Question

Determine the oxidation state of \(\mathrm{N}\) in the compounds in the reaction for the formation of nitric acid. Identify the oxidizing agent and the reducing agent. $$ 3 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow 2 \mathrm{HNO}_{3}(l)+\mathrm{NO}(g) $$

Short answer

In \(\mathrm{NO}_2\), N has an oxidation state of +4; in \(\mathrm{HNO}_3\), N has an oxidation state of +5; and in \(\mathrm{NO}\), N has an oxidation state of +2. \(\mathrm{NO}_2\) acts as both the reducing and the oxidizing agent, reducing to \(\mathrm{NO}\) and oxidizing to \(\mathrm{HNO}_3\).

Step-by-step solution

Assign oxidation states to nitrogen in reactants and products

According to the rules for oxidation states, the oxidation state of oxygen is usually -2, and hydrogen is +1. For the reactant, in \(\mathrm{NO}_2\), knowing that the sum of oxidation states in a molecule equals the overall charge (which is zero for neutral molecules), we can calculate the oxidation state of nitrogen (N) as follows: \(x + 2(-2) = 0\), solving for x gives \(x = +4\). For the product, in nitric acid (\(\mathrm{HNO}_3\)), we calculate \(1 + x + 3(-2) = 0\), which results in an oxidation state for N of +5. For the product \(\mathrm{NO}\) gas, the calculation is \(x + 2(-2) = 0\), yielding \(x = +2\).

Determine the changes in oxidation state for nitrogen

The oxidation state of nitrogen in \(\mathrm{NO}_2\) changes from +4 to +5 in \(\mathrm{HNO}_3\), which indicates that \(\mathrm{NO}_2\) is being oxidized. In \(\mathrm{NO}\), nitrogen has an oxidation state of +2, which is a reduction from its +4 state in \(\mathrm{NO}_2\). Since oxidation state is decreasing from +4 to +2 in \(\mathrm{NO}\), it is the product of the reduction process.

Identify the oxidizing and reducing agents

The oxidizing agent is the species that is reduced, and the reducing agent is the species that is oxidized. Since \(\mathrm{NO}_2\) is being oxidized and is responsible for the other species (\(\mathrm{NO}\)) being reduced, \(\mathrm{NO}_2\) is the reducing agent. Conversely, as the product \(\mathrm{NO}\) is reduced, the initial \(\mathrm{NO}_2\) that was changed into \(\mathrm{NO}\) must be the oxidizing agent.

Key concepts

Assigning Oxidation Numbers

Understanding how to assign oxidation numbers to elements in compounds is critical for studying chemical reactions, particularly redox reactions. Assigning oxidation numbers, or oxidation states, allows chemists to keep track of electron transfers in chemical reactions.

To assign oxidation numbers, a set of rules is followed. Most elements have consistent oxidation states within compounds, such as hydrogen, which is usually given an oxidation state of +1, and oxygen, which is typically assigned -2. However, there are exceptions, such as in peroxides or when bonded with fluorine.

For a molecule like \( \mathrm{NO}_2 \) the oxidation number of nitrogen is found by accounting for the known oxidation states of the oxygen atoms. Since the molecule is neutral, the sum of the oxidation states must be zero. The rule that the sum of oxidation states equals the charge of the molecule is a pivotal concept in determining oxidation numbers. Through simple algebra, the oxidation number for nitrogen in \( \mathrm{NO}_2 \) can be deduced to be +4.

Exercise Improvement Advice

To enhance a student's grasp of assigning oxidation numbers, it's vital to consistently practice with a variety of chemical species and to memorize the standard oxidation states of common elements. Conceptual understanding can be bolstered by visual aids, such as diagrams showing electron transfer, or interactive exercises that allow for trial-and-error learning.

Redox Reactions

Redox reactions are fundamental processes where oxidation and reduction occur simultaneously. In these reactions, one species loses electrons (oxidation) while another gains electrons (reduction). The term 'redox' is a portmanteau of reduction and oxidation.

Recognition of redox reactions involves tracking the movement of electrons by means of oxidation states. For instance, in the given reaction for the formation of nitric acid, a redox process is at play. The nitrogen in \( \mathrm{NO}_2 \) sees an increase in its oxidation state from +4 to +5 when forming \( \mathrm{HNO}_3 \) indicating it has lost an electron – it has been oxidized. Conversely, when forming \( \mathrm{NO} \) gas, its oxidation state is reduced from +4 in \( \mathrm{NO}_2 \) to +2, signifying a gain of electrons.

Understanding redox reactions involves not only recognizing the changes in oxidation states but also knowing the principles of electron transfer and its conservation. This aids in balancing complex redox equations—one of the more challenging tasks for students. Tutoring resources offering step-by-step guidance on electron counting and half-reaction methods greatly benefit learning and application.

Oxidizing and Reducing Agents

To understand the roles of substances in a redox reaction, it is essential to distinguish between oxidizing and reducing agents. An oxidizing agent, also known as an oxidant, gains electrons and becomes reduced in the process. On the other side, a reducing agent, or reductant, donates electrons and is oxidized.

In our example, \( \mathrm{NO}_2 \) is the reducing agent because it loses electrons, as its oxidation state increases from +4 to +5. During this change, another nitrogen species (\( \mathrm{NO} \) gas) forms with a lower oxidation state, meaning \( \mathrm{NO}_2 \) has donated electrons to that species. The \( \mathrm{NO} \) gas, having a lower oxidation state from its precursor, is considered to have been reduced and is thus the product of the oxidizing agent's action.

To fully understand the role of these agents, it's valuable to consider mnemonic devices such as 'LEO says GER' (Lose Electrons Oxidation, Gain Electrons Reduction) which can help students remember the direction of electron flow. Through the identification of oxidizing and reducing agents, one can predict reaction products, balance equations, and understand the energetics of reactions.