In the manufacture of nitric acid by the oxidation of ammonia, the first product is nitric oxide, which is then cxidized to nitrogen dioxide. From the standard reaction enthalpies $$ \begin{gathered} \mathrm{N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{NO}(\mathrm{g}) \\ \Delta H^{\circ}=+180.5 \mathrm{~kJ} \\ \mathrm{~N}_{2}(\mathrm{~g})+2 \mathrm{O}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{NO}_{2}(\mathrm{~g}) \\ \Delta H^{\circ}=+66.4 \mathrm{~kJ} \end{gathered} $$ calculate the standard reaction enthalpy for the oxidation of nitric oxide to nitrogen dioxide: $$ 2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{~g}) \longrightarrow 2 \mathrm{NO}_{2}(\mathrm{~g}) $$

The oxidation of ammonia is a significant chemical process in the production of nitric acid, an essential industrial and laboratory chemical. Nitric acid is used in fertilizers, explosives, and many other chemical processes. The first step in this multi-step reaction involves the conversion of ammonia (NH\(_3\)) directly into nitric oxide (NO) by its reaction with oxygen (O\(_2\)) under high temperature and pressure in the presence of a catalyst. This process of oxidation indicates that ammonia is losing electrons and is being transformed into a more oxidized state in the form of nitric oxide. Understanding this fundamental chemical change is crucial for grasping the overall reaction sequence that leads to the formation of nitric acid.

When teaching this concept, it's vital to emphasize the balanced chemical equation and the accompanying energy changes that occur during this exothermic reaction, where heat is released as the reaction proceeds. Simplifying the concept with clear and concise explanations, while using visual representations, will help students understand the process and its significance in the chemical industry.

Nitric oxide (NO) is a colorless gas that plays an intermediate role in the formation of nitric acid. In the context of our exercise, nitric oxide is produced when nitrogen (N\(_2\)) and oxygen (O\(_2\)) react together. The reaction is endothermic; it absorbs energy in the form of heat, which is indicated by a positive standard reaction enthalpy. With a keen focus on the production of nitric oxide, students should be aware of its attributes, such as its paramagnetic nature and its roles outside industrial chemistry, including its functions as a signaling molecule in biological systems. A fascinating yet straightforward way to approach this is by relating nitric oxide's diverse roles in different fields, thus capturing the interest of students with varied inclinations, from industrial chemistry to environmental science and physiology.

Hess's Law is a powerful tool in thermodynamics that states the total enthalpy change for a reaction is the same, regardless of the number of steps the reaction is carried out in. This law relies on the principle of the conservation of energy and allows us to calculate the enthalpy change of a reaction that might be difficult to measure directly by using known enthalpies of other related reactions. In educational materials, it is essential to outline the concept with practical examples, as seen in our exercise, where Hess's Law lets us determine the enthalpy change of the oxidation of nitric oxide to nitrogen dioxide by manipulating and combining other reactions with known enthalpy changes. By teaching students to break down complex reactions into simpler steps that are easier to analyze, they can use Hess's Law to solve problems that would otherwise seem daunting.

Enthalpy change, denoted by \(\Delta H\), represents the heat absorbed or released during a chemical reaction at constant pressure. It is an essential concept in understanding energy flow within chemical systems. The sign of \(\Delta H\) indicates whether the reaction is exothermic (releasing heat, negative \(\Delta H\)) or endothermic (absorbing heat, positive \(\Delta H\)). In our exercise, the confusing aspect for some learners might be the positive and negative signs of enthalpy changes and how they affect the overall energy of the system. It's helpful to present this concept by explaining that the enthalpy change essentially reflects the differences in energy between reactants and products. By providing relatable examples that compare everyday energy exchanges, educators can bring clarity to this abstract concept. For instance, comparing exothermic reactions to charging a battery and endothermic reactions to using the energy stored in that battery can make it easier for students to grasp the concept of enthalpy change.