Consider the perchlorate \(\left(\mathrm{ClO}_{4}^{-}\right)\) ion. (a) Assign oxidation states to all of the atoms in perchlorate. (b) Explain why the chlorine atom in perchlorate does not follow the halide rule. (c) Do you think perchlorate would be a powerful oxidizing agent or a powerful reducing agent? Explain. (d) Would it be possible to have an even higher oxidation state for the \(C 1\) atom in some other compound? Explain. (Hint: Chlorine is in group VIIA).

To truly grasp the concept of oxidation states, it's essential to recognize that they represent the hypothetical charge an atom would have if it were in a compound, with electrons assigned according to certain rules. In the perchlorate ion, \(\text{ClO}_4^{-}\), assigning oxidation states starts with remembering that oxygen typically carries an oxidation state of -2. Since there are four oxygen atoms, you'll quickly conclude that their combined contribution is -8. As the whole ion has only a -1 charge, the math reveals that the chlorine must have an oxidation state of +7 to balance this out.

In more general terms, a positive oxidation state denotes the potential to accept electrons (and thereby being reduced), whereas a negative state indicates the ability to donate electrons (and thereby being oxidized). Recognizing these will help you discern the role each element can play in redox reactions.

The halide rule, an oxidation states guideline, suggests that halogens like chlorine, fluorine, iodine, and bromine, typically exhibit an oxidation state of -1 when they transform into ions. However, perchlorate throws a curveball, flouting this rule with chlorine's oxidation state of +7. This anomaly is due to the remarkable ability of chlorine to expand beyond the typical octet of electrons through its d-orbitals, allowing it to bond to an excess of electron-withdrawing oxygen atoms.

The Expanded Octet

Highlighting the role of d-orbitals in chlorine's capacity to form larger compounds can help demystify this deviation. By leveraging available d-orbitals for bonding, halogens can surround themselves with more electrons than the octet rule would typically allow, leading to these higher oxidation states.

An oxidizing agent is essentially a chemical catcher of electrons—it receives electrons from other substances in a chemical reaction. Perchlorate's chemistry makes it a strong one owing to the high oxidation state of chlorine (+7). This lofty state of +7 makes chlorine in perchlorate ion restless, eager to 'step down' by gaining electrons and reverting to a more comfortable and stable lower state, such as the common -1 seen in other halides.

The tendency of an element or compound to attract electrons towards itself is referred to as its electronegativity. In a perchlorate molecule, chlorine's high electronegativity is in full display, marking its powerful oxidizing nature. When used, perchlorate relinquishes its high-energy setup by accepting electrons, thereby facilitating redox reactions where other substances are oxidized.