Oxygen (O, Z = 8): The Breath of Life and Architect of Planetary Chemistry

History of the Discovery of Oxygen

Oxygen was discovered independently by several chemists in the 1770s. The Swedish chemist Carl Wilhelm Scheele (1742-1786) isolated oxygen in 1772 by heating various oxides, but his work was not published until 1777. In 1774, the British theologian and chemist Joseph Priestley (1733-1804) produced oxygen by focusing sunlight on mercury oxide using a lens. He observed that a candle burned more vigorously in this gas and that a mouse survived longer in it. It was Antoine Lavoisier (1743-1794) who, between 1775 and 1780, truly understood the nature of oxygen and its role in combustion and respiration, thus overturning the phlogiston theory. Lavoisier named this element oxygen (from the Greek oxys = acid and genes = generate), mistakenly believing that all acids contained oxygen. This discovery marked the beginning of modern chemistry.

Structure and Fundamental Properties

Oxygen (symbol O, atomic number 8) is a non-metal of group 16 (chalcogens) in the periodic table, consisting of eight protons, usually eight neutrons (for the most common isotope), and eight electrons. The three stable isotopes are oxygen-16 \(\,^{16}\mathrm{O}\) (≈ 99.757%), oxygen-17 \(\,^{17}\mathrm{O}\) (≈ 0.038%), and oxygen-18 \(\,^{18}\mathrm{O}\) (≈ 0.205%).
At room temperature, oxygen exists as a diatomic gas (O₂), colorless, odorless, and highly chemically reactive. The O₂ molecule has a unique electronic configuration with two unpaired electrons, giving it paramagnetic properties (liquid oxygen is attracted by a magnet). Oxygen gas constitutes about 21% of Earth's atmosphere by volume and is essential for aerobic respiration. O₂ gas has a density of about 1.429 g/L at standard temperature and pressure.
Oxygen also exists as ozone (O₃), a triatomic allotropic form, pale blue in color, with a characteristic odor, which strongly absorbs ultraviolet rays. The stratospheric ozone layer protects terrestrial life from harmful UV radiation.
The temperature at which the liquid and solid states can coexist (melting point): 54.36 K (−218.79 °C). The temperature at which it transitions from liquid to gas (boiling point): 90.188 K (−182.962 °C). Liquid oxygen has a characteristic pale blue color.

Table of Oxygen Isotopes

Electronic Configuration and Electron Shells of Oxygen

N.B.:
Electron shells: How electrons organize around the nucleus (new window).

Oxygen has 8 electrons distributed across two electron shells. Its full electronic configuration is: 1s² 2s² 2p⁴, or simplified as: [He] 2s² 2p⁴. This configuration can also be written as: K(2) L(6).

Detailed Structure of the Shells

K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 6 electrons distributed as 2s² 2p⁴. The 2s orbitals are complete, while the 2p orbitals contain only 4 out of 6 possible electrons. Thus, 2 electrons are missing to achieve the stable neon configuration with 8 electrons (octet).

Valence Electrons and Oxidation States

The 6 electrons in the outer shell (2s² 2p⁴) are the valence electrons of oxygen. This configuration explains its chemical properties:
By gaining 2 electrons, oxygen forms the O²⁻ ion (oxidation state -2), its most common state in oxides, adopting the stable neon configuration [Ne].
Oxygen can also exhibit an oxidation state of -1 in peroxides (such as H₂O₂, hydrogen peroxide) and superoxides.
In certain compounds with fluorine (OF₂), oxygen exceptionally exhibits a positive oxidation state of +2, being the only element capable of causing oxygen to lose electrons.
The oxidation state 0 corresponds to dioxygen O₂, its natural molecular form, where two oxygen atoms are bonded by a double bond.

Oxygen's electronic configuration, with 6 valence electrons, classifies it among the chalcogens (Group 16) and makes it the second most electronegative element after fluorine (electronegativity of 3.5). This structure gives it characteristic properties: high chemical reactivity (oxygen reacts with almost all elements), strong oxidizing power (second only to fluorine), and the ability to form two covalent bonds to complete its octet. Oxygen primarily forms the oxide ion O²⁻ in ionic compounds but can also establish covalent bonds by sharing its electrons. Dioxygen O₂ is a colorless, odorless, paramagnetic gas essential for the respiration of aerobic organisms. Its molecule has two unpaired electrons, explaining its paramagnetism and reactivity.

Oxygen is fundamentally important: it constitutes about 21% of Earth's atmosphere and approximately 46% of the Earth's crust by mass (the most abundant element). It is essential for life (cellular respiration, ATP energy production), combustion, and countless chemical processes. Oxygen is used industrially in metallurgy (steel production), welding, chemical processes, and in medicine (oxygen therapy). Ozone O₃, an allotropic form, protects the Earth from ultraviolet rays in the stratosphere.

Chemical Reactivity

Oxygen has six valence electrons and is the third most electronegative element (after fluorine and chlorine), making it an extremely powerful oxidant. It forms compounds with practically all other chemical elements, except for the light noble gases (helium, neon, argon). Oxygen typically forms two covalent bonds (as in H₂O, CO₂) or oxide ions O²⁻ in ionic compounds.

Oxidation reactions (combustion, respiration, rust, etc.) involve the transfer of electrons to oxygen. The combustion of organic matter with oxygen releases large amounts of energy in the form of heat and light. This high reactivity is exploited in countless natural and industrial processes, but also makes oxygen potentially dangerous: oxygen-enriched atmospheres significantly increase fire risks.

