Question

Consider the following species. $$ \begin{array}{llll} \mathrm{Cr}^{3+} & \mathrm{Hg}(l) & \mathrm{H}_{2} \text { (acidic) } & \mathrm{Sn}^{2+} \end{array} $$ \(\mathrm{Br}_{2}\) (acidic) Classify each species as oxidizing agent, reducing agent, or both. Arrange the oxidizing agents in order of increasing strength. Do the same for the reducing agents.

Short answer

Question: Arrange the given oxidizing agents and reducing agents in increasing strength order and identify their classifications. Oxidizing agents: 1. Cr^3+ 2. Hg(l) 3. Br2 (acidic) Reducing agents: 1. H2(acidic) 2. Sn^2+ Order of increasing strength: Oxidizing agents: Cr^3+ < Hg(l) < Br2(acidic) Reducing agents: Sn^2+ < H2(acidic)

Step-by-step solution

Identify Oxidizing and Reducing Agents

For this step, consult a standard reduction potential table and determine which species are more likely to undergo reduction or oxidation. From the table, we found: - \(\mathrm{Cr}^{3+}\) : reduction (Cr^3+ + 3e^- -> Cr) - \(\mathrm{Hg}(l)\) : reduction (Hg^2+ + 2e^- -> Hg) - \(\mathrm{H}_{2}\) (acidic) : oxidation (2H+ + 2e^- -> H2) - \(\mathrm{Sn}^{2+}\) : reduction (Sn^2+ + 2e^- -> Sn) - \(\mathrm{Br}_{2}\) (acidic) : reduction (2Br^- -> Br2 + 2e^-) So the oxidizing agents are: \(\mathrm{Cr}^{3+}\), \(\mathrm{Hg}(l)\), and \(\mathrm{Br}_2\). The reducing agents are: \(\mathrm{H}_2\) (acidic) and \(\mathrm{Sn}^{2+}\).

Arrange Oxidizing Agents in Order of Increasing Strength

For this step, consult the standard reduction potential table and find the reduction potentials (E°) for the oxidizing agents. - \(\mathrm{Cr}^{3+}\) : E° = -0.74V - \(\mathrm{Hg}(l)\) : E° = 0.85V - \(\mathrm{Br}_2\) (acidic) : E° = 1.09V The order of increasing strength is: $$\mathrm{Cr}^{3+} < \mathrm{Hg}(l) < \mathrm{Br}_2$$

Arrange Reducing Agents in Order of Increasing Strength

For this step, consult the standard reduction potential table and find the oxidation potentials (E°, which are the negative values of reduction potentials) for the reducing agents. - \(\mathrm{H}_{2}\) (acidic) : E° = 0V - \(\mathrm{Sn}^{2+}\) : E° = -0.14V The order of increasing strength is: $$\mathrm{Sn}^{2+} < \mathrm{H}_{2}(acidic)$$ In conclusion: - The oxidizing agents, in increasing strength order, are: \(\mathrm{Cr}^{3+}\), \(\mathrm{Hg}(l)\), \(\mathrm{Br}_2\) (acidic). - The reducing agents, in increasing strength order, are: \(\mathrm{Sn}^{2+}\), \(\mathrm{H}_{2}\)(acidic).

Key concepts

Electrochemistry

Electrochemistry is a branch of chemistry that studies the relationship between electricity and chemical reactions. It is the groundwork for understanding how batteries work, how metals corrode, and even how certain types of chemical reactions proceed. In electrochemistry, reactions involve the transfer of electrons between species in solution or at the interface of an electrode.

An important concept here is the oxidation-reduction (redox) reaction, where oxidation refers to the loss of electrons, and reduction refers to the gain of electrons. Substances that cause other substances to oxidize, accepting electrons in the process, are called oxidizing agents, while those that cause reduction, losing electrons as a result, are known as reducing agents. By determining which species are likely to donate electrons (reducing agents) and which are likely to accept electrons (oxidizing agents), as done in the exercise, we can predict how a chemical reaction might proceed.

Standard Reduction Potentials

The standard reduction potential (also known as standard electrode potential) is a measure of the tendency of a chemical species to be reduced, that is, to gain electrons and decrease its oxidation state. In electrochemistry, these potentials are measured in volts and are provided under standard conditions (298 K, 1 M concentration, and 1 atm pressure). These values are crucial for determining the direction and spontaneity of redox reactions.

In the given exercise, standard reduction potentials were used to classify various species as oxidizing or reducing agents, and then to arrange them in order of increasing strength. A more positive reduction potential indicates a stronger oxidizing agent because it has a higher tendency to gain electrons. Conversely, a more negative reduction potential signifies a stronger reducing agent, as it will more readily lose electrons. For example, \(\mathrm{Br}_2\) with a higher positive standard reduction potential is a stronger oxidizing agent compared to \(\mathrm{Cr}^{3+}\) and \(\mathrm{Hg}(l)\).

Chemical Reactions in Acidic Solution

When a chemical reaction occurs in acidic solution, it typically implies that the surrounding solution has an excess of \(\mathrm{H}^+\) ions. These \(\mathrm{H}^+\) ions can participate in redox reactions, influencing the behavior of other species. For instance, in the exercise, hydrogen gas (H2) is involved as a reducing agent, with the reaction \(2\mathrm{H}^+ + 2e^- \rightarrow \mathrm{H}_2\) taking place in an acidic environment.

It's essential to consider the acidic conditions when balancing redox equations, as the \(\mathrm{H}^+\) ions and corresponding water molecules are involved in the half-reactions. These conditions are typically set for reactions involving species like \(\mathrm{H}_2\) and halogens, as seen with bromine \(\mathrm{Br}_2\) in the exercise. The \(\mathrm{H}^+\) from the acidic solution can impact the electron transfer process and thus change the standard reduction potentials, which are used to evaluate the strength of oxidizing and reducing agents in the given environment.