which have higher reduction potentials.

Step-by-step solution

Understand Oxidizing Agents

An oxidizing agent is a substance that gains electrons in a chemical reaction, thereby oxidizing another species. The strength of an oxidizing agent is determined by its ability to accept electrons.

Review Electron Configurations

Oxygen has a higher electronegativity and electron affinity compared to chlorine, which means oxygen-containing ions, like sulfate (${\mathrm{SO}}_4^{2-}$) and nitrate (${\mathrm{NO}}_3^{-}$), can stabilize extra electrons better and act as stronger oxidizing agents.

Analyze Chloride Ion

The chloride ion (${\mathrm{Cl}}^{-}$) has a stable electron configuration because it achieves a noble gas configuration by gaining just one electron. This makes it less eager to gain additional electrons and thus not a strong oxidizing agent.

Compare Reduction Potentials

Reduction potential measures the tendency of a species to gain electrons. Both ${\mathrm{SO}}_4^{2-}$ and ${\mathrm{NO}}_3^{-}$ have higher reduction potentials compared to ${\mathrm{Cl}}^{-}$, meaning they are more likely to gain electrons and oxidize other substances.

Key concepts

Electron Configuration

Electron configuration helps us understand how electrons are arranged around an atom. It is crucial for determining the chemical properties of an element, including its reactivity and tendency to participate in redox reactions.
For any atom, electrons exist in different principal energy levels or shells, which are further divided into subshells (s, p, d, f). The arrangement follows a specific order based on increasing energy levels, governed by the principles of quantum mechanics.
The chloride ion (${\text{Cl}}^{-}$), in particular, achieves a stable electron configuration by simply gaining one more electron, reaching the configuration of the noble gas, argon. This means it has filled its outer electron shell, making it more stable and less likely to gain additional electrons.

• In its neutral state, chlorine (Cl) has the electron configuration of $$1s^{2}2s^{2}2p^{6}3s^{2}3p^5$$.
• When it becomes a chloride ion, it gains one electron: $$1s^{2}2s^{2}2p^{6}3s^{2}3p^6$$.

This electron configuration explains why ${\text{Cl}}^{-}$ is a poor oxidizing agent: it's already stabilized.

Reduction Potentials

Reduction potential is a key concept in electrochemistry, measuring the tendency of a chemical species to gain electrons in a redox reaction. A higher reduction potential means a stronger tendency to be reduced, serving as a more potent oxidizing agent.
The standard reduction potential is usually given in volts (V) and is denoted by the symbol $E^0$.
We reference this to the standard hydrogen electrode (SHE), which is set at 0 V. Therefore, a substance with a more positive $E^0$ value is more likely to gain electrons and reduce other species.

• For sulfate ion (${\text{SO}}_4^{2-}$), it tends to have high reduction potential in comparison to the chloride ion.
• Similarly, nitrate ion (${\text{NO}}_3^{-}$) is also known for its strong oxidizing abilities due to its significant reduction potential.

These high reduction potentials signify these ions can readily gain electrons, unlike chloride ions where the reduction potential is much lower, making them less viable as oxidizing agents.

Chloride Ion

The chloride ion, represented as ${\text{Cl}}^{-}$, plays a vital role in distinguishing strong and weak oxidizing agents.
To dive deeper into its nature, let's explore why it isn't a strong oxidizing agent.
As we have seen, it achieves a stable electron configuration similar to the noble gas argon upon gaining one electron. This established stability minimizes its desire to gain further electrons unlike its oxygen-based counterparts like sulfate and nitrate, which have a robust capability to stabilize additional electrons.

• ${\text{Cl}}^{-}$ is more involved in forming ionic bonds, like in NaCl, rather than participating in redox reactions.
• Its participation as an oxidant is notably low in biological and chemical systems.

Hence, the chloride ion is recognized more as a spectator ion than an active participant in redox chemistry.