Scandium (Sc) is a silvery-white transition metal and the first element in the periodic table’s d-block, possessing an atomic number of 21. The chemical behavior and physical characteristics of any element are determined by the arrangement of its electrons. Electrons are negatively charged particles that occupy specific regions around the nucleus. Their organization dictates how an atom interacts with other atoms.
Understanding how these electrons are distributed is necessary to predict an atom’s properties, such as its magnetic behavior. An “unpaired electron” refers to an electron that occupies an orbital alone, without a partner electron spinning in the opposite direction. The count of these solitary electrons provides the answer to Scandium’s specific magnetic and chemical nature.
Atomic Structure and Electron Shells
A neutral atom contains a number of electrons exactly equal to its atomic number, which corresponds to the number of protons in its nucleus. For Scandium (atomic number 21), the atom must accommodate 21 electrons. These electrons are organized into distinct energy levels called electron shells, which are numbered sequentially starting from the one closest to the nucleus.
Within each main shell, electrons reside in specific regions known as subshells, designated by the letters s, p, d, and f. The capacity of each subshell is fixed, determined by the number of orbitals it contains. The s subshell holds a maximum of two electrons, since it contains only one orbital.
The p subshell is composed of three orbitals and holds up to six electrons. The d subshell consists of five orbitals, housing a maximum of ten electrons. The f subshell accommodates up to fourteen electrons across its seven orbitals. Electrons fill these subshells in a particular order, starting with the lowest energy level available.
Mapping the Electrons: Scandium’s Configuration
Scandium’s 21 electrons are systematically placed into the subshells in increasing order of energy, following the building-up principle. The process begins with the lowest energy level, the 1s subshell, filled with the first two electrons (\(1s^2\)).
The next energy levels are the 2s and 2p subshells, accepting two and six electrons, respectively, totaling ten electrons (\(2s^2 2p^6\)). The 3s and 3p subshells are then filled, accounting for eighteen electrons in total (\(3s^2 3p^6\)). This complete configuration corresponds to the stable structure of the noble gas Argon.
The remaining three electrons must be placed into the next available subshells. The 4s subshell is filled before the 3d subshell because it has a slightly lower energy state. The next two electrons occupy the 4s subshell (\(4s^2\)), leaving only one electron left to place.
This final electron is placed into the 3d subshell, the next higher energy level. Therefore, the complete electron configuration for a neutral Scandium atom is \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^1\). This configuration, abbreviated as \([\text{Ar}] 4s^2 3d^1\), details the distribution of all 21 electrons.
Visualizing Orbitals and Counting Unpaired Electrons
The final step in determining the number of unpaired electrons requires visualizing how the electrons are arranged within the individual orbitals of the valence subshells. The full orbitals, such as the \(1s, 2s, 2p, 3s,\) and \(3p\) subshells, are already filled with paired electrons and do not contribute to the unpaired count. Attention must be paid only to the partially filled \(4s\) and \(3d\) subshells.
To correctly fill the orbitals, Hund’s rule must be applied, which dictates the distribution of electrons within a subshell. This rule states that when multiple orbitals are available at the same energy level, electrons will first occupy these orbitals singly. Only once all orbitals in that subshell have one electron will the electrons begin to pair up. This behavior is preferred because electrons repel one another, maximizing their distance.
For Scandium, the \(4s\) subshell contains one orbital occupied by two electrons (\(4s^2\)). These two electrons must be paired within the single \(s\) orbital, spinning in opposite directions. Consequently, the \(4s\) subshell contains zero unpaired electrons.
The \(3d\) subshell is where the final electron resides (\(3d^1\)). The \(d\) subshell is composed of five distinct orbitals, all at the same energy level. Since there is only one electron to place into these five available orbitals, that electron will occupy a single \(d\) orbital alone.
This solitary electron in the \(3d\) subshell is an unpaired electron. Therefore, the total number of unpaired electrons in a ground-state Scandium atom is exactly one. The presence of this single unpaired electron is responsible for Scandium exhibiting paramagnetism, where the element is weakly attracted to an external magnetic field.