Oxygen exhibits only -2 oxidation state while other members of this family show + 4 and +6 oxidation state as well. Explain why is it so?

The behavior of oxygen in terms of its oxidation states is quite fascinating and can be attributed to its unique electronic configuration and position in the periodic table. Oxygen is in Group 16, also known as the chalcogens, and while it primarily exhibits a -2 oxidation state, other elements in this group, such as sulfur, selenium, and tellurium, can show higher oxidation states of +4 and +6. Let’s delve into the reasons behind this difference.

Electronic Configuration and Stability

Oxygen has the electronic configuration of 1s² 2s² 2p⁴. This means it has six valence electrons and needs two more to achieve a stable octet configuration. As a result, it readily gains two electrons, leading to its common -2 oxidation state. This state is particularly stable for oxygen because it allows the atom to fill its outer shell completely.

Comparative Analysis with Other Chalcogens

Now, let’s look at the other members of the chalcogen family:

• Sulfur (S): With an electronic configuration of [Ne]3s²3p⁴, sulfur can also gain two electrons to achieve a -2 state. However, it can also lose electrons to form +4 and +6 oxidation states. This is because sulfur has more available orbitals (3d) that can participate in bonding, allowing it to expand its valence shell.
• Selenium (Se) and Tellurium (Te): These elements follow a similar pattern. They can also achieve higher oxidation states due to their larger atomic size and the availability of d orbitals, which facilitate the formation of multiple bonds with other elements.

Bonding and Oxidation States

The ability of sulfur, selenium, and tellurium to exhibit +4 and +6 oxidation states is largely due to their capacity to form covalent bonds with more than two atoms. For example, in sulfuric acid (H₂SO₄), sulfur is in the +6 oxidation state, bonded to four oxygen atoms. This is possible because sulfur can utilize its d orbitals to accommodate more electrons, something that oxygen cannot do effectively due to its smaller size and lack of available d orbitals.

Summary of Key Points

• Oxygen primarily exhibits a -2 oxidation state due to its need for two additional electrons to achieve a stable octet.
• Other chalcogens can exhibit +4 and +6 states because they have larger atomic sizes and available d orbitals that allow for expanded valence shell configurations.
• The ability to form multiple bonds with other elements is a crucial factor in the oxidation states of sulfur, selenium, and tellurium.

In essence, the unique properties of oxygen, including its electronic configuration and the absence of available d orbitals, limit its oxidation states to -2, while its heavier counterparts can explore a broader range of oxidation states due to their larger size and bonding capabilities. This distinction is a great example of how atomic structure influences chemical behavior.