Name each of the following complex ions and determine the oxidation number of the metal: (a) \(\left[\mathrm{CrCl}_{3}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{OH}_{2}\right)\right]^{+}\); (b) \(\left.\left[\mathrm{Rh}(\mathrm{en})_{3}\right)\right]^{3+}\); (c) \(\left[\mathrm{Fe}(\mathrm{ox})(\mathrm{Br})_{4}\right]^{3-}\); (d) \(\left[\mathrm{Ni}(\mathrm{OH})\left(\mathrm{OH}_{2}\right)_{5}\right]^{2+}\).

Understanding the oxidation states of elements in coordination compounds is pivotal for discerning their nomenclature and properties. For students grappling with oxidation number calculation, here's a straightforward formula: The sum of oxidation numbers of all atoms should equal the overall charge of the compound.

Let's consider a complex ion, such as \(\left[\mathrm{CrCl}_{3}\left(\mathrm{NH}_{3}\right)_{2}\left(\mathrm{OH}_{2}\right)\right]^{+}\). Here, the ammine (\(\mathrm{NH}_{3}\)) and aqua (\(\mathrm{H}_{2}\mathrm{O}\)) ligands have no charge, thus contributing zero to the oxidation number. Chlorido ligands (Cl), on the other hand, each bear a -1 charge. Knowing that the complex bears a +1 charge overall, we can infer that chromium (Cr) must have a +3 oxidation number to balance the -3 from the chlorido ligands and have a sum that equals the charge of the complex ion.

To recap, the steps to calculate the oxidation number involve:

• Listing all ligand charges.
• Applying the formula where the sum of all charges equals the overall charge.
• Inferring the oxidation number for the central atom.

Once mastered, the calculation will become second nature in the study of coordination chemistry.

When moving on to names within coordination compounds, students often stumble upon the various ligand naming conventions. A core principle to remember is that ligands are listed alphabetically regardless of their quantity when naming complex ions.

Common ligands like water, ammonia, carbon monoxide, and halides have specific names when they are a part of coordination complexes. Water becomes 'aqua', ammonia 'ammine', carbon monoxide 'carbonyl', chloride 'chlorido', and so on. Knowledge of these conventions is crucial for writing correct complex ion names.

For examples from our exercise, the ligands include 'chloro' (for Cl), 'ammine' (for \(\mathrm{NH}_{3}\)), 'aqua' (for \(\mathrm{H}_{2}\mathrm{O}\)), 'ethylenediamine' (for en), 'bromido' (for Br), and 'oxalato' (for the oxalate ion \(\mathrm{C}_{2}\mathrm{O}_{4}^{2-}\)). Keeping these names in mind will ensure you always present the ligands in the correct form, such as 'trichlorobis(ammine)(aqua)' in one of our given examples.

Using prefixes like 'bis', 'tris', 'tetrakis', etc., for identical ligands that could cause confusion is also common practice – as seen with 'tris(ethylenediamine)' to describe multiple ethylenediamine ligands.

Finally, mastering coordination compound terminology will enable students to navigate the intricacies of naming these chemicals fluently. Begin with the central metal ion, distinguishing its name based on the overall charge of the compound: for neutral or positively charged complexes, use the regular metal name, whereas for anionic complexes, an '-ate' suffix is affixed to the metal name, such as 'ferrate' for iron.

In the nomenclature, the state of the central ion is designated by Roman numerals in parentheses immediately following the metal’s name – this indicates the oxidation state, a critical piece of information. The combination of ligand names followed by the central metal creates the full name of the coordination compound.

For instance, massive molecules involve terms like 'hexa', 'penta', or 'tetra', reflective of the number of each type of ligand present. These terms combine smoothly to generate names like 'hexachloroplatinate(IV)' for a platinum complex with six chloride ions.

Embracing these naming rules not only assists in homework but paves the way for a deeper understanding of complex ions and their reactions, which are ubiquitous in diverse areas from biological systems to materials science.