How does the electron configuration of transition metals differ from main group elements?

Transition metals differ from main group elements in their electron configuration by filling their d-orbitals after filling their s-orbitals.

In more detail, the electron configuration of an atom describes the distribution of electrons in its atomic orbitals. For main group elements, also known as representative elements, the electron configuration follows a straightforward pattern. They fill their s and p orbitals in a systematic way, moving from one energy level to the next higher one. For example, the electron configuration of oxygen (a main group element) is 1s² 2s² 2p⁴, which follows the order of increasing energy levels.

However, transition metals, which are found in the d-block of the periodic table, have a slightly different pattern. They start filling their d-orbitals after their s-orbitals have been filled. This is because the energy of the 4s orbital is lower than the 3d orbital, so it gets filled first. But when it comes to ionisation, the 4s electrons are lost first before the 3d electrons. This is why the electron configuration of transition metals often ends in d and s orbitals, such as iron (Fe) with an electron configuration of [Ar] 4s² 3d⁶.

Moreover, transition metals can have multiple oxidation states. This is due to the fact that both the s and d orbitals are close in energy, so it's relatively easy for electrons to be removed from both orbitals. This is different from main group elements, which typically have a single oxidation state.

In summary, while main group elements fill their s and p orbitals in a systematic way, transition metals fill their s-orbitals before starting to fill their d-orbitals. This unique electron configuration of transition metals allows them to have multiple oxidation states, distinguishing them from main group elements.

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