Oxygen

And how many large and small stoves are there in the world that burn firewood. If we look beyond home heating, industrial furnaces—like blast furnaces, open-hearth furnaces, and brick kilns—also rely on oxygen. They work with wood, and wood cannot burn without oxygen. If it were not for the green leaves of plants, we would have long since died out from lack of oxygen. When sunlight hits a leaf, it absorbs carbon dioxide from the air and releases oxygen.

This is why oxygen in the Earth's atmosphere is constantly replenished, rather than being depleted. Oxygen is not only found in the air. It combines with another gas, hydrogen, to form water. Approximately half of the atoms in the Earth's crust are oxygen atoms.

Oxygen received its name from the French chemist Antoine Lavoisier, who coined the term "oxygenium" in the late 1770s. He derived it from the Greek words oxys (acid) and genes (forming), because he mistakenly believed that this gas was a necessary component for the creation of all acids. Thus, the name literally translates to "acid-former."

The Path to Discovery

The journey to this name began on August 1, 1774, when the English chemist Joseph Priestley discovered the gas. By heating red mercuric oxide, he observed that it released a substance that caused candles to burn much more intensely than regular air.

However, Priestley did not realize he had found a new element. He thought he had simply identified a particularly pure form of air and called it "dephlogisticated air."

From Discovery to Definition

Priestley shared his findings with Antoine Lavoisier in Paris. Lavoisier took a more strategic approach to the data. He realized that this gas was not just "pure air" but a unique chemical element that was a constituent part of the atmosphere and played a vital role in combustion.

The Legacy of the Name

While Lavoisier’s theory that oxygen was the source of all acidity was later proven wrong (as some acids, like hydrochloric acid, do not contain oxygen), his name for the element became the scientific standard. It replaced Priestley's original name and shifted the focus of chemistry from ancient theories to the modern understanding of elements.

Oxygen is a reactive nonmetal with atomic number 8 and an atomic mass of 15.999 u. It belongs to group 16 (the chalcogens) of the periodic table and is the third most abundant element in the universe by mass, after hydrogen and helium. On Earth, oxygen makes up about 20.95% of the atmosphere by volume, approximately 46.6% of the Earth’s crust by mass, and nearly 65% of the mass of the human body, primarily due to its presence in water and organic compounds.

Under standard conditions, oxygen exists as a diatomic molecule (O₂), which is a colorless, odorless, and tasteless gas. Its melting point is −218.8 °C, and its boiling point is −183.0 °C. In liquid and solid states, oxygen has a pale blue color, which is a result of weak absorption in the red region of the visible spectrum. Oxygen is paramagnetic, meaning it is attracted to a magnetic field, a property that distinguishes it from most other diatomic gases.

Oxygen has a high electronegativity value of 3.44 on the Pauling scale, which explains its strong ability to attract electrons and form compounds. It readily reacts with most elements, forming oxides through exothermic oxidation reactions. This reactivity makes oxygen a key agent in combustion processes, corrosion, and biological respiration. In aerobic organisms, oxygen serves as the final electron acceptor in cellular respiration, allowing the efficient production of ATP. The average adult human consumes roughly 250 mL of O₂ per minute at rest.

The element exhibits several oxidation states, with −2 being the most common. However, oxygen can also display oxidation states of −1 (peroxides), −1/2 (superoxides), and positive values up to +2 in compounds with fluorine. Oxygen forms both covalent and ionic bonds, depending on the nature of the reacting element, and can participate in hydrogen bonding, which strongly influences the physical properties of water and many biological molecules.

Oxygen occurs in multiple allotropes. The most common is molecular oxygen (O₂), while ozone (O₃) is a less stable but highly reactive form. Ozone plays a critical role in the stratosphere, where the ozone layer absorbs most ultraviolet radiation with wavelengths between 200 and 315 nm, protecting life on Earth. Typical ozone concentrations peak at altitudes of 20–30 km above the Earth’s surface.

Oxygen was independently discovered in the early 1770s by Carl Wilhelm Scheele and Joseph Priestley, with Priestley publishing his results in 1774. Antoine Lavoisier later named the element and recognized its role in combustion and respiration, laying the foundation for modern chemical theory. Due to its unique physical and chemical characteristics, oxygen remains essential in environmental systems, industrial processes such as steelmaking and chemical synthesis, and virtually all known forms of complex life.

In the vast majority of known oxygen compounds, oxygen exhibits an oxidation state of −2. An oxidation state of −1 occurs only in a limited number of substances, such as peroxides. Compounds in which oxygen has other oxidation states are rare: −1/2 (superoxides), −1/3 (ozonides), 0 (elemental oxygen and hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and +2 (oxygen difluoride).

