Electronic Configuration, Oxidation States — Definition
Definition
Imagine an atom as a miniature solar system, where the nucleus is the sun and electrons are planets orbiting it. However, unlike planets, electrons don't orbit randomly; they occupy specific energy levels and regions of space called orbitals.
The way these electrons are arranged around the nucleus is called the electronic configuration. It's like an address for every electron in an atom. We follow certain rules to determine this arrangement: the Aufbau principle (electrons fill lower energy orbitals first), Pauli's exclusion principle (no two electrons can have the same set of four quantum numbers, meaning an orbital can hold a maximum of two electrons with opposite spins), and Hund's rule of maximum multiplicity (electrons fill degenerate orbitals singly before pairing up).
For example, the electronic configuration tells us how many electrons are in the outermost shell, known as the valence shell. These valence electrons are the ones involved in chemical bonding and largely dictate an element's chemical behavior.
Now, let's talk about oxidation states. When atoms form chemical bonds, they either gain, lose, or share electrons. The oxidation state (or oxidation number) is a way to keep track of these electron changes.
It's a hypothetical charge an atom would have if all its bonds were completely ionic. For instance, if an atom loses electrons, its oxidation state becomes positive (e.g., +1, +2), indicating it has been oxidized.
If it gains electrons, its oxidation state becomes negative (e.g., -1, -2), indicating it has been reduced. In covalent compounds, where electrons are shared, the oxidation state is assigned by assuming the more electronegative atom attracts all shared electrons.
For Group 16 elements, also known as the chalcogens, their valence electronic configuration is . This means they have six valence electrons. To achieve a stable octet (like noble gases), they typically try to gain two electrons, leading to a common oxidation state of -2.
However, due to varying electronegativity and the availability of d-orbitals in heavier elements, they can exhibit a range of other oxidation states, both negative and positive, which we will explore in detail.