Oxygen is paramagnetic because its molecules contain two unpaired electrons. These electrons occupy separate molecular orbitals, and their magnetic moments make the molecule attracted to an external magnetic field. This behavior is a fundamental prediction of molecular orbital theory, which explains how electrons are distributed across an entire molecule rather than confined to individual atoms.
The result is particularly interesting because oxygen’s magnetic properties cannot be explained correctly by simply drawing its Lewis structure. The familiar structure of O₂ shows a double bond between two oxygen atoms, but it does not reveal the two unpaired electrons responsible for its magnetism. To understand why oxygen behaves this way, it is necessary to examine how atomic orbitals combine to form molecular orbitals and how electrons fill them.
What does paramagnetism mean?
Paramagnetism is a type of magnetism that occurs when a substance has unpaired electrons. Each electron has an intrinsic property called spin, which gives it a magnetic moment. An unpaired electron therefore contributes a magnetic moment that can interact with an external magnetic field.
When a paramagnetic substance is placed in a magnetic field, its magnetic moments tend to align with the field, producing a weak attraction toward the stronger-field region. Once the external field is removed, thermal motion generally randomizes the orientations of these moments, and the substance loses most of its induced magnetization.
Oxygen gas, O₂, displays this behavior because each molecule has two unpaired electrons. Its paramagnetism can be demonstrated experimentally by placing liquid oxygen between the poles of a strong magnet, where it is attracted to the magnetic field.
Not all substances respond to magnetic fields in the same way. Diamagnetic substances, for example, have no permanent net magnetic moment arising from unpaired electrons. When exposed to a magnetic field, they develop a weak induced magnetic response that opposes the applied field. Because all electrons contribute to this response, diamagnetism occurs even in substances whose electrons are completely paired.
The distinction between paramagnetic and diamagnetic behavior depends largely on electron arrangement. Oxygen is a particularly useful example because its magnetic properties provide direct evidence about how electrons occupy molecular orbitals.
Why the Lewis structure does not explain oxygen’s magnetism
A Lewis structure represents chemical bonds and valence electrons using lines and dots. For oxygen, the conventional structure is O=O, with a double bond between the two atoms and two lone pairs on each atom.
This structure is useful for understanding many aspects of oxygen’s chemistry. It represents the double bond and accounts for the total number of valence electrons. However, in its standard form, it places all electrons into pairs.
If every electron in an oxygen molecule were paired, the molecule would be expected to behave as diamagnetic rather than paramagnetic. That prediction conflicts with experiment.
The problem is not that the Lewis structure is useless. Rather, it is designed to emphasize localized bonds and electron pairs, so it does not fully describe the electronic structure of every molecule. In particular, it cannot adequately represent the unpaired electrons in the ground state of O₂ using its simplest conventional form.
Molecular orbital theory provides a more complete explanation by treating electrons as occupying orbitals that extend over the molecule as a whole.
How molecular orbitals form
When two oxygen atoms approach one another, their atomic orbitals interact. Atomic orbitals are regions described by quantum mechanics that characterize the possible states of electrons around an individual atom. When atoms form a molecule, their orbitals can combine to produce molecular orbitals.
These molecular orbitals differ from atomic orbitals because they belong to the molecule rather than to one specific atom. Electrons occupying them may be distributed across both oxygen nuclei.
Combining two atomic orbitals generally produces two molecular orbitals: a bonding orbital and an antibonding orbital.
A bonding molecular orbital is lower in energy than the atomic orbitals from which it forms. Electron density in this orbital helps hold the nuclei together, making the molecule more stable.
An antibonding molecular orbital is higher in energy. Its characteristic wave pattern includes a node between the nuclei, where the orbital wavefunction changes sign and the electron probability density is zero. Occupying an antibonding orbital tends to weaken the bond.
The labels σ (sigma) and π (pi) describe different orbital symmetries. Sigma orbitals have cylindrical symmetry around the internuclear axis, while pi orbitals have electron density distributed on either side of that axis.
For oxygen, interactions among the atomic orbitals produce a series of molecular orbitals with different energies. Electrons fill these orbitals according to the rules of quantum mechanics, and the resulting arrangement determines both the molecule’s bonding and its magnetic properties.
The molecular orbital energy levels of O₂
Each oxygen atom has eight electrons, so an O₂ molecule contains 16 electrons in total. Ten of these are core electrons, while six valence electrons come from each atom, giving the molecule 12 valence electrons to distribute among the valence molecular orbitals.
For a typical introductory treatment of O₂, the valence molecular orbitals are arranged in the following energy order:
- σ(2s)
- σ*(2s)
- σ(2p)
- π(2pₓ) and π(2pᵧ)
- π*(2pₓ) and π*(2pᵧ)
- σ*(2p)
The asterisk indicates an antibonding orbital. The two π orbitals at each of the relevant energy levels are degenerate, meaning they have the same energy in an isolated oxygen molecule.
This ordering is important. In O₂, the two π* antibonding orbitals lie below the σ*(2p) orbital. After the lower-energy orbitals are filled, the remaining electrons enter these two π* orbitals.
The full ground-state molecular orbital configuration, including the core orbitals, can be written as:
1σg² 1σu*² 2σg² 2σu*² 3σg² 1πu⁴ 1πg*²
The subscripts g and u describe how the orbital wavefunction behaves under inversion through the molecule’s center. They are useful labels for a molecule with the symmetry of O₂ but are not necessary for understanding its paramagnetism.
The key part of the configuration is the final 1πg*² term. It indicates that two electrons occupy the pair of degenerate π* antibonding orbitals.