Oxides and other inorganic compounds

Water (H₂O) is an oxide of hydrogen and the most well-known oxygen-containing compound. In a water molecule, hydrogen atoms are covalently bonded to oxygen, but they also experience an additional attraction—about 23.3 kJ/mol per hydrogen atom—to oxygen atoms in neighboring molecules. These hydrogen bonds cause water molecules to remain approximately 15% closer together than would be expected in a simple liquid held together only by van der Waals forces.

Because of its high electronegativity, oxygen forms chemical bonds with nearly all other elements, producing a wide variety of oxides. The surfaces of many metals, including aluminium and titanium, become oxidized when exposed to air, forming a thin oxide layer that passivates the metal and reduces the rate of further corrosion. Many transition metal oxides are non-stoichiometric, meaning they contain slightly less metal than indicated by their ideal formulas. For instance, the mineral FeO (wüstite) is often written as Fe_1−xO, where x is typically close to 0.05.

Oxygen is also present in the atmosphere in small amounts as carbon dioxide (CO₂). The Earth’s crust is largely composed of oxygen-containing minerals, including oxides of silicon (silica SiO₂, found in granite and quartz), aluminium (aluminium oxide Al₂O₃, present in bauxite and corundum), iron (iron(III) oxide Fe₂O₃, found in hematite and rust), and calcium carbonate, which makes up limestone. In addition, the remainder of the crust consists mainly of oxygen compounds such as complex silicates. The Earth’s mantle, which has a much greater mass than the crust, is primarily composed of magnesium and iron silicates.

Water-soluble silicates, including Na₄SiO₄, Na₂SiO₃, and Na₂Si₂O₅, are commonly used as detergents and binding agents.

Oxygen can also function as a ligand for transition metals, forming metal–dioxygen complexes that contain metal–O₂ units. This category includes heme proteins such as hemoglobin and myoglobin. A particularly unusual reaction occurs with PtF₆, which is capable of oxidizing oxygen to form O⁺₂PtF₆⁻, known as dioxygenyl hexafluoroplatinate.

Organic compounds

Some of the most significant classes of oxygen-containing organic compounds include (where “R” represents an organic group): alcohols (R-OH), ethers (R-O-R), ketones (R-CO-R), aldehydes (R-CO-H), carboxylic acids (R-COOH), esters (R-COO-R), acid anhydrides (R-CO-O-CO-R), and amides (R-CO-NR₂). Numerous important organic solvents contain oxygen, such as acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone ((CH₃)₂CO) and phenol (C₆H₅OH) are widely used as precursor materials in the synthesis of many substances.

Other notable oxygen-containing organic compounds include glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are a class of ethers in which the oxygen atom is part of a three-membered ring. Oxygen is also present in nearly all biomolecules that are essential to life or produced by living organisms. It readily reacts with many organic compounds at or below room temperature through a process known as autoxidation. Most oxygen-containing organic compounds are not produced by the direct reaction with O₂. However, some industrially important substances, such as ethylene oxide and peracetic acid, are synthesized through direct oxidation of suitable precursors.

Oxygen plays a fundamental role in sustaining life on Earth and driving many natural and industrial processes. In the biosphere, it is essential for aerobic respiration, a biochemical process in which cells convert nutrients into energy. During cellular respiration, oxygen acts as the final electron acceptor in the electron transport chain, enabling the production of adenosine triphosphate (ATP). This process yields up to 36–38 ATP molecules per glucose molecule, compared to only 2 ATP in anaerobic metabolism, highlighting the efficiency provided by oxygen.

In the Earth’s atmosphere, oxygen maintains a stable concentration of about 20.95%, a level that has remained relatively constant for hundreds of millions of years. Geological evidence suggests that a major increase in atmospheric oxygen, known as the Great Oxidation Event, occurred approximately 2.4 billion years ago. This transformation, driven by photosynthetic cyanobacteria, enabled the evolution of complex multicellular life and fundamentally altered Earth’s chemistry by promoting the formation of oxide minerals and an ozone-rich atmosphere.

Oxygen also plays a critical protective role through the formation of the ozone layer in the stratosphere. Ozone (O₃) absorbs most harmful ultraviolet radiation in the UV-B and UV-C ranges, significantly reducing DNA damage in living organisms. The highest concentration of atmospheric ozone occurs at altitudes of approximately 20–30 km, where it filters solar radiation with wavelengths shorter than 315 nm. Without this oxygen-derived shield, life on land would be severely limited.

