A neutral Nickel atom (Ni), with an atomic number Z=28, possesses twenty-eight electrons in its ground state. These electrons are organized into specific regions of space around the nucleus known as atomic orbitals. An atomic orbital is a mathematical description of the probability of finding an electron within a particular area. To determine exactly how many of these orbitals are utilized by the Nickel atom, we must first understand the fundamental rules governing orbital structure and electron placement.
The Foundation: Defining Atomic Orbitals and Shells
Atomic orbitals are grouped into subshells, which are organized into principal energy levels, or shells, denoted by the principal quantum number \(n\). The first shell is \(n=1\), the second is \(n=2\), and so on, with higher numbers indicating greater distance from the nucleus and higher energy. Within these shells, there are four primary types of subshells, identified by the letters \(s\), \(p\), \(d\), and \(f\).
Each orbital, regardless of its type, can hold a maximum of two electrons. The \(s\) subshell always contains one orbital (2 electrons total capacity). The \(p\) subshell consists of three orbitals (6 electrons). The \(d\) subshell is composed of five orbitals (10 electrons), while the \(f\) subshell contains seven orbitals (14 electrons). The availability of these orbital types expands with the principal quantum number: \(n=1\) has only \(s\), \(n=2\) has \(s\) and \(p\), \(n=3\) has \(s\), \(p\), and \(d\), and \(n=4\) and higher have all four types.
Electron Configuration and Quantum Numbers for Nickel
The arrangement of Nickel’s twenty-eight electrons is governed by the Aufbau principle, which dictates that electrons fill the lowest energy orbitals first. The Pauli exclusion principle limits each orbital to a maximum of two electrons with opposite spins.
For Nickel, the electron filling proceeds systematically through the lower energy levels to achieve its ground state configuration, filling the \(1s\), \(2s\), \(2p\), \(3s\), and \(3p\) subshells completely. The \(2p\) and \(3p\) subshells each contain three orbitals, which are filled according to Hund’s rule.
The filling sequence then continues to the \(4s\) subshell before moving to the \(3d\) subshell. The complete ground-state electron configuration for a neutral Nickel atom is \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8\).
The Final Count: Determining Nickel’s Occupied Orbitals
To find the total number of occupied orbitals in Nickel, we sum the number of orbitals in each subshell that holds at least one electron. We use the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8\) to identify these subshells.
The occupied subshells and their orbital counts are:
- 1s: 1 orbital (full)
- 2s: 1 orbital (full)
- 2p: 3 orbitals (full)
- 3s: 1 orbital (full)
- 3p: 3 orbitals (full)
- 4s: 1 orbital (full)
- 3d: 5 orbitals (contains \(3d^8\), meaning all five orbitals are utilized)
Summing these counts (\(1 + 1 + 3 + 1 + 3 + 1 + 5\)) yields a total of 15 occupied or partially occupied atomic orbitals in a neutral Nickel atom in its ground state.
Why Nickel’s Filling Order Matters (Transition Metal Context)
Nickel is a transition metal, characterized by the \(4s\) orbital filling before the \(3d\) orbital, following the Aufbau principle. Initially, the \(4s\) orbital has a slightly lower energy than the \(3d\) orbital. The proximity of the \(4s\) orbital to the nucleus allows for this lower initial energy state.
However, once the \(3d\) subshell begins to fill, the relative energies shift due to electron-electron repulsion and shielding effects. This causes the energy of the \(3d\) orbitals to drop below the \(4s\) orbital, which dictates the order of electron loss during ionization.
When Nickel forms ions, such as \(\text{Ni}^{2+}\), electrons are removed from the orbital with the highest principal quantum number first, which is the \(4s\) orbital. Therefore, two electrons are lost from \(4s^2\), resulting in the \(\text{Ni}^{2+}\) ion having the configuration \(3d^8\). This behavior is a defining characteristic of transition metals.