eduKate Learning Manual · Chemistry × Physics · Secondary → JC · Observe → Model → Predict → Check
Wait, What? Oxygen Is Magnetic
The oxygen molecules in air are not tiny bar magnets, but molecular oxygen is paramagnetic: it is attracted into a magnetic field. Cool oxygen until it becomes a pale blue liquid and the effect becomes dramatic enough to demonstrate directly.
This is more than a curiosity. Oxygen’s magnetism is one of the classic clues that a simple Lewis structure does not tell the whole electronic story of a molecule.
Atomic orbitals combine → molecular orbitals form → electrons occupy them according to energy and spin rules → two electrons remain unpaired in O₂ → the molecule is paramagnetic.
The Big Question
How can a molecule as familiar as O₂ reveal a limit of one bonding model and point us toward a deeper one?
Quick Answer
In molecular orbital theory, atomic orbitals combine to create orbitals spread across the molecule. For ground-state O₂, the electron configuration leaves two electrons unpaired in separate antibonding π* orbitals. Unpaired electron spins create a net magnetic response, so O₂ is attracted to an applied magnetic field.
What You Will Learn
- what paramagnetism and diamagnetism mean
- why paired and unpaired electrons matter
- how bonding and antibonding molecular orbitals are formed
- why O₂ has bond order 2 in a simple MO treatment
- why the observed magnetism of oxygen supports molecular orbital theory
- how models are judged by the predictions they can make
Part 1 — Magnetism Can Come from Electrons
Electron spin is associated with a magnetic moment. When electrons are paired with opposite spins in the same orbital, their spin magnetic effects largely cancel. When one or more electrons remain unpaired, the substance can show paramagnetism.
This gives us a measurable bridge from an invisible electronic arrangement to a macroscopic observation.
Part 2 — Why a Lewis Structure Is Useful but Incomplete
A familiar Lewis structure for O₂ places a double bond between two oxygen atoms. This is useful for electron counting and many reaction problems. But if interpreted too literally as a complete electronic picture, it does not naturally predict the observed two unpaired electrons in ground-state oxygen.
That does not make Lewis structures “wrong.” It means a model has a domain of usefulness. Molecular orbital theory resolves a question that the simpler representation does not.
Part 3 — Build Molecular Orbitals
When atomic orbitals interact, combinations can produce molecular orbitals of different energies. A constructive combination concentrates electron density in a way that stabilises the molecule and is called bonding. A corresponding higher-energy combination is antibonding and is marked with an asterisk.
Electrons fill the available molecular orbitals according to familiar quantum rules: lower-energy orbitals are occupied first, each orbital accommodates at most two electrons with opposite spins, and degenerate orbitals are singly occupied before pairing.
Part 4 — The Two Electrons That Change the Prediction
In the usual valence molecular-orbital diagram for O₂, the highest occupied orbitals are two degenerate π* antibonding orbitals. Each contains one electron. Those two electrons are unpaired.
That predicts paramagnetism before we ever place oxygen near a magnet. This is a strong example of what scientific models are for: not merely to redraw what we already know, but to generate testable consequences.
A Quantitative Window — Bond Order
A simple molecular-orbital bond order is:
bond order = ½(number of bonding electrons − number of antibonding electrons)
For O₂, the valence-electron count gives bond order 2. This agrees with the familiar idea that oxygen has a strong double-bond-like interaction, while the orbital occupancy simultaneously predicts two unpaired electrons.
Part 5 — How Do We Know Oxygen Is Paramagnetic?
Magnetic susceptibility can be measured quantitatively. NIST reports high-precision measurements of gaseous oxygen’s magnetic susceptibility that agree closely with modern theoretical calculations. The observation is therefore not merely a classroom demonstration; it is a measurable property of molecular oxygen.
Liquid oxygen provides a visually striking version because many O₂ molecules are concentrated in a small volume. A strong magnetic field can attract the liquid into the region of higher field strength.
Part 6 — Follow One Electron
Imagine one of the last electrons placed into the O₂ molecular-orbital diagram. Lower orbitals have already been filled. Two equivalent π* orbitals remain available. Instead of pairing immediately with an electron already in one orbital, the electron occupies the other degenerate orbital with parallel spin in the simple ground-state picture.
That occupancy changes what the whole molecule does in a magnetic field. One quantum-scale placement contributes to a bulk property we can measure.
