The oxidation state of oxygen in \(\mathrm{O}_{2}\left[\mathrm{PtF}_{6}\right]\) is: (a) \(-1 / 2\) (b) \(+2\) (c) \(+1 / 2\) (d) \(+1\)

Oxygen is a versatile element, known primarily for its general oxidation state of -2 when it combines with other elements. However, in the special case of peroxide compounds, we encounter what is known as a peroxide linkage. This unique feature occurs when two oxygen atoms are directly connected to each other, forming an O-O bond within a molecule.

In peroxides, unlike in most other oxygen-containing compounds, each oxygen atom has an oxidation state of -1. This deviates from the common -2 state due to the shared electrons between the two oxygen atoms in the peroxide bond. Essentially, each oxygen atom keeps one additional electron from its usual share, leading to this less negative oxidation state.

Hydrogen peroxide (H₂O₂) is a classic example of a peroxide, where the O-O peroxide linkage is clearly present. The distinctive '-1' oxidation state in peroxides is a crucial exception to remember, as it plays into further understanding of oxidation states in compounds and molecular interactions.

Understanding the calculation of oxidation states is fundamental for grasping many concepts in chemistry, especially when delving into redox reactions and the balancing of chemical equations. The oxidation state can be thought of as the hypothetical charge an atom would carry if all its bonds were ionic, with the electrons being donated to the more electronegative element.

To determine oxidation states, we follow a set of rules:

• For an atom in its elemental form, the oxidation state is always zero.
• For a monatomic ion, the oxidation state is equal to the ion's charge.
• Fluorine always has an oxidation state of -1 in its compounds.
• Hydrogen is usually +1 (except when it forms hydrides with metals, where it is -1).
• Oxygen is usually -2, but with exceptions like in peroxides.

When these rules result in multiple possible oxidation states, additional information, such as the typical charges of other elements in the compound or the overall charge of the molecule, is used to pinpoint the correct values. It is a systematic approach where the known oxidation states can help deduce the unknown ones, leading to a clear understanding of the electron distribution in a molecule.

While oxygen typically exhibits an oxidation state of -2, there are notable exceptions that every chemistry student should recognize. These include:

• In peroxides such as H₂O₂, oxygen has an oxidation state of -1 due to the peroxide linkage, as discussed previously.
• In superoxides, which contain the superoxide anion (O₂^-), each oxygen atom has an oxidation state of -1/2.
• When bonded to fluorine, which is more electronegative, oxygen can have a positive oxidation state, as fluorine will always be -1.

Oxygen in compounds with less electronegative elements than carbon, such as OF₂, is an example where these rules come in very handy. Here, oxidation states are essential not only for balancing equations but also for understanding the underlying chemical properties and reactivities of different substances.