Electron configuration is the systematic arrangement of electrons in the atomic orbitals, which are designated by letters like s, p, d, and f. This organizational structure is fundamental because it dictates an atom’s chemical behavior, including how it will bond with other atoms and the overall stability of the resulting compounds. For the vast majority of elements, this electron placement follows a reliable and predictable set of rules.
Understanding Standard Electron Configuration Rules
The placement of electrons in orbitals is generally governed by three principles, which together define the standard filling pattern. The Aufbau principle states that electrons will occupy the lowest energy orbitals available first before moving to higher energy ones. This principle dictates the general sequence of filling, such as 1s, 2s, 2p, 3s, 3p, and then 4s before 3d.
Another guideline is Hund’s Rule, which applies when multiple orbitals within the same subshell, such as the three p orbitals or five d orbitals, have equal energy. This rule mandates that electrons will fill each of these degenerate orbitals singly before any orbital is double-occupied with a second, paired electron. This maximizes the number of unpaired electrons with parallel spin, which is a more stable arrangement for the atom. The Pauli Exclusion Principle completes the set, stating that any orbital can hold a maximum of two electrons, and these two electrons must have opposite spins.
Identifying the Two Primary Exceptions
Two elements in the first row of the transition metals are the most commonly cited exceptions to the standard filling rules: Chromium (Cr) and Copper (Cu).
Chromium, with an atomic number of 24, would be predicted to have the configuration [Ar] 4s2 3d4 by following the Aufbau principle. The actual, observed electron configuration is [Ar] 4s1 3d5, showing a shift of one electron from the 4s orbital to the 3d orbital.
Copper, with an atomic number of 29, is predicted to have the configuration [Ar] 4s2 3d9. However, its actual, observed configuration is [Ar] 4s1 3d10. In both cases, the final, most stable configuration is achieved by promoting a single electron from the 4s orbital into the higher-energy 3d orbital. This small energy cost is overcome by a significant gain in stability for the overall atomic structure.
The Stability Principle Behind the Anomalies
The underlying reason for these anomalies is the quantum mechanical principle of enhanced stability associated with certain orbital fillings. Specifically, subshells that are exactly half-filled (d5) or completely filled (d10) possess greater stability than configurations that are only partially filled. This extra stability is a result of two primary factors: electron symmetry and exchange energy.
A half-filled or fully-filled subshell results in a more symmetrical distribution of electron charge around the nucleus, which is inherently a more stable state for a physical system. Exchange energy refers to the stabilizing energy released when two or more electrons with the same spin are present in different degenerate orbitals and can effectively swap positions. The number of possible exchanges is maximized when a subshell is exactly half-filled (d5 in Chromium) or completely filled (d10 in Copper).
The stability gained from maximizing exchange energy and achieving spherical symmetry is large enough to overcome the slight energy difference between the 4s and 3d orbitals. For Chromium, the electron is promoted to achieve the highly stable half-filled d5 configuration. For Copper, the electron moves to complete the d10 shell, which is the most stable configuration possible for the d subshell. This extra stability lowers the overall energy of the atom, making the anomalous configuration the true ground state.