Why is liquid oxygen attracted to a magnet?
Pour liquid oxygen between the poles of a strong magnet and it does not fall through. It bridges the gap and hangs there, pale blue, until it boils away.
This is a genuine problem for the Lewis structure taught in every introductory course. That structure shows O=O, a double bond, with two lone pairs on each atom and every electron paired. A molecule with all electrons paired is diamagnetic — weakly repelled by a magnetic field, not attracted to it.
The observation says oxygen has two unpaired electrons. The Lewis model says it has none. The model is wrong, and this is the classic demonstration of why molecular orbital theory was needed.
In the MO picture, the atomic orbitals of the two oxygen atoms combine into molecular orbitals spread over the whole molecule. Filling them in energy order puts the last two electrons into two degenerate antibonding π* orbitals. By Hund’s rule — the same rule that keeps carbon’s 2p electrons apart — they occupy separate orbitals with parallel spins rather than pairing in one.
So O₂ has a bond order of 2, matching the double bond Lewis predicted, but with two unpaired electrons that Lewis cannot accommodate. Both facts are true at once, and only the MO treatment produces both.
This is worth dwelling on because it is one of the few places where a first-year model fails visibly rather than subtly. You can watch it fail.
Why is oxygen’s atomic weight a range?
Oxygen is one of fourteen elements for which IUPAC publishes an interval rather than a single value: [15.99903, 15.99977].
Natural oxygen is a mixture of three stable isotopes — ¹⁶O at about 99.76%, ¹⁸O at 0.20%, and ¹⁷O at 0.04% — and the proportions vary measurably between sources. Evaporation preferentially removes the lighter ¹⁶O from liquid water, so water vapour, rain, snow and ice each carry a different ratio, and the difference depends on temperature.
That temperature dependence turned oxygen isotopes into the backbone of palaeoclimatology. The δ¹⁸O ratio in an ice core, a foraminifera shell or a cave formation records the temperature at which the water condensed. Much of what is known about ice ages comes from this single measurement.
There is a historical footnote worth knowing. Until 1961, chemists and physicists used different atomic mass scales: chemists set natural oxygen to exactly 16, physicists set ¹⁶O to exactly 16. Because natural oxygen contains heavier isotopes, the two scales differed by about 275 parts per million, and every published atomic mass had to be labelled with which convention it used. The switch to carbon-12 as the standard settled it, for the specific reason that carbon’s isotopic variation was smaller and better characterised.
How do you read oxygen’s electron configuration?
Oxygen’s ground state is [He] 2s² 2p⁴: a filled 1s core, two electrons in 2s, and four in 2p.
Four electrons across three 2p orbitals means one orbital holds a pair and two hold single electrons. Hund’s rule again determines this — the two singles have parallel spins — and it is the direct atomic-level cause of the paramagnetism described above, carried through into the molecule.
Two electrons short of a filled shell, with a small atomic radius and a high effective nuclear charge, is a combination that makes oxygen an aggressive electron acceptor. Its electronegativity of 3.44 is second only to fluorine’s. The practical consequence is that oxygen forms compounds with almost every element, and the release of energy when it does is what powers respiration, combustion and most of industrial chemistry.
Its ground state term symbol is ³P₂. The superscript 3 confirms the two unpaired 2p electrons.
Why did oxygen change the Earth?
For roughly the first two billion years of Earth’s history there was effectively no free oxygen in the atmosphere. The air was nitrogen, carbon dioxide, methane and water vapour.
Cyanobacteria changed that. Oxygenic photosynthesis splits water and releases O₂ as waste, and the waste accumulated. It was absorbed first by dissolved iron in the oceans, precipitating the banded iron formations that supply most of the world’s iron ore today, then by rock weathering, and only afterwards did it begin to build up in the air. The Great Oxidation Event, around 2.4 billion years ago, was very likely the largest mass extinction in Earth’s history: for the anaerobic life that dominated at the time, oxygen was poison.
It also made complex life possible. Aerobic respiration extracts roughly fifteen times more energy per glucose molecule than fermentation, and that energy budget is what large multicellular organisms run on. The ozone layer that formed from accumulated O₂ then screened enough ultraviolet radiation for life to leave the water.
The element that is now about 21% of the atmosphere and 46% of the Earth’s crust by mass arrived as a pollutant.
What does the ozone layer do?
Ozone in the stratosphere, between roughly 15 and 35 kilometres up, absorbs ultraviolet radiation in the UV-B and UV-C bands. Without it, the surface radiation environment would be lethal to most life on land.
The chemistry is a cycle. Ultraviolet light splits O₂ into two oxygen atoms; each combines with another O₂ to make ozone; further ultraviolet splits the ozone back apart. The cycle absorbs energy at every step, and the ozone is continually destroyed and remade rather than accumulating.
Chlorofluorocarbons broke it. CFCs are inert enough to survive the journey to the stratosphere intact, where ultraviolet finally breaks them and releases chlorine atoms. A single chlorine atom catalyses the destruction of thousands of ozone molecules before it is removed, because it emerges from each reaction unchanged and available to do it again.
The Antarctic ozone hole was identified in 1985, and the Montreal Protocol phasing out CFCs was agreed in 1987 — an unusually fast response. It is generally regarded as the most successful international environmental agreement ever concluded, and stratospheric ozone is now measurably recovering, with full recovery projected around the middle of this century.
Ground-level ozone is a separate matter entirely: there it is a pollutant, a component of smog, and damaging to lungs and crops.