Transition metals exhibit distinct differences from main group elements in their electron configurations, primarily characterized by the filling of their d-orbitals after their s-orbitals.

To elaborate, the electron configuration of an atom details how electrons are distributed across its atomic orbitals. Main group elements, often referred to as representative elements, follow a clear and systematic pattern in their electron configurations, filling their s and p orbitals in a sequential manner as they progress from one energy level to the next. For instance, the electron configuration of oxygen, a main group element, is given by:

This configuration adheres to the ascending order of energy levels.

In contrast, transition metals, located in the d-block of the periodic table, exhibit a different electron filling sequence. They begin to fill their d-orbitals only after their s-orbitals have been completely filled. This occurs because the energy of the $4s$ orbital is lower than that of the $3d$ orbital, leading to the $4s$ orbital being filled first. However, during the process of ionization, the electrons in the $4s$ orbital are removed before those in the $3d$ orbital. For example, the electron configuration of iron (Fe) is represented as:

Moreover, a notable feature of transition metals is their ability to exhibit multiple oxidation states. This capability arises from the close energy levels of the s and d orbitals, making it relatively easy for electrons to be removed from either orbital. This is in stark contrast to main group elements, which typically display only a single oxidation state.

In summary, while main group elements fill their s and p orbitals in a systematic order, transition metals first fill their s-orbital before commencing with their d-orbitals. This unique electron configuration not only distinguishes transition metals from main group elements but also provides them with the ability to exhibit multiple oxidation states.