Electron configuration describes the arrangement of electrons within an atom’s orbitals. While general principles dictate how electrons typically fill these orbitals, some elements deviate from these rules. These exceptions reveal subtleties in electron behavior.
Understanding Standard Electron Configuration
Standard electron configuration is guided by fundamental principles. The Aufbau principle states that electrons first occupy the lowest energy orbitals available, such as the 4s orbital before the 3d orbital. Hund’s rule dictates that electrons will individually occupy each orbital within a subshell before any orbital is doubly occupied, and all electrons in singly occupied orbitals will have parallel spins. The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers, meaning each orbital can hold a maximum of two electrons with opposite spins. These rules provide a framework for predicting electron configurations for most elements.
The Driving Force Behind Exceptions
Exceptions to standard electron configurations arise from the enhanced stability of half-filled or fully-filled d and f subshells, which results in a lower overall energy for the atom. The energy difference between certain orbitals, such as the 4s and 3d orbitals, is relatively small. This close energy spacing allows an electron to transfer from an s-orbital to a d-orbital if doing so leads to a more stable half-filled (e.g., d⁵) or fully-filled (e.g., d¹⁰) configuration.
This stability is attributed to two main factors: symmetrical electron distribution and exchange energy. A symmetrical arrangement of electrons in half-filled or fully-filled subshells reduces electron-electron repulsion, making the atom more stable. Electrons with parallel spins in degenerate orbitals can exchange positions, and this exchange releases energy. The more possible exchanges, the greater the stability, which is maximized in half-filled and fully-filled subshells. Electron shielding, where inner electrons reduce the effective nuclear charge felt by outer electrons, also influences orbital energies and contributes to these stability preferences.
Key Elements That Deviate
Chromium (Cr), atomic number 24, is an exception. Its expected configuration is [Ar] 4s² 3d⁴. One 4s electron promotes to the 3d orbital to achieve a stable half-filled 3d subshell, resulting in [Ar] 4s¹ 3d⁵.
Copper (Cu), atomic number 29, is another exception. Its expected configuration is [Ar] 4s² 3d⁹. Instead, copper adopts [Ar] 4s¹ 3d¹⁰, as a 4s electron moves to the 3d orbital to achieve a completely filled 3d subshell. This fully-filled d-orbital configuration is energetically more favorable.
Other Common Deviations
Beyond chromium and copper, other elements in the same groups often display similar exceptions.
Molybdenum (Mo)
Molybdenum (Mo), located directly below chromium in the periodic table, follows a similar pattern. Its configuration is [Kr] 5s¹ 4d⁵, rather than the expected [Kr] 5s² 4d⁴.
Silver (Ag) and Gold (Au)
Silver (Ag), found below copper, has a configuration of [Kr] 5s¹ 4d¹⁰ instead of the expected [Kr] 5s² 4d⁹. Gold (Au) also exhibits this behavior, with a configuration of [Xe] 4f¹⁴ 5d¹⁰ 6s¹ rather than [Xe] 4f¹⁴ 5d⁹ 6s².
Heavier Elements
Heavier elements and those in the lanthanide and actinide series can also present more intricate exceptions. Niobium (Nb), for instance, has an actual configuration of [Kr] 5s¹ 4d⁴, deviating from an expected [Kr] 5s² 4d³. This is influenced by the close energy levels of the 4d and 5s orbitals and factors like orbital size and electron repulsion. Palladium (Pd) is another exception, with a configuration of [Kr] 4d¹⁰ and no electrons in its 5s orbital. These diverse examples show that while general rules exist, the pursuit of atomic stability can lead to varied electron arrangements.