Oxygen forms oxides with all elements except the light noble gases. These oxides can be basic (metallic oxides like CaO), acidic (non-metallic oxides like SO₂, CO₂), or amphoteric (like Al₂O₃). Water (H₂O), the most important oxygen compound, covers 71% of Earth's surface and is essential to all known life.

In living organisms, oxygen is used in aerobic cellular respiration to oxidize organic molecules (glucose) and produce energy (ATP). This respiration also generates reactive oxygen species (free radicals) that can damage cells, against which organisms have developed antioxidant systems. Paradoxically, oxygen, essential for aerobic life, is also an oxidative poison at high concentrations.

Industrial and Technological Applications of Oxygen

• in the medical field for oxygen therapy, anesthesia, and intensive care;
• for steel production (basic oxygen steelmaking) by removing impurities from molten iron;
• as an oxidizer in rocket engines (liquid oxygen mixed with hydrogen or kerosene);
• in oxy-fuel welding and cutting;
• for oxygen enrichment in wastewater treatment (aeration);
• in the production of chemicals (ethylene, propylene, ethylene oxide, etc.);
• for bleaching paper pulp in the paper industry;
• in life support systems for submarines, space stations, and spacesuits;
• as an oxidant in industrial combustion processes to improve energy efficiency;
• in glass manufacturing (oxygen-enriched combustion to achieve high temperatures);
• for water disinfection and purification (ozonation);
• in aquaculture to maintain adequate dissolved oxygen levels;
• as an oxidant in medical and industrial waste incineration systems;
• in the production of technical glass and ceramics;
• for coal gasification and synthesis gas production;
• in the chemical industry for the catalytic oxidation of organic compounds;
• as a bleaching agent in the textile industry;
• in hyperbaric chambers to treat decompression sickness and certain infections;
• for breathing at high altitudes (aviation, mountaineering) and in underwater diving;
• in the production of non-ferrous metals (copper, zinc, nickel) by pyrometallurgy.

Role in Astrophysics and Cosmology

Oxygen is the third most abundant element in the observable universe (after hydrogen and helium) and the first heavy element by cosmic abundance. It represents about 1% of the total baryonic mass of the universe. Unlike primordial elements (H, He, Li), oxygen is entirely produced by stellar nucleosynthesis.

Oxygen is mainly synthesized in massive stars (greater than 8 solar masses) by the carbon and helium burning processes. The triple-alpha reaction produces carbon-12, which then captures a helium-4 nucleus to form oxygen-16. At higher temperatures (about 1 billion kelvin), carbon fusion also produces oxygen. Massive stars develop a layered structure (like an onion) with zones of combustion of different elements, including an oxygen-rich layer.

Oxygen is massively dispersed into the interstellar medium during type II supernova explosions. These cataclysmic events eject the oxygen-rich outer layers of the star at speeds of thousands of kilometers per second, enriching the interstellar medium for future generations of stars and planets. Oxygen represents a significant fraction of the mass ejected by supernovae, making these events the main sources of galactic oxygen.

In the interstellar medium, oxygen exists in several forms: atomic (O, O⁺, O⁺⁺), molecular (O₂, which is rare and difficult to detect), and incorporated into many molecules such as H₂O (water ice), CO (carbon monoxide, the second most abundant molecule after H₂), CO₂, OH, and complex organic molecules. Doubly ionized atomic oxygen (O⁺⁺) emits characteristic spectral lines in planetary nebulae and HII regions, allowing astronomers to map the distribution of oxygen in galaxies.

The isotopic ratio ¹⁶O/¹⁸O in different astronomical objects (meteorites, comets, interstellar dust, presolar grains) reveals crucial information about nucleosynthesis processes in different types of stars and the history of chemical enrichment in our galaxy. Isotopic anomalies of oxygen discovered in certain refractory inclusions of primitive meteorites indicate the contribution of different stellar sources to the material that formed the solar system.

In planetary atmospheres, oxygen plays a central role. On Earth, atmospheric oxygen (21% O₂) is almost entirely of biological origin, produced by photosynthesis of plants, algae, and cyanobacteria for about 2.4 billion years (the "Great Oxygenation Event"). This accumulation of oxygen profoundly transformed Earth's chemistry and allowed the evolution of complex aerobic life. The spectroscopic detection of molecular oxygen and ozone in the atmosphere of an exoplanet could constitute a potential biosignature, although abiotic processes could also produce oxygen under certain conditions.

N.B.:
The "oxygen paradox" illustrates the dual nature of this essential element. Molecular oxygen (O₂) is absolutely vital for aerobic organisms, allowing efficient cellular energy production via mitochondrial respiration. However, oxygen is also a powerful oxidative poison: its reactive derivatives (superoxide radicals, hydrogen peroxide, hydroxyl radicals) damage proteins, lipids, and DNA. Aerobic organisms have had to develop sophisticated antioxidant mechanisms (enzymes such as superoxide dismutase, catalase, peroxidase; antioxidant molecules such as vitamins C and E) to protect themselves from oxygen toxicity while exploiting its energetic potential. Oxygen is also responsible for cellular aging through the accumulation of oxidative damage over time. This remarkable duality—vital and toxic simultaneously—reflects the complex evolutionary history of life on Earth and the biological compromises necessary to exploit the enormous energetic potential of oxygen.