Problem 45 - Chemistry: The Central Science

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: \(\mathrm{Zn}, \mathrm{I}^{-}, \mathrm{Sn}^{2+}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Al}\).

Step-by-Step Solution

Answer

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

Step 1: Find the standard reduction potentials

Refer to a table of standard reduction potentials to find the values for each of the given agents. We'll be looking at their reduction half-reactions: \( \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-} + 14 \mathrm{H}^{+} + 6 \mathrm{e}^- \rightarrow 2 \mathrm{Cr}^{3+} + 7 \mathrm{H}_{2} \mathrm{O}\) : \(E^{\circ} = 1.33 \ V\) \( \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{H}^{+} + 2 \mathrm{e}^- \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) : \(E^{\circ} = 1.77 \ V\) \( \mathrm{Cu}^{2+} + 2 \mathrm{e}^- \rightarrow \mathrm{Cu}\) : \(E^{\circ} = 0.34 \ V\) \( \mathrm{Cl}_{2} + 2 \mathrm{e}^- \rightarrow 2 \mathrm{Cl}^{-}\) : \(E^{\circ} = 1.36 \ V\) \( \mathrm{O}_{2} + 4 \mathrm{H}^{+} + 4 \mathrm{e}^- \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) : \(E^{\circ} = 1.23 \ V\)

Step 2: Arrange in increasing order

Now that we know the standard reduction potentials for the given oxidizing agents, we can arrange them in order of increasing strength. Remember, the higher the reduction potential, the stronger the oxidizing agent: \(\mathrm{Cu}^{2+}, \mathrm{O}_{2}, \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}, \mathrm{Cl}_{2}, \mathrm{H}_{2} \mathrm{O}_{2}\). (b) Arrange the following in order of increasing strength as reducing agents in acidic solution: \(\mathrm{Zn}, \mathrm{I}^{-}, \mathrm{Sn}^{2+}, \mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Al}\).

Step 1: Find the standard reduction potentials

Refer to a table of standard reduction potentials to find the values for each of the given agents. We'll be looking at their oxidation half-reactions: \( \mathrm{Zn} \rightarrow \mathrm{Zn}^{2+} + 2 \mathrm{e}^-\) : \(E^{\circ} = -0.76 \ V\) (in reverse direction) \( \mathrm{I}_{2} + 2 \mathrm{e}^- \rightarrow 2 \mathrm{I}^{-}\) : \(E^{\circ} = -0.54 \ V\) (in reverse direction) \( \mathrm{Sn}^{4+} + 2 \mathrm{e}^- \rightarrow \mathrm{Sn}^{2+}\) : \(E^{\circ} = -0.15 \ V\) (in reverse direction) \( \mathrm{H}_{2} \mathrm{O}_{2} + 2 \mathrm{H}^{+} + 2 \mathrm{e}^- \rightarrow 2 \mathrm{H}_{2} \mathrm{O}\) : \(E^{\circ} = 1.77 \ V\) (in reverse direction) \( \mathrm{Al}^{3+} + 3 \mathrm{e}^- \rightarrow \mathrm{Al}\) : \(E^{\circ} = -1.66 \ V\) (in reverse direction)

Step 2: Arrange in increasing order

Now that we know the standard reduction potentials for the given reducing agents, we can arrange them in order of increasing strength. Remember, the lower the reduction potential, the stronger the reducing agent: \(\mathrm{H}_{2} \mathrm{O}_{2}, \mathrm{Sn}^{2+}, \mathrm{I}^{-}, \mathrm{Zn}, \mathrm{Al}\).

Key Concepts

Standard Reduction Potential

The standard reduction potential, represented by the symbol \(E^\circ\), is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Each half-reaction has a standard reduction potential, which can be found in tables of standard electrode potentials. The value is determined under standard conditions, which includes a temperature of 298 K, a 1 M concentration for each ion participating in the reaction, and a 1 bar pressure for gases.

In the context of redox reactions, the more positive the \(E^\circ\) value, the greater the species' affinity for electrons—indicating a strong oxidizing agent. Conversely, a more negative \(E^\circ\) value suggests a stronger reducing agent, since it indicates a greater tendency to lose electrons. When comparing \(E^\circ\) values, we can predict the direction of electron flow between two half-reactions in a redox reaction, with electrons moving from the half-reaction with the lower \(E^\circ\) to the higher.

Redox Reactions

Redox reactions are a family of chemical reactions that involve the transfer of electrons between two species. The term 'redox' is a portmanteau for reduction-oxidation. All redox processes involve the movement of electrons from one atom or molecule to another, with one species being reduced as it gains electrons, and the other oxidized as it loses electrons.

These reactions are characterized by two half-reactions: the oxidation half-reaction, where electrons are lost; and the reduction half-reaction, where electrons are gained. It’s important to note that in a redox reaction, there must be a change in the oxidation states of the species involved. The substance that loses electrons is known as the reducing agent (or reductant), and the substance that gains electrons is the oxidizing agent (or oxidant).

Chemical Oxidation

Chemical oxidation refers to the process in which a substance loses electrons and increases its oxidation state. This involves the transfer of electrons from the substance being oxidized to another substance. Oxidation can occur in the presence of oxygen leading to common forms like rusting, but it includes any reaction where atoms increase their oxidation state.

Oxidizing agents, or oxidants, are substances that drive the oxidation process by gaining electrons and themselves becoming reduced. In the exercise provided, substances like \(\mathrm{Cl}_2\) and \(\mathrm{H}_2\mathrm{O}_2\) act as potent oxidizing agents due to their high standard reduction potentials, thus they readily accept electrons from other substances. The strength of an oxidizing agent can often be inferred from its standard reduction potential.

Chemical Reduction

Chemical reduction is the counterpart to oxidation, focusing on the gain of electrons by a substance, which as a result decreases its oxidation state. During reduction, a reducing agent, or reductant, donates electrons to another species, resulting in the oxidation of the reducing agent.

Reducing agents are typically elements or compounds that readily lose or donate electrons, making them excellent electron donors. This can be seen in the exercise where metals like \(\mathrm{Zn}\) and \(\mathrm{Al}\) function as reducing agents. These substances have relatively low (negative) standard reduction potentials when written as oxidation half-reactions, meaning they have a strong tendency to lose electrons and cause reduction in other substances.