How is oxidation defined in chemical reactions?

Oxidation in chemical reactions is defined as the loss of electrons. This fundamental concept is rooted in redox (reduction-oxidation) reactions, where one species undergoes oxidation and another undergoes reduction. When a substance loses electrons, its oxidation state increases, reflecting the change that occurs during the reaction.

For instance, consider the reaction of metals with oxygen. A metal atom, such as sodium, can lose an electron to form a sodium ion. This process is oxidation, and it results in the metal becoming more positively charged. In this context, the ability to identify the loss of electrons not only helps in understanding redox reactions but also in balancing equations and predicting the behavior of elements in different reactions.

The other options do not accurately describe oxidation. The gain of protons pertains to a different type of reaction that does not involve the electron transfer we associate with oxidation and reduction. Similarly, gaining electrons is characteristic of reduction, while the loss of neutrons does not influence the oxidation state of an atom and is related to nuclear reactions rather than chemical changes.