Question

(a) Assuming standard conditions, arrange the following in order of increasing strength as oxidizing agents in acidic solution: \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Cu}^{2+}, \mathrm{Cl}_{2}, \mathrm{O}_{2} .\) (b) Arrange the following in order of increasing strength as reducing agents in acidic solution: \(\text {Zn,}\) \(\mathrm{I}^{-}, \mathrm{Sn}^{2+}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Al} .\)

Short answer

(a) The oxidizing agents in order of increasing strength in acidic solution are: \(\mathrm{Cu}^{2+}, \mathrm{O}_{2}, \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}, \mathrm{Cl}_{2},\) and \(\mathrm{H}_{2} \mathrm{O}_{2}.\) (b) The reducing agents in order of increasing strength in acidic solution are: \(\mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{I}^{-}, \mathrm{Sn}^{2+}, \text {Zn},\) and \(\mathrm{Al}.\)

Step-by-step solution

Find the standard reduction potentials for the given half-cell reactions.

First, we need to find the standard reduction potentials (E°) for the given half-cell reactions. We can do this by consulting a table of standard reduction potentials. (a) For the oxidizing agents in acidic solution: - \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+ 14\mathrm{H}^{+}+ 6\mathrm{e}^{-} \longrightarrow 2\mathrm{Cr}^{3+} + 7H_{2}O\), E° = 1.33 V - \(\mathrm{H}_{2} \mathrm{O}_{2} + 2\mathrm{H}^{+} + 2\mathrm{e}^{-} \longrightarrow 2\mathrm{H}_{2}\mathrm{O}\), E° = 1.78 V - \(\mathrm{Cu}^{2+} + 2\mathrm{e}^{-} \longrightarrow \mathrm{Cu}\), E° = 0.34 V - \(\mathrm{Cl}_{2} + 2\mathrm{e}^{-} \longrightarrow 2\mathrm{Cl}^-\), E° = 1.36 V - \(\mathrm{O}_{2} + 4\mathrm{H}^{+} + 4\mathrm{e}^{-} \longrightarrow 2\mathrm{H}_{2}\mathrm{O}\), E° = 1.23 V (b) For the reducing agents in acidic solution: - \(\text {Zn} + 2\mathrm{H}^{+} \longrightarrow \mathrm{Zn}^{2+}+ \mathrm{H}_{2}\), E° = -0.76 V - \(2\mathrm{I}^{-} \longrightarrow \mathrm{I}_{2}+ 2\mathrm{e}^{-}\), E° = 0.54 V - \(\mathrm{Sn}^{2+} + 2\mathrm{e}^{-} \longrightarrow \mathrm{Sn}\), E° = -0.14 V - \(\mathrm{H}_{2} \mathrm{O}_{2} + 2\mathrm{H}^{+} + 2\mathrm{e}^{-} \longrightarrow 2\mathrm{H}_{2}\mathrm{O}\), E° = 1.78 V - \(\mathrm{Al} + 3\mathrm{H}^{+} \longrightarrow \mathrm{Al}^{3+} + \frac{3}{2}\mathrm{H}_{2}\), E° = -1.66 V

Arrange the substances in order of increasing strength as oxidizing agents in acidic solution.

Now that we have the standard reduction potentials, we can arrange the given substances in order of increasing strength as oxidizing agents. We're looking for the substances with the most positive E° values, as these will be the strongest oxidizing agents. (a) The oxidizing agents in order of increasing strength are as follows: - \(\mathrm{Cu}^{2+}\) (0.34 V) - \(\mathrm{O}_{2}\) (1.23 V) - \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\) (1.33 V) - \(\mathrm{Cl}_{2}\) (1.36 V) - \(\mathrm{H}_{2} \mathrm{O}_{2}\) (1.78 V)

Arrange the substances in order of increasing strength as reducing agents in acidic solution.

Similarly, for reducing agents, we're looking for the substances with the most negative E° values, as these will be the strongest reducing agents. (b) The reducing agents in order of increasing strength are as follows: - \(\mathrm{H}_{2} \mathrm{O}_{2}\) (1.78 V) - \(\mathrm{I}^{-}\) (0.54 V) - \(\mathrm{Sn}^{2+}\) (-0.14 V) - \(\text {Zn}\) (-0.76 V) - \(\mathrm{Al}\) (-1.66 V)

Key concepts

Oxidizing Agents

In redox (reduction-oxidation) reactions, oxidizing agents play a critical role. They are substances that accept electrons during the chemical reaction. When tasked as an oxidizer, they gain electrons and therefore, undergo reduction. This seems opposite at first, because while they are gaining electrons, they are "causing" the oxidation of another species.

• An important feature of oxidizing agents is their tendency to attract and accept electrons.
• They are typically associated with positive oxidation states and high reduction potential values.
• Examples include elements like chlorine (\( \text{Cl}_2 \)); compounds such as hydrogen peroxide (\( \text{H}_2\text{O}_2 \)); and ions like dichromate (\( \text{Cr}_2\text{O}_7^{2-} \)).

The strength of an oxidizing agent is directly related to its standard reduction potential. The more positive the potential, the stronger the oxidizing agent as it indicates a greater tendency to gain electrons.

Reducing Agents

Reducing agents, on the other side of the redox equation, are substances that donate electrons to another compound. They get oxidized themselves because they lose electrons in the process.

• Reducing agents often possess negative oxidation states and are characterized by negative reduction potentials.
• They include metals, which are generally good electron donors. Familiar examples include zinc (\( \text{Zn} \)), aluminum (\( \text{Al} \)), and iodide ions (\( \text{I}^- \)).

The effectiveness of a reducing agent is determined by its ability to release electrons. Thus, substances with more negative standard reduction potentials are stronger reducing agents as they readily relinquish their electrons in a reaction.

Standard Reduction Potentials

Standard reduction potentials (\( E^{\circ} \)) are used to predict the direction and feasibility of redox reactions. It’s a measure of the tendency of a chemical species to be reduced. These potentials are established under standard conditions: 1 molar concentration, 1 atmosphere pressure, and a temperature of 25°C.

• An electrode potential table is a helpful resource, listing values that aid in comparing the strengths of oxidizers and reducers.
• A positive \( E^{\circ} \) value indicates the half-reaction tends to proceed in the forward direction, behaving as an effective oxidizing agent. Conversely, a negative \( E^{\circ} \) suggests a propensity to release electrons, characteristic of reducing agents.

Utilizing these potentials allows chemists to calculate the cell potential and deduce important details about the electrochemistry involved in a reaction.

Electrochemistry

Electrochemistry is a branch of chemistry that studies the relationship between electricity and chemical reactions. It is pivotal in understanding redox processes since electron transfer is at the core of both chemical changes and electrical movement. Considering how oxidizing agents and reducing agents interact with electrons, electrochemistry connects with various practical applications:

• Electrochemical cells, like batteries, are designed based on redox reactions. They convert chemical energy into electrical energy, relying on the flow of electrons from a reducing agent to an oxidizing agent through an external circuit.
• Electroplating and corrosion are also processes investigated within electrochemistry, utilizing or preventing unwanted redox reactions.

Understanding electrochemistry introduces insights into how electrical energy is stored, transferred, and used, linking chemical reactions with practical technological advancements.