The electron configuration of an atom describes the arrangement of electrons within its atomic orbitals. This organization dictates an element’s chemical behavior, including its reactivity and the types of bonds it can form. Determining this configuration involves a set of established rules that predict the most stable, lowest-energy arrangement for the electrons. Understanding this structure is fundamental to predicting the properties of all chemical elements.
Fundamentals of Electron Configuration
To denote how electrons are distributed, scientists use a notation based on the four types of orbitals: \(s, p, d,\) and \(f\). Each orbital type has a fixed capacity for electrons; the \(s\) subshell can hold up to 2 electrons, the \(p\) subshell can hold 6, the \(d\) subshell can hold 10, and the \(f\) subshell can hold 14 electrons. The number preceding the orbital letter indicates the principal energy level, or shell, such as \(1s\) or \(4s\).
The distribution follows the Aufbau principle, which states that electrons occupy the lowest available energy levels first. For example, the \(1s\) orbital is filled before the \(2s\) orbital, and the \(2p\) orbital is filled before the \(3s\) orbital. A shorthand notation uses the symbol of the preceding noble gas in square brackets, such as \([\text{Ar}]\), to represent the core electrons.
The Expected Configuration of Copper
Copper has an atomic number of 29, meaning a neutral copper atom contains 29 electrons. Following the standard Aufbau filling order, the first 18 electrons fill the inner shells up to the configuration of Argon, \([\text{Ar}]\). The remaining 11 electrons are placed in the next available orbitals, the \(4s\) and \(3d\) subshells.
The standard order dictates that the \(4s\) subshell, which is slightly lower in energy, should be filled before the \(3d\) subshell. This process assigns two electrons to \(4s\) and the remaining nine electrons to \(3d\). Therefore, the strictly predicted configuration is \([\text{Ar}] 4s^2 3d^9\).
The Actual Configuration and Stability Rationale
The definitive, stable ground state electron configuration for a neutral copper atom is an exception to the strict Aufbau principle filling order. Instead of the predicted \([\text{Ar}] 4s^2 3d^9\) arrangement, the actual configuration is \([\text{Ar}] 4s^1 3d^{10}\). This means one electron is promoted from the \(4s\) subshell to the \(3d\) subshell.
This minor shift provides a substantial boost in atomic stability due to the enhanced stability of completely filled subshells. A \(d\) subshell is most stable when it is exactly half-filled (\(d^5\)) or completely filled (\(d^{10}\)). The energy gained by achieving a completely filled \(3d^{10}\) subshell is greater than the energy required to move one electron from the \(4s\) orbital.
Two main factors contribute to this increased stability: symmetry and exchange energy. A fully filled \(d^{10}\) subshell has a highly symmetrical distribution of electrons, which is favored in nature. Furthermore, the \(d^{10}\) configuration allows for a greater number of electron exchanges, where electrons with the same spin can switch positions, and this higher total exchange energy makes the atom significantly more stable than the \(3d^9\) arrangement.
Configuration of Copper Ions
Copper commonly forms two stable ions: copper(I), \(\text{Cu}^+\), and copper(II), \(\text{Cu}^{2+}\). When a transition metal atom ionizes, electrons are removed first from the orbital with the highest principal quantum number (\(n\)), which is the outermost shell. For copper’s configuration of \([\text{Ar}] 4s^1 3d^{10}\), the \(4s\) electron is removed first because \(4s\) has a higher \(n\) value than \(3d\).
The monovalent copper(I) ion, \(\text{Cu}^+\), is formed by the loss of the single \(4s\) electron, resulting in the configuration \([\text{Ar}] 3d^{10}\). The divalent copper(II) ion, \(\text{Cu}^{2+}\), is formed by the loss of a second electron, which is removed from the \(3d\) subshell. The final configuration for the \(\text{Cu}^{2+}\) ion is \([\text{Ar}] 3d^9\).