# Electron Configurations of d-block Elements The electronic configuration of transition metals generally follows the *Aufbau principle*. However, there are notable exceptions to this rule, particularly for chromium (Cr) and copper (Cu). According to the IUPAC definition, transition elements are those that form at least one stable ion with a partially filled d-subshell. ## Exceptions to Aufbau Principle **Chromium (Cr):** The expected configuration would be $[Ar] 3d^4 4s^2$. However, its actual configuration is $[Ar] 3d^5 4s^1$. **Copper (Cu):** The expected configuration would be $[Ar] 3d^9 4s^2$. Its actual configuration is $[Ar] 3d^{10} 4s^1$. ### Reason for Stability These exceptions occur because atoms or ions with a half-filled $3d$ sub-level ($3d^5$) or a fully-filled $3d$ sub-level ($3d^{10}$) exhibit enhanced stability. This increased stability is attributed to two main factors: 1. **Reduced Inter-electronic Repulsion:** A half-filled or fully-filled sub-level configuration minimizes the repulsive forces between electrons within the same orbital compared to configurations like $4s^2$ alongside an incomplete $3d$ sub-shell. 2. **Symmetrical Electron Cloud:** These configurations lead to a more symmetrical distribution of electron density, which provides more effective shielding of the nucleus. --- ## Forming Cations of d-block Elements A crucial rule for d-block elements when forming cations is that electrons are *always* removed from the highest energy sub-shell first. For the first-row d-block elements, this means electrons are removed from the $4s$ orbital before any electrons are removed from the $3d$ orbital. The electron configurations for some first-row d-block elements and their common ions are provided in the table below: | Element | Electronic configuration | Ion | Electronic configuration | | :--- | :--- | :--- | :--- | | Sc | $[Ar] 3d^1 4s^2$ | $Sc^{3+}$ | $[Ar]$ | | Ti | $[Ar] 3d^2 4s^2$ | $Ti^{2+}$ | $[Ar] 3d^2$ | | V | $[Ar] 3d^3 4s^2$ | $V^{2+}$ | $[Ar] 3d^3$ | | Cr | $[Ar] 3d^5 4s^1$ | $Cr^{2+}$ | $[Ar] 3d^4$ | | | | $Cr^{3+}$ | $[Ar] 3d^3$ | | Mn | $[Ar] 3d^5 4s^2$ | $Mn^{2+}$ | $[Ar] 3d^5$ | | Fe | $[Ar] 3d^6 4s^2$ | $Fe^{2+}$ | $[Ar] 3d^6$ | | | | $Fe^{3+}$ | $[Ar] 3d^5$ | | Co | $[Ar] 3d^7 4s^2$ | $Co^{2+}$ | $[Ar] 3d^7$ | | Ni | $[Ar] 3d^8 4s^2$ | $Ni^{2+}$ | $[Ar] 3d^8$ | | Cu | $[Ar] 3d^{10} 4s^1$ | $Cu^{1+}$ | $[Ar] 3d^{10}$ | | | | $Cu^{2+}$ | $[Ar] 3d^9$ | | Zn | $[Ar] 3d^{10} 4s^2$ | $Zn^{2+}$ | $[Ar] 3d^{10}$ | For more information on the properties of these transition elements, see . --- ## Relative Energies of 4s and 3d Sub-shells The energy difference between the $3d$ and $4s$ orbitals is not constant across the period of transition metals. ### In Calcium (Ca) The $4s$ orbital has a lower energy than the $3d$ orbital. This is because the $4s$ electrons experience the full attraction of the unscreened nucleus, making them more "penetrating" (closer to the nucleus on average) than $3d$ electrons. ### Across the 3d Period (after Calcium) As $3d$ electrons are added, the $4s$ electrons are increasingly shielded by these $3d$ electrons. This shielding causes a slight decrease in the energy of the $4s$ electrons, which is reflected in small rises in the first and second ionization energies across the period. Conversely, the $3d$ electrons experience a higher effective nuclear charge due to less effective shielding from other $3d$ electrons and the increasing nuclear charge. This stronger attraction by the nucleus results in a potential increase in the energy required to remove $3d$ electrons, which is typically seen as a higher rise in the *third* ionization energy compared to the second.