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Electronic configuration

The general electronic configuration of the d-block atoms is (n − 1)d n s n p. Here "" is the electronic configuration of the last noble gas preceding the atom in question, and n is the highest principal quantum number of an occupied orbital in that atom. For example, Ti (Z = 22) is in period 4 so that n = 4, the first 18 electrons have the same configuration of Ar at the end of period 3, and the overall configuration is 3d 4s. The period 6 and 7 transition metals also add core (n − 2)f electrons, which are omitted from the tables below. The p orbitals are almost never filled in free atoms (the one exception being lawrencium due to relativistic effects that become important at such high Z), but they can contribute to the chemical bonding in transition metal compounds.

The Madelung rule predicts that the inner d orbital is filled after the valence-shell s orbital. The typical electronic structure of transition metal atoms is then written as n s (n − 1)d. This rule is approximate, but holds for most of the transition metals. Even when it fails for the neutral ground state, it accurately describes a low-lying excited state.

The d subshell is the next-to-last subshell and is denoted as (n − 1)d subshell. The number of s electrons in the outermost s subshell is generally one or two except palladium (Pd), with no electron in that s sub shell in its ground state. The s subshell in the valence shell is represented as the n s subshell, e.g. 4s. In the periodic table, the transition metals are present in ten groups (3 to 12).

The elements in group 3 have an n s (n − 1)d configuration, except for lawrencium (Lr): its 7s 7p configuration exceptionally does not fill the 6d orbitals at all. The first transition series is present in the 4th period, and starts after Ca (Z = 20) of group 2 with the configuration 4s, or scandium (Sc), the first element of group 3 with atomic number Z = 21 and configuration 4s 3d, depending on the definition used. As we move from left to right, electrons are added to the same d subshell till it is complete. Since the electrons added fill the (n − 1)d orbitals, the properties of the d-block elements are quite different from those of s and p block elements in which the filling occurs either in s or in p orbitals of the valence shell.The electronic configuration of the individual elements present in all the d-block series are given below:

A careful look at the electronic configuration of the elements reveals that there are certain exceptions to the Madelung rule. For Cr as an example the rule predicts the configuration 3d 4s, but the observed atomic spectra show that the real ground state is 3d 4s. To explain such exceptions, it is necessary to consider the effects of increasing nuclear charge on the orbital energies, as well as the electron–electron interactions including both Coulomb repulsion and exchange energy. The exceptions are in any case not very relevant for chemistry because the energy difference between them and the expected configuration is always quite low.

The (n − 1)d orbitals that are involved in the transition metals are very significant because they influence such properties as magnetic character, variable oxidation states, formation of coloured compounds etc. The valence s and p orbitals (n s and n p) have very little contribution in this regard since they hardly change in the moving from left to the right in a transition series.In transition metals, there are greater horizontal similarities in the properties of the elements in a period in comparison to the periods in which the d orbitals are not involved. This is because in a transition series, the valence shell electronic configuration of the elements do not change. However, there are some group similarities as well.