Question

Consider the following species. \(\begin{array}{llll}\mathrm{Cr}^{3+} & \mathrm{Hg}(l) & \mathrm{H}_{2} \text { (acidic) } & \mathrm{Sn}^{2+} & \mathrm{Br}_{2} \text { (acidic) }\end{array}\) 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: Identify the oxidizing agents and reducing agents in the given list, then arrange them in order of increasing strength: \(\mathrm {Cr^{3+}, Hg(l), H_{2}(acidic), Sn^{2+}, Br_{2}(acidic)}\). Answer: Oxidizing agents (in order of increasing strength): 1. \(\mathrm{Cr^{3+}}\) 2. \(\mathrm{Br_{2}(acidic)}\) Reducing agents (in order of increasing strength): 1. \(\mathrm{Hg(l)}\) 2. \(\mathrm{H_{2}(acidic)}\) \(\mathrm{Sn^{2+}}\) can act as both an oxidizing and reducing agent, depending on the situation.

Step-by-step solution

Identify Oxidizing and Reducing Agents

To identify whether a species is an oxidizing or reducing agent, we will examine its potential to accept or donate electrons in a redox reaction. 1. \(\mathrm{Cr}^{3+}\): Chromium will likely accept electrons to be reduced to \(\mathrm{Cr}^{2+}\) or Cr. 2. \(\mathrm{Hg}(l)\): Mercury may readily lose two electrons to become \(\mathrm{Hg}^{2+}\). 3. \(\mathrm{H}_{2} \text {(acidic)}\): Hydrogen will probably lose electrons to form \(\mathrm{H}^{+}\). 4. \(\mathrm{Sn}^{2+}\): Tin may either accept electrons to form Sn or lose two electrons to become \(\mathrm{Sn}^{4+}\). 5. \(\mathrm{Br}_{2} \text {(acidic)}\): Bromine will likely accept electrons to form \(\mathrm{Br}^{-}\). Now we know whether each species acts as an oxidizing agent, reducing agent, or both.

Arrange Oxidizing Agents in Order of Increasing Strength

Based on the standard electrode potential table, we can arrange the oxidizing agents in order of increasing strength as follows: 1. \(\mathrm{Cr}^{3+}\) 2. \(\mathrm{Br}_{2} \text {(acidic)}\)

Arrange Reducing Agents in Order of Increasing Strength

Now, using the standard electrode potential table, we can arrange the reducing agents in order of increasing strength as follows: 1. \(\mathrm{Hg}(l)\) 2. \(\mathrm{H}_{2} \text {(acidic)}\)

List the Oxidizing and Reducing Agents and their Strength Order

Finally, we can summarize the results as follows: Oxidizing Agents (in order of increasing strength): 1. \(\mathrm{Cr}^{3+}\) 2. \(\mathrm{Br}_{2} \text {(acidic)}\) Reducing Agents (in order of increasing strength): 1. \(\mathrm{Hg}(l)\) 2. \(\mathrm{H}_{2} \text {(acidic)}\) \(\mathrm{Sn}^{2+}\) can act as both an oxidizing and reducing agent, depending on the situation. Keep in mind that these rankings of oxidizing and reducing agent strength are based on their standard electrode potentials, and could change based on other factors, such as different concentrations or temperatures.