Transition metals are characterized by their ability to exhibit multiple oxidation states, a property that arises from the presence of unpaired electrons in their d orbitals.

These elements occupy a central position in the periodic table and are distinguished by their capacity to form compounds with varying oxidation states. This variation is largely attributed to their electron configuration. Transition metals possess partially filled d orbitals, which can accommodate up to ten electrons. Notably, the energy levels of these d orbitals are very similar to those of the outermost s orbital, which can hold a maximum of two electrons.

When a transition metal forms a compound, it can lose electrons from both the s and d orbitals. The specific number of electrons lost determines the oxidation state of the metal. Because both the s and d orbitals provide multiple electrons for loss, transition metals are capable of exhibiting a wide range of oxidation states. For instance, iron can exist in the +2 oxidation state when it loses two electrons or in the +3 oxidation state when it loses three electrons.

Additionally, the energy difference between the 4s and 3d orbitals is relatively small, allowing both energy levels to participate in bond formation. This flexibility enables transition metals to form a diverse array of compounds with varying properties. It is important to note that while the ability to exist in multiple oxidation states is not exclusive to transition metals, they are particularly remarkable for the range and diversity of oxidation states they can display.

In summary, the distinctive electron configuration of transition metals, characterized by their partially filled d orbitals and the close energy levels of the s and d orbitals, allows them to lose different numbers of electrons, leading to the exhibition of multiple oxidation states. This property is fundamental to their chemistry and significantly contributes to their versatility in forming a variety of compounds.