Assign oxidation numbers to each element in (a) methane (b) sulfurous acid (c) sodium oxide (d) the dihydrogen phosphate ion \(\left(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right)\)

Understanding oxidation state rules is essential for mastering redox chemistry and determining how electrons are transferred between atoms in chemical reactions. The oxidation state, often referred to as oxidation number, is a theoretical charge on an atom if electrons were completely transferred.

Some key rules are:

• An element in its pure form, like O2 or H2, always has an oxidation number of 0.
• For monatomic ions, such as Na+ or Cl-, the oxidation number is the same as the ion's charge.
• In molecules or compounds, the total sum of oxidation numbers must equal the overall charge. For example, in a neutral molecule, the sum is zero.
• Typically, oxygen has an oxidation number of -2, and hydrogen +1, which are essential in balancing equations.
• Alkali metals (group 1) have an oxidation number of +1, and alkaline earth metals (group 2) have +2.

When applying these rules to complex molecules, it helps to start with the most electronegative elements, assigning their common oxidation states, and then work towards the less electronegative ones. This method simplifies the process and ensures accuracy in assigning the correct states to each element.

Redox chemistry revolves around the concept of oxidation and reduction reactions, which involve the transfer of electrons between substances. The species that loses electrons is oxidized, while the one gaining electrons is reduced. Assigning accurate oxidation numbers is crucial for identifying which species is oxidized and which is reduced during a redox reaction.

In the methane molecule, for instance, carbon is oxidized as it has a -4 oxidation state, implying it has gained electrons compared to its elemental state. Contrastingly, in sulfurous acid, sulfur is in a +4 oxidation state, suggesting it has lost electrons compared to its elemental form. The shifts in oxidation states reveal the underlying redox processes.

These shifts also dictate the stoichiometry in redox reactions, as the number of electrons lost by the oxidizing agent must equal the number gained by the reducing agent. Balancing redox reactions involves a careful analysis of the changes in oxidation states and ensures the conservation of both mass and charge.

Chemical nomenclature refers to the systematic naming of chemical compounds and ions based on their structure and composition. Understanding chemical names is not just a matter of memorization; it provides insights into the very nature of the substances involved. For instance, in the name 'dihydrogen phosphate ion,' 'dihydrogen' signifies two hydrogen atoms are present, 'phosphate' means a phosphorous atom bonded to four oxygen atoms is central to the ion, and the suffix 'ion' indicates a net charge is present.

Assigning oxidation numbers also plays a part in nomenclature, especially when an element can have multiple possible oxidation states. For sulfur in sulfurous acid, the +4 oxidation number is lower than the maximum +6 it could have, thus the 'ous' suffix is used, distinguishing it from sulfuric acid where sulfur has a +6 state.

In nomenclature, familiarity with oxidation states can help predict chemical properties and reactivity. Moreover, naming compounds correctly is an integral part of communication in the scientific community, allowing for clear and precise discussions of chemical behavior and reaction mechanisms.