Why Is Oxygen Paramagnetic? The MO Theory Answer

Pour liquid oxygen — a pale blue liquid that boils at −183 °C — between the poles of a strong magnet, and something remarkable happens. Instead of flowing straight through, it clings to the magnet, bridging the gap between the poles. Pour liquid nitrogen through the same magnet, and it passes straight by. This simple demonstration exposes a flaw in the Lewis structure of oxygen and is one of the most convincing pieces of evidence for molecular orbital theory. Here’s the full explanation.

Paramagnetism and diamagnetism

Electrons have a property called spin, which makes each one behave like a tiny magnet.

• When electrons are paired in an orbital, their spins are opposite and their magnetic effects cancel.
• An unpaired electron has an uncancelled magnetic moment.

This leads to two kinds of magnetic behaviour:

Behaviour	Electrons	Response to a magnet	Examples
Paramagnetic	One or more unpaired	Attracted into a magnetic field	O₂, NO, Fe³⁺ ions
Diamagnetic	All paired	Very weakly repelled	N₂, H₂O, most organic compounds

(Ferromagnetism — the strong, permanent magnetism of iron — is a different, collective effect in solids.)

So if oxygen is attracted to a magnet, O₂ molecules must contain unpaired electrons. How many? Measurements of the strength of the attraction show two per molecule.

The problem with the Lewis structure

Draw O₂ using the rules for Lewis structures (see Lewis dot structures):

• Total valence electrons: 2 × 6 = 12.
• Structure: O=O, with two lone pairs on each oxygen.

Every electron is in a pair: two bonding pairs and four lone pairs. This predicts O₂ should be diamagnetic.

Could a different Lewis structure fix it? You could draw a single bond with an unpaired electron on each oxygen (•O–O•), which would be paramagnetic. But each oxygen would then have only seven electrons, and the bond order would be 1. That contradicts other evidence: O₂’s bond length (121 pm) and bond enthalpy (498 kJ mol⁻¹) clearly fit a double bond, not a single one (a typical O–O single bond is 148 pm and about 146 kJ mol⁻¹; see bond length and bond strength).

We need a model that gives O₂ a double bond and two unpaired electrons at the same time. Lewis structures can’t do that.

The MO theory solution

In molecular orbital theory, the atomic orbitals of the two oxygen atoms combine to form molecular orbitals spread over the whole molecule (see molecular orbital theory for beginners).

For O₂ (as for F₂), the valence molecular orbitals in order of increasing energy are:

σ2s < σ*2s < σ2p < π2p (two orbitals, equal energy) < π*2p (two orbitals, equal energy) < σ*2p

Now fill in O₂’s 12 valence electrons:

1. σ2s: 2 electrons
2. σ*2s: 2 electrons (running total 4)
3. σ2p: 2 electrons (6)
4. π2p: 4 electrons, two in each of the two π orbitals (10)
5. π*2p: 2 electrons left, going into two orbitals of equal energy (12)

Here’s the key step. Hund’s rule says that when electrons enter orbitals of equal energy, they occupy them singly, with parallel spins, before pairing up. Pairing two electrons in the same orbital costs extra energy because of their mutual repulsion.

So the last two electrons go one into each π* orbital, with the same spin:

π*2p: ↑ ↑

That’s two unpaired electrons. O₂ is paramagnetic. ✓

And the bond order still comes out as 2

Count bonding and antibonding electrons:

• Bonding: σ2s (2) + σ2p (2) + π2p (4) = 8
• Antibonding: σ*2s (2) + π*2p (2) = 4

Bond order = ½(8 − 4) = 2. ✓

So MO theory gives exactly what the evidence demands: a double bond (consistent with the bond length and energy) and two unpaired electrons (consistent with the magnetism). It does this naturally, without any special adjustments — which is why the paramagnetism of oxygen is often called the classic triumph of MO theory.

Comparing N₂, O₂ and F₂

Molecule	Valence electrons	π* electrons	Bond order	Unpaired	Behaviour in a magnet
N₂	10	0	3	0	Passes through (diamagnetic)
O₂	12	2 (one in each π*)	2	2	Attracted (paramagnetic)
F₂	14	4 (both π* full)	1	0	Diamagnetic

Oxygen is special because it sits exactly at the point where the two π* orbitals are half filled. One fewer pair of electrons (N₂) and they’re empty; one more pair (F₂) and they’re full.

Triplet and singlet oxygen

Chemists describe the spin state of a molecule by its multiplicity. Ordinary O₂, with two unpaired electrons of parallel spin, is called triplet oxygen. It’s the ground state — the lowest-energy form, and the oxygen in the air you breathe.

If energy is added (for example, by light in the presence of a dye), the two π* electrons can pair up with opposite spins, giving singlet oxygen. Singlet oxygen lies about 94 kJ mol⁻¹ above the ground state and is far more reactive. It’s used in photodynamic therapy against some cancers and is one reason light and dyes together can damage biological molecules.

Why oxygen’s unpaired electrons matter

Oxygen’s spin state has a surprising benefit for life. Most organic molecules — sugars, fats, proteins, wood — have all their electrons paired. For triplet O₂ to react directly with them, electron spins would have to flip during the reaction, which quantum rules make unlikely. The result is that reactions of O₂ with organic matter are thermodynamically favourable but kinetically slow at room temperature.

That’s why paper, wood and your own body don’t spontaneously burst into flames in air, even though their combustion releases lots of energy. A spark or flame supplies enough energy to create radicals, which react readily with triplet O₂ and start a chain reaction.

Living things take advantage of O₂’s special electronic structure too:

• Haemoglobin and other oxygen-binding proteins use iron (or copper) ions, which have unpaired electrons of their own and can interact with O₂’s unpaired electrons (see haemoglobin).
• In the electron transport chain, O₂ accepts electrons one at a time at a metal-containing enzyme. Occasionally a single electron escapes to O₂ early, forming superoxide (O₂⁻) — a source of oxidative stress (see the chemistry of ageing).

A demonstration you can see

In university labs, the liquid oxygen demonstration is done with a strong electromagnet or a powerful permanent magnet. The pale blue liquid oxygen is held suspended between the poles until it evaporates. Liquid nitrogen poured the same way simply falls through. (This involves cryogenic liquids at below −180 °C and a strong oxidiser, so it’s for trained demonstrators only — liquid oxygen can make organic materials ignite violently.)

Common misconceptions

• “O₂ is magnetic like iron.” It’s paramagnetic — attracted to a magnet only while in the field — not a permanent magnet.
• “Unpaired electrons mean a single bond.” MO theory shows O₂ has both a bond order of 2 and two unpaired electrons.
• “The Lewis structure of O₂ is simply wrong.” It correctly shows a double bond; it just can’t show unpaired electrons in antibonding orbitals.
• “Paramagnetism needs a metal.” O₂, NO and other radicals are paramagnetic non-metal molecules.

Key takeaways

• Paramagnetism means unpaired electrons; liquid O₂ is attracted to a magnet, so O₂ has unpaired electrons.
• The Lewis structure O=O predicts all electrons paired — it can’t explain the magnetism.
• In MO theory, O₂’s last two electrons occupy two degenerate π* orbitals singly (Hund’s rule), giving two unpaired electrons.
• The bond order is still 2, matching O₂’s bond length and strength.
• O₂’s triplet ground state makes its reactions with organic matter slow at room temperature.

For more practice, see bond order from molecular orbital diagrams.