The magnetic behavior of an ion, such as the beryllium ion (\(\text{Be}^{2+}\)), is determined by the arrangement of its electrons within their orbitals. To classify \(\text{Be}^{2+}\) as diamagnetic or paramagnetic, we must examine its electron configuration. The presence or absence of single, unpaired electrons dictates the ion’s magnetic classification.
Understanding Diamagnetism and Paramagnetism
A material’s magnetic property is categorized based on how it interacts with an external magnetic field. When all electrons in an atom or ion are paired, the substance is classified as diamagnetic. The magnetic moment created by one spinning electron is perfectly canceled out by its partner spinning in the opposite direction, resulting in a net magnetic moment of zero.
Diamagnetic substances are consequently repelled very weakly by an external magnetic field. Conversely, a substance that contains at least one unpaired electron is called paramagnetic. An unpaired electron creates a tiny, permanent magnetic moment. These paramagnetic materials are weakly attracted to an external magnetic field because their individual magnetic moments tend to align with the applied field.
Determining the Electron Count of Beryllium Ion
The neutral beryllium atom (Be) is located in Group 2 of the periodic table and has an atomic number of 4. This atomic number indicates that a neutral beryllium atom possesses four electrons.
The species in question is the beryllium ion, which carries a \(2+\) charge (\(\text{Be}^{2+}\)). The positive charge signifies that the neutral atom has lost electrons. Specifically, a \(2+\) charge means that two electrons have been removed from the neutral atom. Therefore, the \(\text{Be}^{2+}\) ion has only two electrons remaining, which will be used to determine the ion’s electron configuration.
Electron Configuration and the Final Answer
With a confirmed count of two electrons, the electron configuration for the \(\text{Be}^{2+}\) ion can be written by following the rules of orbital filling. Electrons fill the lowest energy levels first, starting with the \(1s\) orbital. The \(1s\) orbital can hold a maximum of two electrons.
Both available electrons in the \(\text{Be}^{2+}\) ion will occupy this single \(1s\) orbital, resulting in the electron configuration \(1s^2\). According to the Pauli Exclusion Principle, two electrons sharing the same orbital must have opposite spins, meaning they are paired. Since the \(\text{Be}^{2+}\) ion has only the completely filled \(1s\) orbital and zero unpaired electrons, the \(\text{Be}^{2+}\) ion is diamagnetic.