The interaction of matter with an external magnetic field reveals fundamental properties of its atomic structure. This behavior is categorized into two primary types: paramagnetism and diamagnetism, determined by the arrangement of electrons within an atom or ion. Understanding whether a substance is attracted to or repelled by a magnet provides insight into its subatomic configuration. This analysis focuses on the iron(III) ion, \(\text{Fe}^{3+}\), to determine its magnetic classification.
Understanding Paramagnetism and Diamagnetism
The magnetic behavior of any substance is directly dependent on the spin of its constituent electrons. Electrons possess an intrinsic property called spin, which creates a tiny magnetic moment, effectively making each electron a miniature magnet. When electrons are paired within an atomic orbital, they must have opposite spins, a concept called the Pauli Exclusion Principle. The opposing spins cancel out their magnetic moments, resulting in a net magnetic spin of zero for that pair.
A substance is classified as diamagnetic when all of its electrons are paired, meaning there is no net magnetic moment. Diamagnetic materials exhibit a very slight repulsion when placed in an external magnetic field. Conversely, an atom or ion is considered paramagnetic if it possesses one or more unpaired electrons. These single electrons, with their uncancelled spin, produce a net magnetic moment for the atom.
The presence of unpaired electrons causes the material to be weakly attracted to an external magnetic field. The strength of this attraction, or the degree of paramagnetism, increases with the number of unpaired electrons. Determining the magnetic property of an ion requires mapping out its electron configuration to count any unpaired electrons.
Electron Configuration of the Neutral Iron Atom
Iron (\(\text{Fe}\)) is a transition metal with an atomic number of 26, meaning a neutral iron atom contains 26 electrons. To establish the groundwork for the \(\text{Fe}^{3+}\) ion, the electron arrangement of the neutral atom must first be determined. Electrons fill the lowest energy levels first, following the Aufbau principle. For neutral iron, the core electrons match the configuration of the noble gas Argon (\(\text{Ar}\)), which accounts for 18 electrons.
The remaining eight electrons occupy the valence shell orbitals. The order of filling dictates that the \(4s\) orbital, with its lower energy, fills before the \(3d\) orbital. This results in a ground-state electron configuration of \([\text{Ar}] 4s^2 3d^6\) for the neutral iron atom. This configuration shows two electrons in the \(4s\) orbital and six electrons distributed across the five \(3d\) orbitals.
Determining the Electron Configuration of the \(\text{Fe}^{3+}\) Ion
The formation of a cation, a positively charged ion, involves the loss of electrons from the neutral atom. The \(\text{Fe}^{3+}\) ion is created when a neutral iron atom loses three electrons. A specific rule governs the removal of electrons in transition metals like iron: electrons are always removed from the orbital with the highest principal quantum number (\(n\)) first, which represents the outermost shell.
In the case of iron, the \(4s\) orbital has a principal quantum number of \(n=4\), while the \(3d\) orbital has \(n=3\). Therefore, the two electrons in the \(4s\) orbital are removed first, despite the \(3d\) orbital being filled later. This leaves the ion with a \(+2\) charge and the configuration \([\text{Ar}] 3d^6\). To achieve the \(\text{Fe}^{3+}\) ion, one more electron must be removed, and this third electron comes from the \(3d\) orbital.
The resulting electron configuration for the \(\text{Fe}^{3+}\) ion is \([\text{Ar}] 3d^5\). The \(d\) subshell contains five orbitals, and according to Hund’s Rule, electrons fill each orbital singly before pairing. Since a half-filled subshell is energetically favorable, the five electrons in the \(3d\) subshell each occupy their own orbital. Consequently, all five electrons are unpaired.
The Magnetic Property of \(\text{Fe}^{3+}\)
The electron configuration of \([\text{Ar}] 3d^5\) for the iron(III) ion provides the answer to its magnetic classification. With five electrons distributed singly across the five \(d\) orbitals, \(\text{Fe}^{3+}\) possesses five unpaired electrons.
Since the \(\text{Fe}^{3+}\) ion contains five unpaired electrons, it is strongly paramagnetic. These uncancelled electron spins align themselves with an external magnetic field, causing the ion to be attracted to the magnet.
The magnetic property of the iron ion is a direct consequence of its oxidation state, which dictates the resulting orbital arrangement. The half-filled \(3d^5\) subshell of \(\text{Fe}^{3+}\) maximizes its magnetic moment, classifying it as a paramagnetic species. This strong attraction is a measurable physical property utilized in various chemical and material science applications.