The arrangement of electrons within an atom’s energy levels and sublevels determines its chemical behavior. To predict this arrangement, known as the electron configuration, chemists rely on the Aufbau principle, which dictates a systematic order for filling atomic orbitals. This principle is highly effective for most elements, providing a stable, low-energy configuration. However, the complex interactions between electrons in heavier elements cause a handful of atoms to deviate from this expected sequence.
Understanding the Standard Electron Filling Rules
The Aufbau principle states that electrons will always occupy the lowest available energy orbital before moving into higher energy orbitals. The filling order is not strictly numerical; for instance, the 4s orbital, part of the fourth principal energy level, is filled with electrons before the 3d orbitals, which belong to the third level. This sequence is followed because the 4s orbital, due to its shape and penetration near the nucleus, is actually slightly lower in energy than the 3d orbitals in neutral atoms. The standard electron filling process also incorporates the Pauli exclusion principle, which limits any single orbital to a maximum of two electrons with opposite spins. Furthermore, Hund’s rule is applied within a subshell, requiring electrons to occupy degenerate orbitals singly before any orbital is filled with a second electron.
Identifying the Core Exceptions
The most widely cited exceptions to the expected configuration occur in the first row of the transition metals. Chromium (Cr, Z=24) is predicted to have the configuration [Ar]4s\(^2\) 3d\(^4\). However, its observed, lower-energy configuration is [Ar]4s\(^1\) 3d\(^5\). This shift involves promoting one electron from the 4s orbital into the 3d subshell.
Copper (Cu, Z=29) exhibits a similar deviation. The expected configuration would be [Ar]4s\(^2\) 3d\(^9\), but the actual configuration is [Ar]4s\(^1\) 3d\(^{10}\). These two elements are the standard examples taught in chemistry, but the phenomenon is not limited to them. Molybdenum (Mo, Z=42) and Silver (Ag, Z=47) also follow this pattern to achieve d\(^5\) and d\(^{10}\) configurations, respectively.
Other exceptions are found in the third transition series, including Gold (Au, Z=79) and certain elements in the Lanthanide and Actinide series. Palladium (Pd, Z=46) is a notable special case, as it completely empties its 5s orbital, resulting in a configuration of [Kr]4d\(^{10}\) instead of the expected [Kr]5s\(^2\) 4d\(^8\).
Why Electron Configurations Deviate
The primary reason elements bypass the standard filling order is the superior stability associated with half-filled (d\(^5\)) and fully-filled (d\(^{10}\)) electron subshells. An atom gains stability by rearranging its electrons to achieve one of these configurations, even if it means moving an electron to a seemingly higher-energy orbital like 3d. This increased stability is a result of two main quantum mechanical factors: symmetry and exchange energy.
The symmetrical distribution of electrons in a half-filled or fully-filled subshell minimizes electron-electron repulsion, which lowers the overall energy of the atom. Exchange energy is the stabilizing energy released when electrons with the same spin can exchange positions within a subshell. The number of possible exchanges is maximized when the subshell is half-filled or fully-filled, providing a significant energy benefit that outweighs the small energy cost of moving an electron from the s orbital.
The energy difference between the ns and (n-1)d orbitals, such as the 4s and 3d levels, is very slight in transition metals. This minimal energy gap makes it energetically favorable for an electron to jump from the s orbital to the d orbital to achieve the highly stable d\(^5\) or d\(^{10}\) arrangement.
Patterns in the D-Block and F-Block
The majority of electron configuration exceptions are concentrated in the d-block (transition metals) and f-block (inner transition metals) of the periodic table. The exceptions in the d-block primarily involve the promotion of one electron from the outermost s orbital to the adjacent d subshell.
The general pattern is for elements that would otherwise have a d\(^4\) or d\(^9\) configuration to promote an s-electron to complete the d\(^5\) (half-filled) or d\(^{10}\) (fully-filled) subshell. This tendency is a strong indicator of where deviations from the principle are most likely to be observed.
In the f-block, the Lanthanide and Actinide series exhibit complex exceptions because the 4f and 5d orbitals have extremely similar energy levels. This overlap causes some elements to place an electron in the d orbital before the f orbital, or vice versa, creating a series of subtle but significant deviations from the standard filling rule.