Question

Specify which of the following are oxidation-reduction reactions, and identify the oxidizing agent, the reducing agent, the substance being oxidized, and the substance being reduced. a. \(\mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q)\) b. \(\mathrm{HCl}(g)+\mathrm{NH}_{3}(g) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(s)\) c. \(\mathrm{SiCl}_{4}(l)+2 \mathrm{H}_{2} \mathrm{O}(l) \rightarrow 4 \mathrm{HCl}(a q)+\mathrm{SiO}_{2}(s)\) d. \(\mathrm{SiCl}_{4}(l)+2 \mathrm{Mg}(s) \rightarrow 2 \mathrm{MgCl}_{2}(s)+\mathrm{Si}(s)\) e. \(\mathrm{Al}(\mathrm{OH})_{4}^{-}(a q) \rightarrow \mathrm{AlO}_{2}^{-}(a q)+2 \mathrm{H}_{2} \mathrm{O}(i)\)

Short answer

In conclusion, only reactions a and d are oxidation-reduction reactions. For reaction a: - Oxidizing agent: Ag¹⁺(aq) - Reducing agent: Cu(s) - Substance being oxidized: Cu(s) - Substance being reduced: Ag¹⁺(aq) For reaction d: - Oxidizing agent: SiCl₄(l) - Reducing agent: Mg(s) - Substance being oxidized: Mg(s) - Substance being reduced: SiCl₄(l)

Step-by-step solution

Identify reaction type and changes in oxidation numbers

To see if the reaction is redox, we will calculate the oxidation numbers of elements before and after the reaction: Cu: 0 -> +2 Ag: +1 -> 0 Cu has increased in oxidation number (lost electrons) and thus, undergoes oxidation, turning to Cu²⁺(aq). Ag¹⁺(aq) has decreased in oxidation number (gained electrons) and thus undergoes reduction, turning to Ag(s).

Identify the oxidizing and reducing agents

The oxidizing agent is Ag¹⁺(aq), and the reducing agent is Cu(s). The substance being oxidized is Cu(s), and the substance being reduced is Ag¹⁺(aq). b.

Identify reaction type and changes in oxidation numbers

Checking the oxidation numbers for each element: Cl: -1 -> -1 H: +1 -> +1 N: -3 -> -3 No change in oxidation numbers, hence, this is not an oxidation-reduction reaction. c.

Identify reaction type and changes in oxidation numbers

Checking the oxidation numbers for each element: Si: +4 -> +4 Cl: -1 -> -1 O: -2 -> -2 H: +1 -> +1 No change in oxidation numbers, hence, this is not an oxidation-reduction reaction. d.

Identify reaction type and changes in oxidation numbers

Checking the oxidation numbers for each element: Si: +4 -> 0 Mg: 0 -> +2 Cl: -1 -> -1 Si has decreased in oxidation number (gained electrons) and thus undergoes reduction, turning to Si(s). Mg undergoes oxidation, losing electrons, and having its oxidation number increase to form Mg²⁺(aq).

Identify the oxidizing and reducing agents

The oxidizing agent is SiCl₄(l), and the reducing agent is Mg(s). The substance being oxidized is Mg(s), and the substance being reduced is SiCl₄(l). e.

Identify reaction type and changes in oxidation numbers

Checking the oxidation numbers for each element: Al: +3 -> +3 O: -2 -> -2 H: +1 -> +1 No change in oxidation numbers, hence, this is not an oxidation-reduction reaction. In conclusion, a and d are oxidation-reduction reactions, with the oxidizing agents being Ag¹⁺(aq) and SiCl₄(l), the reducing agents are Cu(s) and Mg(s), the substance being oxidized being Cu(s) and Mg(s), and the substance being reduced is Ag¹⁺(aq) and SiCl₄(l), respectively.

Key concepts

Oxidation Numbers

Oxidation numbers are a helpful way to track how electrons are transferred in a chemical reaction. Think of them as imaginary charges that help us figure out if an element has gained or lost electrons. In oxidation-reduction (redox) reactions, some elements will see a change in their oxidation numbers.

• If an element's oxidation number increases, this means it lost electrons, which is called oxidation.
• If an element's oxidation number decreases, it gained electrons, meaning it underwent reduction.

By calculating oxidation numbers before and after a reaction, we can identify which elements have been oxidized and reduced. For example, when examining the reaction \( \mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \), copper's oxidation number changes from 0 to +2, indicating it has undergone oxidation.

Oxidizing Agent

The oxidizing agent in a redox reaction is the substance that accepts electrons. It is reduced in the process, which means its oxidation number decreases.

• The oxidizing agent becomes reduced by taking electrons from another substance.
• It helps in oxidizing another element by gaining electrons itself.
• For example, in the reaction \( \mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \), Ag⁺ ions are the oxidizing agents because they accept electrons from copper, reducing to Ag(s).

Recognizing the oxidizing agent is essential for understanding the electron transfer process in redox reactions.

Reducing Agent

The reducing agent in a redox reaction is the substance that donates electrons. Its oxidation number increases as it is oxidized.

• The reducing agent is oxidized as it gives away electrons.
• It helps in reducing another substance by losing electrons itself.
• In the reaction \( \mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \), copper is the reducing agent because it donates electrons to Ag⁺, becoming Cu²⁺.

Understanding the role of the reducing agent helps clarify how substances interchange electrons in a redox reaction.

Redox Reaction

A redox reaction, short for oxidation-reduction reaction, involves the transfer of electrons between substances.

• One substance is oxidized (loses electrons), while another is reduced (gains electrons).
• Identifying redox reactions involves looking for changes in oxidation numbers.
• Reactions such as \( \mathrm{Cu}(s)+2 \mathrm{Ag}^{+}(a q) \rightarrow 2 \mathrm{Ag}(s)+\mathrm{Cu}^{2+}(a q) \) are classic examples of redox processes.

In these reactions, you will often need to identify what is being oxidized and reduced and determine the oxidizing and reducing agents. Redox reactions play a crucial role in energy production, metabolism, and many industrial processes.