Think Like a Scientist — Model Competition
Suppose two models both explain that oxygen atoms form O₂. Model A predicts all electrons paired. Model B predicts two unpaired electrons. An experiment shows O₂ is paramagnetic.
Which model has gained evidence?
Model B. The key is not that it looks more sophisticated. It made the prediction that survived contact with measurement.
Observation vs Inference
Observation: oxygen is attracted into a magnetic field; quantitative susceptibility is positive.
Inference: the electronic ground state contains unpaired electron spin, consistent with the molecular-orbital description of O₂.
Common Misconceptions and Repairs
- “Oxygen is ferromagnetic like iron.” Repair: O₂ is paramagnetic; it does not behave like a permanent ferromagnet.
- “A double bond means every electron is paired.” Repair: that conclusion comes from treating a Lewis picture as a complete orbital model.
- “Antibonding electrons mean there is no bond.” Repair: bond strength depends on the balance between bonding and antibonding occupancy; O₂ still has positive bond order.
- “More complicated models are automatically better.” Repair: a model earns value by explaining and predicting evidence within its domain.
Checkpoint Questions
- What makes a species paramagnetic?
- Why can electron pairing reduce net magnetic response?
- Where are the two unpaired electrons in the simple MO picture of ground-state O₂?
- What is the bond order of O₂ in the standard valence MO treatment?
- Why does oxygen magnetism provide evidence about electronic structure?
Apply It — Compare O₂ and N₂
Use molecular-orbital occupancy to predict whether N₂ should be paramagnetic or diamagnetic. Then ask what experimental magnetic behaviour would count as a check on the prediction.
Answer Key
1. One or more unpaired electrons. 2. Opposite spins in a paired orbital largely cancel their spin magnetic moments. 3. In the two degenerate π* antibonding orbitals. 4. Two. 5. Magnetic susceptibility is a macroscopic measurement that depends on microscopic electron states. N₂ is predicted to be diamagnetic in the standard MO picture because its occupied valence orbitals contain paired electrons.
Can You Explain WHY?
Explain why oxygen’s attraction to a magnet is evidence about how electrons occupy molecular orbitals. A strong answer should connect orbital energies → electron filling → unpaired spin → magnetic susceptibility → observation.
Singapore Secondary and JC Science Bridge
Secondary Chemistry builds models of atomic structure, bonding and properties. JC Chemistry deepens the relationship between electron arrangement, structure and reactivity. Molecular orbital theory is a useful higher-resolution window showing why different representations answer different questions. It also links Chemistry to Physics through magnetic fields and quantum behaviour.
Deep Science Windows
- Triplet oxygen: ground-state O₂ has total electron spin corresponding to a triplet state.
- Singlet oxygen: electronically excited oxygen states have different spin arrangements and chemical behaviour.
- Spectroscopy: electronic states can be probed through absorption and emission, not only magnetism.
- Quantum chemistry: more advanced calculations refine the simple orbital diagrams used for teaching.
Evidence Boundaries
The classroom MO diagram is a model, not a literal picture of tiny tracks occupied by electrons. Real molecular electronic structure is described by quantum mechanics, and advanced treatments include electron correlation and more detailed wavefunctions. The core evidence remains robust: ground-state molecular oxygen has unpaired electron spin and is paramagnetic.
Manual Summary — KNOW → CONNECT → EXPLAIN → APPLY → CHECK
- KNOW: O₂ is paramagnetic.
- CONNECT: unpaired electrons produce a magnetic response.
- EXPLAIN: MO theory predicts two unpaired electrons in π* orbitals.
- APPLY: predict magnetic behaviour from electron occupancy.
- CHECK: compare model predictions with susceptibility measurements.
Teaching Guide for Parents, Tutors and Teachers
Why this opening works: oxygen feels familiar and non-magnetic. The contradiction creates a reason to learn a deeper bonding model.
- Central model: electron occupancy produces observable properties.
- Teaching sequence: magnetic observation → paired/unpaired spin → molecular orbitals → O₂ filling → prediction → evidence.
- Diagnostic question: “What observation would force us to improve a bonding model?”
- If stuck: begin with paired versus unpaired arrows before drawing a full MO diagram.
- Ready for more: compare O₂, O₂⁺, O₂⁻, N₂ and the singlet/triplet distinction.
Quiet Teaching Standard: do not turn the orbital diagram into decorative memorisation. Every arrow should answer a physical question.