Question

Specify which of the following equations represent oxidation– reduction reactions, and indicate the oxidizing agent, the reducing agent, the species being oxidized, and the species being reduced a. \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)\) b. \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)\) c. \(\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)\) d. \(2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\)

Short answer

The redox reactions are a, b, and c. The oxidizing and reducing agents are: a. Oxidizing agent: Water; Reducing agent: Methane b. Oxidizing agent: Silver nitrate; Reducing agent: Copper c. Oxidizing agent: Hydrochloric acid; Reducing agent: Zinc

Step-by-step solution

Identify oxidation-reduction (redox) reactions

For each equation, assign oxidation numbers to the elements. If the oxidation numbers change from reactants to products, the equation is a redox reaction. a. \(CH_4(g)+H_2O(g) \rightarrow CO(g)+3H_2(g)\) - Assign oxidation numbers: C: -4 to +2, H: +1 to 0, O: -2 - Oxidation numbers change, so this is a redox reaction. b. \(2AgNO_3(aq)+Cu(s) \rightarrow Cu(NO_3)_2(aq)+2Ag(s)\) - Assign oxidation numbers: Ag: +1 to 0, N: +5, O: -2, Cu: 0 to +2 - Oxidation numbers change, so this is a redox reaction. c. \(Zn(s)+2HCl(aq) \rightarrow ZnCl_2(aq)+H_2(g)\) - Assign oxidation numbers: Zn: 0 to +2, H: +1 to 0, Cl: -1 - Oxidation numbers change, so this is a redox reaction. d. \(2H^+(aq)+2CrO_4^{2-}(aq) \rightarrow Cr_2O_7^{2-}(aq)+H_2O(l)\) - Assign oxidation numbers: H: +1, Cr: +6 to +6, O: -2 - Oxidation numbers do not change, so this is not a redox reaction.

Determine oxidizing and reducing agents

For each redox reaction, identify the species that cause the oxidation and reduction by determining which species gained or lost electrons. a. \(CH_4(g)+H_2O(g) \rightarrow CO(g)+3H_2(g)\) - Oxidation: C is oxidized from -4 to +2 (loses electrons) - Reduction: H is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Water - Reducing agent: Methane b. \(2AgNO_3(aq)+Cu(s) \rightarrow Cu(NO_3)_2(aq)+2Ag(s)\) - Oxidation: Cu is oxidized from 0 to +2 (loses electrons) - Reduction: Ag is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Silver nitrate - Reducing agent: Copper c. \(Zn(s)+2HCl(aq) \rightarrow ZnCl_2(aq)+H_2(g)\) - Oxidation: Zn is oxidized from 0 to +2 (loses electrons) - Reduction: H is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Hydrochloric acid - Reducing agent: Zinc

Summary

The redox reactions are a, b, and c. The oxidation and reduction species, along with their corresponding agents, are as follows: a. Oxidizing agent: Water; Reducing agent: Methane b. Oxidizing agent: Silver nitrate; Reducing agent: Copper c. Oxidizing agent: Hydrochloric acid; Reducing agent: Zinc

Key concepts

Oxidizing Agent

An oxidizing agent, also known as an oxidant, is a substance that causes another substance to lose electrons. In other words, it facilitates the process of oxidation by itself gaining electrons. This can seem a bit counterintuitive at first because although it "oxidizes" another component, it is actually reduced in the reaction. In the examples provided, the oxidizing agents are:

• Water (\(H_2O\)) in the reaction with methane (\(CH_4\)).
• Silver nitrate (\(AgNO_3\)) in the reaction with copper (\(Cu\)).
• Hydrochloric acid (\(HCl\)) in the reaction with zinc (\(Zn\)).

The role of an oxidizing agent is crucial in redox reactions as it allows for the transfer of electrons. Recognizing the oxidizing agent in a reaction can involve tracking the change in oxidation numbers, which indicate gain of electrons.

Reducing Agent

The reducing agent, or reductant, is the opposite of the oxidizing agent. It is the species that donates electrons and reduces another species. While it reduces the other substance, it gets oxidized itself. Identifying the reducing agent helps to understand who gives up electrons in a chemical reaction. Here are the reducing agents from the provided reactions:

• Methane (\(CH_4\)) donates electrons to water (\(H_2O\)).
• Copper (\(Cu\)) donates electrons to silver nitrate (\(AgNO_3\)).
• Zinc (\(Zn\)) donates electrons to hydrochloric acid (\(HCl\)).

In these reactions, the atoms of the reducing agents increase their oxidation numbers, highlighting their role in electron donation.

Oxidation Numbers

Oxidation numbers (or states) provide a simple way to keep track of electrons in a compound, permitting the identification of oxidizing or reducing agents without guessing. Here's how they change in each example:

• In the reaction with methane, carbon changes from an oxidation number of -4 to +2, while hydrogen changes from +1 to 0. This show carbon's loss and hydrogen's gain of electrons.
• For copper and silver nitrate, copper changes from 0 to +2, and silver changes from +1 to 0, revealing electron transfer between copper and silver.
• In zinc's reaction with hydrochloric acid, zinc changes from 0 to +2 and hydrogen changes from +1 to 0, indicating zinc's electrons going to hydrogen.

Oxidation numbers are crucial for identifying redox reactions and understanding how electrons shift between species. By monitoring how these numbers change, chemists can decode complex reactions and predict product outcomes.