Question

Apply The following equations show redox reactions that are sometimes used in the laboratory to generate pure nitrogen gas and pure dinitrogen monoxide gas (nitrous oxide, \(\mathrm{N}_{2} \mathrm{O} )\) $$\mathrm{NH}_{4} \mathrm{NO}_{2}(\mathrm{s}) \rightarrow \mathrm{N}_{2}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$$ $$\mathrm{NH}_{4} \mathrm{NO}_{3}(\mathrm{s}) \rightarrow \mathrm{N}_{2} \mathrm{O}(\mathrm{g})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l})$$ a. Determine the oxidation number of each element in the two equations, and then make diagrams showing the changes in oxidation numbers that occur in each reaction. b. Identify the atom that is oxidized and the atom that is reduced in each of the two reactions. c. Identify the oxidizing and reducing agents in each of the two reactions. d. Write a sentence telling how the electron transfer taking place in these two reactions differs from that taking place here $$2 \mathrm{AgNO}_{3}+\mathrm{Zn} \rightarrow \mathrm{Zn}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}$$

Short answer

In the given redox reactions, nitrogen is both oxidized and reduced within the same reaction, with \(\mathrm{NH}_{4}\) acting as the reducing agent and \(\mathrm{NO}_{2}\) or \(\mathrm{NO}_{3}\) as the oxidizing agent. In contrast, in the reaction \(2 \mathrm{AgNO}_{3}+\mathrm{Zn} \rightarrow \mathrm{Zn}(\mathrm{NO}_{3})_{2}+2 \mathrm{Ag}\), electron transfer involves the reduction of silver ions by zinc, where nitrogen retains the same oxidation state throughout the reaction.

Step-by-step solution

a. Determine the oxidation number of each element

First, let's determine the oxidation numbers by following the general rules: 1. The oxidation number of an element in its elemental form is 0. 2. The oxidation number of hydrogen is generally +1 when bonded to non-metals, and -1 when bonded to metals. 3. The oxidation number of oxygen is generally -2. 4. The sum of the oxidation numbers in a compound must be equal to its overall charge. For the first equation \(\mathrm{NH}_{4}\mathrm{NO}_{2} (\mathrm{s}) \rightarrow \mathrm{N}_{2} (\mathrm{g}) + 2 \mathrm{H}_{2}\mathrm{O}(\mathrm{l})\): - In \(\mathrm{NH}_{4}\), N has an oxidation number of -3 and H has +1. - In \(\mathrm{NO}_{2}\), N has an oxidation number of +3 and O has -2. - In \(\mathrm{N}_{2}\), N has an oxidation number of 0. - In \(\mathrm{H}_{2}\mathrm{O}\), H has an oxidation number of +1 and O has -2. For the second equation \(\mathrm{NH}_{4}\mathrm{NO}_{3}(\mathrm{s}) \rightarrow \mathrm{N}_{2}\mathrm{O}(\mathrm{g}) + 2\mathrm{H}_{2}\mathrm{O}(\mathrm{l})\): - In \(\mathrm{NH}_{4}\), N has an oxidation number of -3 and H has +1. - In \(\mathrm{NO}_{3}\), N has an oxidation number of +5 and O has -2. - In \(\mathrm{N}_{2}\mathrm{O}\), N has an oxidation number of +1 (for the central N) and -2 (for the two peripheral N). #b. Identify the atom that is oxidized and the atom that is reduced in each of the two reactions.

b. Identify the atom that is oxidized and the atom that is reduced

Now, we can identify which atom is oxidized and which is reduced by analyzing the changes in oxidation numbers: For the first reaction: - Nitrogen in \(\mathrm{NH}_{4}\) goes from -3 to 0 (oxidized). - Nitrogen in \(\mathrm{NO}_{2}\) goes from +3 to 0 (reduced). For the second reaction: - Nitrogen in \(\mathrm{NH}_{4}\) goes from -3 to +1 (oxidized). - Nitrogen in \(\mathrm{NO}_{3}\) goes from +5 to -2 (reduced). #c. Identify the oxidizing and reducing agents in each of the two reactions.