In natural cycles, oxygen is a key component of the carbon, nitrogen, and sulfur cycles. It enables the oxidation of organic matter during decomposition and plays a central role in the weathering of rocks, which regulates atmospheric carbon dioxide over geological timescales. Dissolved oxygen in oceans, typically ranging from 4 to 8 mg/L, is essential for aquatic life, and reductions below 2 mg/L can lead to hypoxic conditions, causing mass die-offs in marine ecosystems.

Oxygen is equally vital in industrial and technological applications. It is widely used in steel production, where high-purity oxygen increases furnace temperatures and improves efficiency. In medicine, supplemental oxygen has been used since the early 20th century to treat respiratory and cardiac conditions, and modern medical oxygen has a typical purity of ≥99%. Oxygen is also indispensable in aerospace technology, where liquid oxygen (LOX), with a boiling point of −183 °C, serves as a powerful oxidizer in rocket propulsion systems.

Through its roles in biology, atmospheric protection, geochemical cycles, and industry, oxygen remains one of the most influential elements on Earth. Its presence shapes the planet’s environment, enables complex life, and supports technologies that define modern civilization.

Oxygen is one of the most important industrial gases, playing a central role in manufacturing, energy production, and chemical processing. Large-scale industrial oxygen production began in the late 19th century after the development of cryogenic air separation. Today, millions of tons of oxygen are produced annually worldwide, with purities commonly exceeding 99.5% for industrial use.

In the steel and metallurgical industries, oxygen is essential for converting iron into steel. The basic oxygen furnace (BOF), introduced commercially in the 1950s, uses high-purity oxygen to remove carbon and impurities from molten iron. This process significantly reduces processing time—from several hours to less than 40 minutes—and improves fuel efficiency. Oxygen is also widely used in non-ferrous metal production, including copper, lead, and zinc refining.

Oxygen plays a key role in combustion-based industries by increasing flame temperatures and improving energy efficiency. In oxy-fuel combustion, replacing air with pure oxygen can raise flame temperatures above 3,000 °C. This technique is used in glass manufacturing, cement production, and metal cutting, where higher temperatures lead to faster processing and lower fuel consumption.

In the chemical industry, oxygen is a critical reactant in large-scale oxidation processes. It is used in the production of chemicals such as ethylene oxide, nitric acid, and hydrogen peroxide. For example, global ethylene oxide production exceeds 25 million tons per year, and oxygen-based processes are favored because they are more efficient and produce fewer nitrogen-containing by-products compared to air-based reactions.

Oxygen is also essential in environmental and energy-related industries. In wastewater treatment, dissolved oxygen levels are carefully controlled—typically between 2 and 4 mg/L—to support aerobic microorganisms that break down organic pollutants. In power generation, oxygen-enriched combustion and oxy-fuel technologies are used to improve efficiency and reduce emissions, particularly in carbon capture and storage (CCS) systems.

In aerospace and energy storage technologies, oxygen serves as a powerful oxidizer. Liquid oxygen (LOX), with a boiling point of −183 °C, has been used in rocket engines since the 1940s, enabling high thrust and efficient fuel combustion. Across modern industry, oxygen remains indispensable for increasing productivity, improving energy efficiency, and supporting advanced technologies.

We know that oxygen is produced by plants, but even though plants don't grow in the winter, why doesn't oxygen run out? First, not all parts of the world experience winter, so oxygen is definitely being released in other warmer places.

Secondly, there are trees that grow in the winter, such as conifers like pine, which also release oxygen in the winter. Finally, in the oceans, at the bottom of frozen seas and lakes, various plants grow (for example, algae), which also contribute significantly to the formation of air.

However, due to the rapid exchange of air masses on planet Earth, oxygen is always present in sufficient quantities everywhere. This gas makes up 20-21% of the volume of the Earth's atmosphere, and this figure is constant in any corner of the Earth.

Furthermore, there might be a slightly higher concentration of oxygen in the winter because cold air is naturally denser. A drop in bicycle tire pressure in the winter illustrates how cold temperatures cause air to become denser, reducing its volume or pressure in a fixed container.

The combustion process occurs when various substances combine with oxygen. Oxygen is the second most active non-metal after fluorine, so it undergoes a very active chemical reaction with any substance.

If many substances ignite easily in regular air (20-21% oxygen), their combustion will be much more intense in pure or saturated oxygen (50-100%). Substances that burn poorly or not at all in normal air, such as iron, will ignite when exposed to pure oxygen.

