Think of ammonium nitrite as \(\left(\mathrm{NH}_{4}^{+}\right)\left(\mathrm{NO}_{2}^{-}\right)\) and assign oxidation numbers within each ion. Now think of ammonium nitrate as \(\mathrm{N}_{2} \mathrm{H}_{4} \mathrm{O}_{2},\) and assign oxidation numbers. Which assignment makes more sense?

Understanding the chemical makeup and behavior of ammonium nitrite is fundamental in the study of redox chemistry. Ammonium nitrite, with the formula \( \left(\mathrm{NH}_4^+\right)\left(\mathrm{NO}_2^-\right) \), consists of two ions: the ammonium ion (\(\mathrm{NH}_4^+\)) and the nitrite ion (\(\mathrm{NO}_2^-\)).

The oxidation state of nitrogen in the ammonium ion is -3 because hydrogen, which is less electronegative, is assigned a +1 oxidation state by convention. Therefore, thinking in terms of electron transfer, the nitrogen in ammonium effectively has 'gained' electrons compared to its elemental state. On the other hand, the nitrogen in the nitrite ion has an oxidation state of +3, as each oxygen contributes a -2 oxidation state. This illustrates the concept that nitrogen can exhibit multiple oxidation states depending on its chemical environment.

Oxidation states are essential in predicting the reactivity of compounds, particularly in reactions where electron transfer occurs. For instance, ammonium nitrite is known to decompose into nitrogen gas (\(\mathrm{N}_2\)) and water, showcasing typical redox behavior.

Ammonium nitrate is another nitrogenous compound commonly used as a fertilizer and has significant implications in redox chemistry. It is often represented by the formula \(\mathrm{NH}_4\mathrm{NO}_3\), highlighting that it is made up of an ammonium ion and a nitrate ion.

Unlike ammonium nitrite, ammonium nitrate's nitrogen atoms have different oxidation states in its two constituent ions. The oxidation state of nitrogen in ammonium is still -3, but in the nitrate ion (\(\mathrm{NO}_3^-\)), the nitrogen has an oxidation state of +5. This reflects the higher number of oxygen atoms associated with the nitrogen in the nitrate ion compared to the nitrite ion.

The ability for the same element, such as nitrogen, to exhibit various oxidation states is a fundamental principle of redox chemistry and is essential for understanding the transformation and energy exchange in chemical reactions. Ammonium nitrate’s stability and its ability to release oxygen make it a component in certain explosives when combined with fuel oils, showcasing a drastic potential for oxidation-reduction reactions.

Chemical notation is a standardized way of representing chemical compounds and elements, which greatly facilitates the communication of information between chemists and students. In the context of oxidation numbers, chemical notation plays a crucial role.

For example, \(\left(\mathrm{NH}_4^+\right)\left(\mathrm{NO}_2^-\right)\) and \(\mathrm{NH}_4\mathrm{NO}_3\) not only tell us the composition of ammonium nitrite and ammonium nitrate, respectively, but they also imply the electrical neutrality of these compounds. The pluses and minuses indicate the charges on the ions, leading us to consider how these charges balance out and help us in assigning oxidation states.

Significant in oxidation state determination is recognizing that in chemical compounds, the sum of the oxidation numbers of all atoms must equal the overall charge of the compound. Misinterpreting notation or failing to apply these principles can lead to incorrect assignments of oxidation states and misunderstanding of the compound's characteristics and potential reactions.

Redox chemistry, short for reduction-oxidation chemistry, is the study of reactions where electrons are transferred between substances. The term 'redox' itself comes from the concept that one reactant is reduced (gains electrons) and the other is oxidized (loses electrons).

Oxidation numbers are not just bookkeeping devices but insightful indicators for predicting redox reactions. In the case of ammonium nitrite and ammonium nitrate, understanding the oxidation states of nitrogen helps us predict its behavior. Nitrogen exhibits a versatile range of oxidation states from -3 to +5, indicating its potential to participate in diverse redox reactions.

In teaching redox chemistry, it is important to explain that the assignment of oxidation numbers, while seeming arbitrary at times, follows specific rules. These rules are designed to reflect the electron distribution in molecules and ions in an idealized and simplified manner. They help predict the transfer of electrons, which is the central theme of redox reactions, and thus drive the transformation of matter and the flow of energy within chemical systems.