Electron configuration details the specific arrangement of electrons in an atom’s energy levels and orbitals. This configuration is a powerful tool for predicting an atom’s chemical behavior and bonding tendencies. While determining the electron configuration for a neutral atom follows a straightforward set of rules, the process changes significantly when an atom gains or loses electrons to form an ion. Configuring ions requires a modified approach, particularly for positively charged ions, or cations, where the removal of electrons does not always follow the reverse of the filling order.
Foundation: Electron Configuration for Neutral Atoms
The arrangement of electrons in a neutral atom is governed by three fundamental principles. The Aufbau principle states that electrons must occupy the lowest energy orbitals available before filling higher-energy orbitals, creating the predictable sequence (\(1s, 2s, 2p, 3s\), etc.). The electron distribution is written using spectroscopic notation, listing the principal quantum number (\(n\)), the orbital type (\(s\), \(p\), \(d\), \(f\)), and the number of electrons (superscript).
The Pauli Exclusion Principle limits the occupancy of each orbital to a maximum of two electrons, and these electrons must possess opposite spins. Furthermore, Hund’s Rule addresses how electrons fill orbitals of the same energy, known as degenerate orbitals, such as the three \(p\) orbitals or five \(d\) orbitals. This rule requires that every degenerate orbital must receive one electron with parallel spin before any orbital can be doubly occupied, maximizing the number of unpaired electrons to achieve the lowest-energy, most stable configuration.
Step-by-Step for Anions: Adding Valence Electrons
Anions are negatively charged ions formed when a neutral atom gains one or more electrons, typically to achieve a stable, noble gas electron configuration. The process for determining the electron configuration of an anion is relatively simple because the newly added electrons follow the standard Aufbau filling order. The extra electrons are simply placed into the next available orbital of the parent atom’s configuration, which is usually a partially filled valence orbital.
For example, a neutral oxygen atom (O) has the electron configuration \(1s^2 2s^2 2p^4\). To become the oxide ion (O\(^{2-}\)), the atom gains two electrons. These two electrons are added to the partially filled \(2p\) subshell, which can hold a maximum of six electrons. The resulting electron configuration for the stable oxide ion is \(1s^2 2s^2 2p^6\), which is isoelectronic with the noble gas neon.
The Critical Rule for Cations: Removing Valence Electrons
Cations, or positively charged ions, are formed when a neutral atom loses one or more electrons, and this process introduces the primary complexity in ion configuration. The fundamental rule for cation formation is that electrons are always removed from the orbital with the highest principal quantum number (\(n\)) first, because these are the outermost, highest-energy electrons. This removal order takes precedence over the order in which the orbitals were originally filled.
For main group elements, the rule is straightforward, as the highest principal quantum number corresponds to the last filled orbital. For instance, neutral sodium (Na) is \(1s^2 2s^2 2p^6 3s^1\), and forming the sodium ion (Na\(^+\)) involves removing the single valence electron from the \(3s\) orbital. The resulting stable configuration is \(1s^2 2s^2 2p^6\), which mirrors the configuration of neon.
Transition Metal Cations
The rule becomes non-intuitive when dealing with transition metals, which involve \(s\) and \(d\) orbitals with closely spaced energies. For a neutral transition metal like iron (Fe), the Aufbau principle dictates the configuration as \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^6\). The \(4s\) orbital is filled before the \(3d\) orbital because it is slightly lower in energy in the neutral atom.
However, when iron forms a cation, the \(4s\) electrons are always removed before the \(3d\) electrons, even though the \(3d\) orbital was filled later. The principal quantum number of the \(4s\) orbital (\(n=4\)) is higher than the \(3d\) orbital (\(n=3\)), making the \(4s\) electrons the outermost and therefore the easiest to remove. This occurs because the \(3d\) orbitals contract and become lower in energy relative to the \(4s\) orbital once the electrons begin to fill the \(d\) subshell.
To form the iron(II) ion (Fe\(^{2+}\)), the two electrons are removed from the \(4s\) orbital first. This yields the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^6\). If a third electron is lost to form the iron(III) ion (Fe\(^{3+}\)), that electron must be removed from the \(3d\) subshell. The resulting configuration is \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5\), which is stable due to the half-filled \(3d\) subshell. This demonstrates the distinction: the order of filling applies only to neutral atoms, while the order of removal is determined by the highest principal quantum number.
Presenting the Result: Shorthand Notation and Orbital Diagrams
Once the correct electron configuration for the ion is determined, it can be presented in a concise format using the noble gas shorthand notation. This method replaces the inner core electrons with the symbol of the preceding noble gas in square brackets. For example, the full configuration of Fe\(^{3+}\) (\(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5\)) can be simplified to \([Ar] 3d^5\), since the argon atom has the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6\).
Another visual method for representing the ion’s electron arrangement is the orbital diagram. This diagram uses a series of boxes or lines to represent each specific orbital, with arrows used to represent the electrons within them. Arrows pointing up and down indicate electrons with opposite spins, adhering to the Pauli Exclusion Principle. The diagram visually confirms Hund’s Rule by showing electrons distributed singly across degenerate orbitals before any pairing occurs.