But if we use liquid oxygen, it burns even faster. If we mix liquid oxygen with powder-like substances (sawdust, moss, peat) and ignite them, they will explode like a regular bomb because they burn so quickly.

The composition of oxygen in the Earth's crust (by weight) is almost 50%. It is the third most abundant element in the universe, after hydrogen and helium. In the human body, oxygen makes up 65%.

The human body is adapted to an oxygen saturation of 20-21 percent at atmospheric pressure. If the saturation drops by a few percent, a person feels unwell. At 17% saturation, a person quickly becomes tired and dizzy, at 13% saturation, they lose consciousness, and at 7%, they die.

What happens if you use highly saturated oxygen? In some cases, it is useful. For example, in medicine, it is used to treat asthma, tuberculosis, poisoning with poisonous gases, and some other diseases.

However, breathing highly concentrated oxygen is harmful; very high concentrations can lead to high blood pressure and oxygen toxicity. Divers were the first to suffer from this problem.

They suffered from severe pain, vomiting, and twitching limbs as they breathed from high-pressure oxygen tanks. Following these incidents, the oxygen concentration in divers' breathing apparatus was reduced.

• This element is essential to human life and, despite being all around us, was only discovered in the 18th century.
• It causes substances to burn but does not burn itself.
• Trees and plants on Earth produce only half of the air on our planet. The other half, according to scientists, is produced by algae and phytoplankton in the world's oceans, which are capable of photosynthesis.
• Oxygen accounts for about two-thirds of the weight of most living things. This is because living organisms, including humans, are largely composed of water, and oxygen constitutes 88.9% of water by weight.
• An average passenger plane consumes 50-75 tons of oxygen during a 9-hour flight. At the same time, 25,000-50,000 hectares of forest produce the same amount of oxygen.
• All plants on Earth produce 3 trillion tons of oxygen per year.
• One tree produces an average of 125 kg of oxygen per year. This is enough to provide two people with clean air.
• The human brain can survive for 4-6 minutes without oxygen, after which it dies.
• Due to the abundance of oxygen during the Carboniferous period, insects of that time were gigantic.
• The human brain alone consumes about 20% of the body's total oxygen intake.

Oxygen. The word itself is simple, yet it describes the single most critical element sustaining complex life on Earth. It is an odorless, colorless gas, an invisible necessity that permeates our atmosphere and fuels the engine of every aerobic organism, from the smallest bacterium to the largest whale. To contemplate a world without oxygen is to contemplate a world utterly devoid of the life we know.

The most immediate and fundamental role of oxygen lies in cellular respiration. This biochemical process is the primary mechanism by which living cells extract energy from food. In the mitochondria, the cell's powerhouses, oxygen acts as the final electron acceptor in the electron transport chain. This sophisticated reaction generates adenosine triphosphate (ATP), the universal energy currency of the cell. Without oxygen, this highly efficient energy production system collapses, and life can only rely on the far less potent process of anaerobic respiration, which is insufficient for the demands of multicellular life.

Beyond the internal chemistry of our bodies, oxygen is a planetary architect. The massive shift in atmospheric composition known as the Great Oxidation Event, occurring billions of years ago, was a pivotal moment. The proliferation of photosynthetic organisms—chiefly cyanobacteria—began flooding the atmosphere with oxygen. This toxic byproduct to early anaerobic life ultimately paved the way for the evolution of more complex, oxygen-dependent organisms. Today, the balance of the ecosystem, maintained by the delicate cycle of photosynthesis (producing oxygen) and respiration (consuming oxygen), is a testament to its enduring influence. Forests, oceans, and even microscopic algae are our vital, ongoing oxygen factories.

Furthermore, oxygen is crucial for combustion and industrial processes. While our focus is often biological, oxygen's reactivity is harnessed by humanity daily. It is essential for generating heat and power, driving most engines, and producing steel and other vital materials. Even the protective ozone layer, which shields the Earth from harmful solar ultraviolet radiation, is a molecule formed from three oxygen atoms (O3). This layer is a natural sunscreen, safeguarding the very life that consumes O2.

In conclusion, oxygen is more than just a breathable gas; it is the linchpin of terrestrial existence. It is the final spark in our cellular energy process, the geological force that shaped our atmosphere, and the environmental buffer protecting us from space. Its necessity is absolute and unceasing. As we take our next breath, we perform the most basic, yet most profound, act of connection to the planet's vast, invisible life-support system. Without oxygen, the flicker of life instantly extinguishes.