The interaction of matter with a magnetic field is a fundamental property determined by the behavior and arrangement of electrons within its atoms or ions. When a substance is placed near a magnet, the resulting attraction or repulsion reveals its magnetic character. This magnetic response is dictated by specific quantum mechanical rules governing how electrons occupy their energy levels. To determine if the Cobalt(II) ion (\(\text{Co}^{2+}\)) is attracted to or repelled by a magnetic field, we must examine its precise electron configuration. The magnetic state rests entirely on whether the ion possesses unpaired electrons in its atomic orbitals.
Understanding Paramagnetism and Diamagnetism
The magnetic behavior of any substance is classified into two types: diamagnetism and paramagnetism. Diamagnetism is characterized by a weak repulsion from an external magnetic field. This occurs because all electrons in a diamagnetic substance are paired, meaning the spin of each electron is perfectly canceled out by its partner. This results in a net magnetic moment of zero for the atom or ion.
Paramagnetism is a state where the material is weakly attracted to an external magnetic field. This attraction is caused by the presence of one or more unpaired electrons within the atom or ion. An unpaired electron creates a small, permanent magnetic moment that aligns with the applied external field. The external magnetism is strong enough to induce a temporary, weak alignment, causing the attraction.
Electron Configuration of the Cobalt(II) Ion
The first step in determining the magnetic state is to establish the precise electron configuration of the Cobalt(II) ion. The neutral Cobalt atom (Co) has an atomic number of 27, meaning it possesses 27 protons and 27 electrons. Its ground-state electron configuration is \([\text{Ar}]4s^2 3d^7\). This configuration shows that the outermost electrons are found in the \(4s\) and \(3d\) subshells.
The Cobalt(II) ion (\(\text{Co}^{2+}\)) is formed when the neutral atom loses two electrons. When ionizing transition metals, electrons are always removed from the orbital with the highest principal quantum number (\(n\)) first. This is because these are the valence electrons furthest from the nucleus.
In the case of Cobalt, the \(4s\) orbital has a principal quantum number of \(n=4\), which is higher than the \(3d\) orbital’s \(n=3\). Consequently, the two electrons are removed from the \(4s\) orbital. Removing the two \(4s\) electrons leaves the \(\text{Co}^{2+}\) ion with the final electron configuration of \([\text{Ar}]3d^7\). This configuration, with seven electrons residing in the \(3d\) subshell, dictates the ion’s magnetic properties.
Determining the Magnetic State of \(\text{Co}^{2+}\)
The \(3d^7\) configuration means seven electrons are distributed among the five degenerate \(d\) orbitals. Two fundamental quantum mechanical rules govern this arrangement. The Pauli Exclusion Principle states that a maximum of two electrons can occupy any single orbital, and those two electrons must have opposite spins.
Hund’s Rule dictates the filling order within orbitals of the same energy. It asserts that electrons will occupy each orbital singly before any pairing occurs, and all singly occupied orbitals will have parallel spins. This arrangement minimizes electron-electron repulsion and maximizes the number of unpaired electrons.
Applying these rules to the seven \(d\) electrons, the first five electrons occupy the five \(d\) orbitals individually, all with parallel spins. The remaining two electrons must then pair up in two of the half-filled orbitals. This leaves three of the five \(d\) orbitals with only a single electron. Because the Cobalt(II) ion has three unpaired electrons, it possesses a net magnetic moment and is classified as paramagnetic. This means \(\text{Co}^{2+}\) is weakly attracted to an external magnetic field.