Why the two electrons remain unpaired
The explanation follows from two fundamental principles of electron behavior: the Pauli exclusion principle and Hund’s rule.
The Pauli exclusion principle states that no two electrons in the same atom or molecule can have identical values for all four quantum numbers. In a single orbital, this means that at most two electrons can occupy the orbital, and they must have opposite spins.
Hund’s rule describes how electrons fill orbitals of equal energy. When several degenerate orbitals are available, electrons generally occupy them singly with parallel spins before pairing up, provided this arrangement gives the lowest-energy state.
In oxygen, the two π* antibonding orbitals have the same energy. The two remaining valence electrons could, in principle, occupy one orbital as a pair. However, the ground-state arrangement places one electron in each orbital, with their spins parallel.
This arrangement reduces the energy associated with electron–electron interactions compared with the corresponding paired arrangement. As a result, the lowest-energy electronic state has two unpaired electrons.
The simplified valence-level occupancy can be represented as:
- σ(2s): 2 electrons
- σ*(2s): 2 electrons
- σ(2p): 2 electrons
- π(2pₓ) and π(2pᵧ): 4 electrons
- π*(2pₓ) and π*(2pᵧ): 2 electrons, one in each orbital
The final pair of electrons is the essential feature. Each occupies a different antibonding π orbital, and their parallel spins give the molecule a nonzero electronic magnetic moment.
Oxygen is therefore paramagnetic not because it has an unusual number of electrons, but because quantum mechanics dictates how those electrons occupy the available molecular orbitals.
How the molecular orbital configuration explains the O₂ bond
Molecular orbital theory also explains why oxygen has a double bond, despite having electrons in antibonding orbitals.
The bond order is calculated by subtracting the number of electrons in antibonding molecular orbitals from the number in bonding molecular orbitals and dividing the result by two:
Bond order = ½ × (number of bonding electrons − number of antibonding electrons)
For O₂, there are eight electrons in bonding valence orbitals and four in antibonding valence orbitals. Therefore:
Bond order = ½ × (8 − 4) = 2
A bond order of 2 is consistent with a double bond between the oxygen atoms.
The antibonding electrons weaken the bond relative to what it would be if those orbitals were empty, but they do not cancel the bonding effect completely. The net result is a stable molecule with a bond order of two.
This illustrates an important point: antibonding electrons do not automatically prevent a molecule from forming. What matters is the balance between bonding and antibonding occupancy.
It also shows that bonding and magnetism are connected through the same electron arrangement. The molecular orbitals that determine the strength of the O–O bond also determine whether the molecule has unpaired electrons.
Why a magnetic field attracts oxygen
The two unpaired electrons in O₂ give the molecule a net electronic magnetic moment. In the presence of an external magnetic field, the molecule’s magnetic states respond to the field, producing an energetic preference for alignment with it.
In a nonuniform magnetic field, this interaction creates a net force toward the region of stronger field for a paramagnetic substance. This is why oxygen can be drawn toward a strong magnet.
The attraction is relatively weak under ordinary conditions, and oxygen gas does not behave like a permanent magnet. Its molecular magnetic moments are not permanently aligned in the absence of an applied field.
Temperature also matters. Thermal motion tends to disrupt alignment, so the magnetic response of a paramagnetic substance generally becomes weaker as temperature rises, all else being equal. Oxygen’s observed magnetic behavior reflects the balance between magnetic interactions and thermal energy.
Liquid oxygen makes the effect easier to demonstrate because it is much denser than gaseous oxygen. A visible sample can be held in the strong field between magnet poles, providing a striking demonstration of a property that originates from electron occupancy at the molecular scale.
Why molecular orbital theory succeeds where simpler models fail
The paramagnetism of oxygen is one of the classic examples showing why molecular orbital theory is essential for understanding molecular electronic structure.
A simple Lewis structure emphasizes paired electrons in localized bonds and lone pairs. Although it captures many useful features of O₂, it does not represent the two unpaired electrons in the ground state.
Molecular orbital theory instead accounts for all the relevant electrons and the energies of the orbitals available to them. It predicts that the two highest occupied valence electrons enter separate, degenerate π* antibonding orbitals. Hund’s rule favors their parallel spins, leaving both electrons unpaired.
This conclusion is not merely a convenient theoretical picture. It agrees with the experimentally established paramagnetism of oxygen and with other spectroscopic evidence about its electronic ground state.
More advanced descriptions use quantum-mechanical methods to account for electron interactions and the detailed structure of oxygen’s electronic states. Nevertheless, the standard molecular orbital model captures the central explanation accurately: the arrangement of electrons in the two π* orbitals produces the unpaired electrons responsible for oxygen’s magnetic behavior.
The essential connection between electrons and magnetism
Oxygen’s paramagnetism emerges from the same quantum rules that govern chemical bonding. Its 12 valence electrons fill molecular orbitals in order of increasing energy, subject to the Pauli exclusion principle and Hund’s rule. Because the two final electrons occupy separate, equal-energy π* antibonding orbitals, they remain unpaired in the ground state.
Those unpaired electrons give O₂ a net electronic magnetic moment, making oxygen attracted to an external magnetic field. At the same time, the balance of bonding and antibonding electrons gives the molecule a bond order of two.
The explanation is a powerful lesson in chemistry: a molecule’s properties cannot always be understood from its Lewis structure alone. To explain oxygen’s magnetism, it is necessary to understand not only how many electrons a molecule contains, but also which molecular orbitals those electrons occupy and how quantum mechanics governs their arrangement.