c. Identify the oxidizing and reducing agents

The terms "oxidizing agent" and "reducing agent" refer to the species that cause another species to be oxidized or reduced. In general, the oxidizing agent is the species getting reduced while the reducing agent is the species getting oxidized. For the first reaction: - Oxidizing agent: \(\mathrm{NO}_{2}\) - Reducing agent: \(\mathrm{NH}_{4}\) For the second reaction: - Oxidizing agent: \(\mathrm{NO}_{3}\) - Reducing agent: \(\mathrm{NH}_{4}\) #d. Write a sentence telling how the electron transfer taking place in these two reactions differs from that taking place here.

d. Comparing the electron transfer

The electron transfer in the first two reactions involves nitrogen-containing species, with nitrogen being both oxidized and reduced within the same reaction, whereas in the provided reaction \(2 \mathrm{AgNO}_{3}+\mathrm{Zn} \rightarrow \mathrm{Zn}(\mathrm{NO}_{3})_{2}+2 \mathrm{Ag}\), the electron transfer involves the reduction of silver ions by zinc, with nitrogen remaining in the same oxidation state throughout the reaction.

Key concepts

Oxidation Number

The concept of oxidation number is crucial in understanding redox reactions. It represents the hypothetical charge an atom would have if all bonds were completely ionic. Calculating oxidation numbers helps identify which elements are oxidized or reduced in a reaction.
To determine oxidation numbers follow these rules:

• The oxidation number for any element in its standard state is 0.
• For hydrogen, it's typically +1, except when bonded with metals, where it's -1.
• Oxygen is usually -2 in compounds.
• The total oxidation numbers for all atoms in a neutral compound sum to zero.

For example, in the reaction \({ NH}_{4}{NO}_{2} \rightarrow { N}_{2} + 2 { H}_{2}\o \), nitrogen's state changes from -3 in \({ NH}_{4}\) to +3 and then 0. Another nitrogen atom begins at +3 and reduces to 0.

Oxidizing Agent

An oxidizing agent enables another element to lose electrons by accepting electrons itself. Thus, the oxidizing agent is reduced in the reaction. Its role is vital because it drives the electron flow that leads to oxidation.
In \({ NH}_{4}{ NO}_{2} \rightarrow { N}_{2} + 2{ H}_{2}{ O}\), \({ NO}_{2}\) acts as the oxidizing agent. Here, nitrogen in \({ NO}_{2}\) goes from an oxidation state of +3 to 0 as it gains electrons and becomes reduced. \({ NO}_{3}\) performs a similar role in the reaction \({ NH}_{4}{ NO}_{3} \rightarrow { N}_{2}{ O} + 2{ H}_{2}{ O}\), reducing the oxidation state of nitrogen from +5 to +1.

Reducing Agent

Reducing agents are substances that donate electrons to another species, causing the reduction of that species while the reducing agent itself is oxidized. Identifying the reducing agent is essential to comprehend the flow of electrons in redox reactions.
In the reaction \({ NH}_{4}{ NO}_{2} \rightarrow { N}_{2} + 2{ H}_{2}{ O}\), \({ NH}_{4}\) is the reducing agent. It loses electrons, with nitrogen going from -3 to 0. Similarly, in \({ NH}_{4}{ NO}_{3}\), nitrogen in \({ NH}_{4}\) is oxidized from -3 to +1, demonstrating that \({ NH}_{4}\) serves as the reducing agent in the reaction, donating electrons to the nitrogen in \({ NO}_{3}\).

Electron Transfer

Electron transfer is a fundamental component of redox reactions. It describes how electrons are transferred from the reducing agent to the oxidizing agent, altering their oxidation states. Understanding electron transfer reveals the mechanics of energy conversion at the atomic level.
In the given reactions, electron transfer primarily involves nitrogen atoms within the compounds, unlike the distinct, separate electron transfer observed in the reaction \({ 2AgNO}_{3} + Zn \rightarrow { Zn}(NO_{3})_{2} + 2Ag\). In this reaction, zinc transfers electrons directly to silver ions, resulting in the reduction of silver and the oxidation of zinc. Here, distinct elements transfer electrons between each other, standing out from the internal nitrogen electron exchange seen in the decomposition of ammonium nitrate and nitrite.