The common oxygen molecule, O2, is a diatomic gas that plays a central role in atmospheric and biological processes. While many stable molecules exhibit no interaction with a magnetic field, oxygen presents a surprising exception. The simple answer to the question of its magnetic nature is that O2 is paramagnetic. This designation means the invisible gas possesses an inherent magnetic character, causing it to be weakly attracted to an external magnetic field. This unusual property is a consequence of the specific arrangement of electrons within the molecule’s structure.
Defining Paramagnetism and Diamagnetism
The magnetic behavior of any substance is fundamentally determined by the spin and arrangement of its electrons. Materials are generally categorized into two main groups based on how they respond to a magnetic field. Paramagnetism describes a weak attraction to an external magnetic field that disappears when the field is removed. This attraction arises from the presence of one or more unpaired electrons within the material’s atomic or molecular structure.
Conversely, diamagnetism is a property found in all substances, though it is only noticeable when paramagnetism is absent. Diamagnetic materials are weakly repelled by a magnetic field because all of their electrons are fully paired up in their orbitals. In a paired electron configuration, the magnetic moments of the two electrons, which spin in opposite directions, cancel each other out completely. This leaves no net magnetic moment to align with an external field, resulting in the slight repulsion.
The Role of Unpaired Electrons in Oxygen
The paramagnetism of the oxygen molecule is a direct result of its electron configuration, which features two unpaired electrons. For many years, traditional chemical bonding theories, such as Valence Bond Theory, incorrectly predicted that the O2 molecule would be diamagnetic. These early models suggested that all the valence electrons would be neatly paired, thus failing to account for oxygen’s experimentally observed magnetic attraction. The failure of these simpler theories highlighted a significant gap in the understanding of how electrons behave in molecules.
These two unpaired electrons give oxygen its unique magnetic signature. When an external magnetic field is applied, the magnetic moments generated by these lone electrons align themselves with the field. This alignment results in a net magnetic moment for the entire molecule, causing the weak attraction characteristic of a paramagnetic substance. This molecular-level magnetic property is present regardless of whether the oxygen is in its gaseous, liquid, or solid state.
Molecular Structure: The O2 Orbital Diagram
Understanding oxygen’s paramagnetism requires the application of Molecular Orbital (MO) Theory, a more sophisticated model of chemical bonding. Unlike Valence Bond Theory, which focuses on localized electron pairs, MO Theory treats the entire molecule as a single entity with its own set of energy levels, known as molecular orbitals. These orbitals are formed by the combination of the atomic orbitals from the two oxygen atoms, creating both lower-energy bonding orbitals and higher-energy antibonding orbitals.
For the O2 molecule, which contains a total of twelve valence electrons (six from each oxygen atom), these electrons are systematically filled into the newly formed molecular orbitals starting from the lowest energy level. The filling sequence involves orbitals designated as sigma (\(\sigma\)) and pi (\(\pi\)), with corresponding antibonding orbitals denoted by an asterisk (\(\sigma^\) and \(\pi^\)). The first ten valence electrons fill the \(\sigma_{2s}\), \(\sigma_{2s}^\), \(\sigma_{2p}\), and the two \(\pi_{2p}\) bonding orbitals.
The crucial difference for oxygen arises when placing the final two valence electrons. These two electrons must occupy the next available energy level, which consists of the two \(\pi_{2p}^\) antibonding orbitals. According to Hund’s Rule, electrons will occupy separate orbitals singly before they pair up. This rule ensures the electrons maximize their total spin.
The eleventh and twelfth valence electrons each enter a different \(\pi_{2p}^\) orbital, and they both possess the same spin orientation. This molecular orbital configuration, with its two unpaired electrons in the antibonding orbitals, explains why the oxygen molecule is paramagnetic.
Observing Oxygen’s Magnetic Properties
Oxygen’s paramagnetism, predicted by Molecular Orbital Theory, is confirmed through laboratory experiments. The most common demonstration involves cooling oxygen gas until it liquefies at approximately \(-183^\circ\text{C}\) (\(-297^\circ\text{F}\)). When liquid oxygen is poured between the pole faces of a strong magnet, it does not simply flow through like most other liquids.
Instead, the liquid oxygen is attracted to the magnetic field and is held suspended between the magnetic poles until it evaporates. In contrast, liquid nitrogen, which has all paired electrons, flows straight through the magnetic field without attraction, confirming its diamagnetic nature. This physical suspension of liquid oxygen serves as a real-world example of the force exerted by the magnetic moment of its two unpaired electrons.
Beyond the classroom demonstration, oxygen’s magnetic property has practical implications in technology and medicine. For instance, the behavior of oxygen in magnetic fields is relevant in processes like the separation of atmospheric gases for industrial use. Furthermore, the principles of magnetism, linked to electron spin, form the basis for Magnetic Resonance Imaging (MRI), a technique that uses powerful magnetic fields to create detailed images of the human body.