The electron configuration of an atom is a systematic notation that describes how electrons are distributed among the various atomic shells and orbitals. For Nickel (\(\text{Ni}\)), which has an atomic number of 28, the configuration must account for all 28 electrons in a neutral atom. This arrangement dictates the chemical behavior of the element, particularly how it interacts with other atoms to form bonds.
The Governing Principles of Electron Arrangement
The placement of electrons follows three governing principles that ensure the lowest energy, most stable arrangement. The Aufbau principle states that electrons fill orbitals starting from the lowest available energy level before occupying higher ones. This establishes the sequence of orbital filling, such as \(1s\) before \(2s\), and \(2p\) before \(3s\).
The Pauli Exclusion Principle limits the number of electrons in any single orbital to a maximum of two, and these two electrons must have opposite spins. This rule determines the capacity of each subshell, allowing \(s\) subshells to hold two electrons, \(p\) subshells six, and \(d\) subshells ten.
When electrons encounter orbitals of the same energy, such as the three orbitals in a \(p\) subshell or the five orbitals in a \(d\) subshell, Hund’s rule applies. This rule requires that electrons occupy each degenerate orbital singly before any pairing occurs. Furthermore, all singly occupied orbitals must have electrons with the same spin, leading to the maximum number of unpaired electrons and a more stable configuration.
Step-by-Step Derivation of Nickel’s Configuration
To determine the electron configuration for neutral Nickel, with its 28 electrons, the process begins by sequentially filling the orbitals according to the Aufbau principle. The first ten electrons fill the \(1s\), \(2s\), and \(2p\) subshells (\(1s^2 2s^2 2p^6\)). The next eight electrons fill the \(3s\) and \(3p\) subshells (\(3s^2 3p^6\)), accounting for 18 electrons in total.
Following the \(3p\) subshell, the next lowest energy level is the \(4s\) subshell, which accepts the next two electrons to become \(4s^2\). The remaining eight electrons are then placed into the \(3d\) subshell, resulting in \(3d^8\). The full electron configuration for neutral Nickel is therefore \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8\).
While the electrons fill the \(4s\) orbital before the \(3d\) orbital, the convention for writing the final configuration often places the \(3d\) subshell before the \(4s\) subshell. This is done to group subshells by their principal quantum number (\(n\)), resulting in the common textbook presentation: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^8 4s^2\). Both notations accurately represent the distribution of the 28 electrons in the neutral atom.
Notations: Full and Condensed Representations
The full notation, such as \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8\), details the electron count in every subshell, but it can be lengthy for heavier elements. The core electrons, which are the inner-shell electrons that do not participate in chemical bonding, are often represented using a condensed notation. This simplified form uses the symbol of the preceding noble gas in square brackets to represent the stable, filled inner shells.
The noble gas that comes before Nickel is Argon (\(\text{Ar}\)), which has a total of 18 electrons and the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6\). The condensed configuration for Nickel replaces these first 18 electrons with \([\text{Ar}]\) in brackets. The remaining electrons are those in the outermost \(4s\) and \(3d\) subshells.
The resulting condensed notation for neutral Nickel is \([\text{Ar}] 4s^2 3d^8\). This notation is practical because it focuses attention on the valence electrons, which are involved in chemical reactions. These outermost electrons are the primary determinants of an element’s chemical properties.
Configuration of Nickel Ions and Stability Factors
Nickel, as a transition metal, commonly forms positive ions by losing electrons, with the \(\text{Ni}^{2+}\) ion being the most prevalent in chemical compounds. When an atom forms a cation, electrons are removed from the orbitals with the highest principal quantum number (\(n\)) first, regardless of the filling order. Although the \(3d\) subshell is filled after the \(4s\) subshell, the \(4s\) orbital belongs to the fourth energy level (\(n=4\)) and is considered the outermost shell.
Consequently, when Nickel ionizes, the two electrons are removed from the \(4s\) orbital first, not the \(3d\) orbital. The configuration for the \(\text{Ni}^{2+}\) ion, which has 26 electrons, is \([\text{Ar}] 3d^8\). The \(\text{Ni}^{3+}\) ion is formed by removing a third electron, which must come from the \(3d\) subshell after the \(4s\) electrons are gone.
The resulting configuration for the \(\text{Ni}^{3+}\) ion, which has 25 electrons, is \([\text{Ar}] 3d^7\). The stability of these ions is influenced by the electron arrangement in the \(d\) subshell. While half-filled (\(d^5\)) and fully-filled (\(d^{10}\)) subshells offer exceptional stability, the \(d^8\) configuration of \(\text{Ni}^{2+}\) is still a relatively stable arrangement for a transition metal ion.