Oxygen and its compounds play a key role in many of the important processes of life and industry. Oxygen in the biosphere is essential in the processes of respiration and metabolism, the means by which animals derive the energy needed to sustain life. Furthermore, oxygen is the most abundant element at the surface of the Earth. In combined form it is found in ores, earths, rocks, and gemstones, as well as in all living organisms.

Oxygen is a gaseous chemical element in Group VA of the periodic table. The chemical symbol for atomic oxygen is O, its atomic number is 8, and its atomic weight is 15.9994. Elemental oxygen is known principally in the gaseous form as the diatomic molecule, which makes up 20.95% of the volume of dry air. Diatomic oxygen is colorless, odorless, and tasteless.

Two 18th-century scientists share the credit for first isolating elemental oxygen: Joseph PRIESTLEY (1733-1804), an English clergyman who was employed as a literary companion to Lord Shelburne at the time of his most significant experimental work, and Carl Wilhelm SCHEELE (1742-86), a Swedish pharmacist and chemist. It is generally believed that Scheele was the first to isolate oxygen, but that Priestley, who independently achieved the isolation of oxygen somewhat later, was the first to publicly announce his findings.

The interpretation of the findings of Priestley and the resultant clarification of the nature of oxygen as an element was accomplished by the French scientist Antoine-Laurent LAVOISIER (1743-94). Lavoisier’s experimental work, which extended and improved upon Priestley’s experiments, was principally responsible for the understanding of COMBUSTION and the establishment of the law of conservation of matter.

Lavoisier gave oxygen its name, which is derived from two Greek words that mean “acid former.” Lavoisier held the mistaken idea that oxides, when dissolved in water, would form only acids. It is true that some oxides when dissolved in water do form acids; for example, sulfur dioxide forms sulfurous acid. Some oxides, however, such as sodium oxide, dissolve in water to form bases, as in the reaction to form sodium hydroxide; therefore oxygen was actually inappropriately named.


Oxygen is formed by a number of nuclear processes that are believed to occur in stellar interiors. The most abundant isotope of oxygen, with mass 16, is thought to be formed in hydrogen-burning stars by the capture of a proton by the isotopes of nitrogen and fluorine, with the subsequent emission of, respectively, a gamma ray and an alpha particle. In helium-burning stars the isotope of carbon with mass 12 is thought to capture an alpha particle to form the isotope with mass 16 with the emission of a gamma ray.

In the terrestrial environment oxygen accounts for about half of the mass of the Earth’s crust, 89% of the mass of the oceans, and 23% of the mass (and 21% of the volume) of the atmosphere. Most of the Earth’s rocks and soils are principally silicates. The silicates are an amazingly complex group of materials that typically consist of greater than 50 (atomic) percent oxygen in combination with silicon and one or more metallic elements.

Several important ores are principally oxides of the desired metals, such as the important iron-bearing minerals hematite, magnetite, and limonite and the most important aluminum-bearing mineral, BAUXITE (a mixture of hydrated aluminum oxides and iron oxide).


Three naturally occurring isotopes of oxygen have been found: one with mass 16 (99. 759% of all oxygen), one with mass 17 (0.037%); and one with mass 18 (0.204%). The rarer isotopes, principally the latter, find their major use in labeling experiments used by scientists to follow the steps of chemical reactions.

If oxygen at a pressure of one atmosphere is cooled, it will liquefy at 90.18 K (-182.97 deg C; -297.35 deg F), the normal boiling point of oxygen, and it will solidify at 54.39 K (-218.76 deg C; -361.77 deg F), the normal melting point of oxygen. The liquid and solid forms of oxygen have a pale blue color. Several different structures are known for solid oxygen: solid type III, from the lowest temperatures achievable to 23.66 K; type II, from 23.66 to 43.76 K; and type I, from 43.76 to 54.39 K. The critical temperature for oxygen, the temperature above which it is impossible to liquefy the gas no matter how much pressure is applied, is 154.3 K (-118.9 deg C; -181.9 deg F). The pressure of liquid and gaseous oxygen coexisting in equilibrium at the critical temperature is 49.7 atmospheres.

Oxygen gas exhibits a slight but important solubility in water. Molecular oxygen dissolved in water is required by aquatic organisms for their metabolic processes and is ultimately responsible for the oxidation and removal of organic wastes in water. The solubilities of gases depend on the temperature of the solution and the pressure of the gas over the solution. At 20 deg C (68 deg F) and an oxygen pressure of one atmosphere, the solubility of O(2) in water is about 45 grams of oxygen per cubic meter of water, or 45 ppm (parts per million).

Molecular diatomic oxygen is a fairly stable molecule requiring a dissociation energy (the energy required to dissociate one mole of molecular oxygen in its ground state into two moles of atomic oxygen in its ground state) of 493.6 kilojoules per mole. The molecule is dissociated by ultraviolet radiation of any wavelength shorter than 193 nm. Solar radiation striking stratospheric oxygen dissociates it into atomic oxygen for this reason. The atomic oxygen formed in this fashion is capable of reacting with oxygen to form OZONE.


Many direct, uncatalyzed reactions of oxygen do not occur rapidly at room temperature. This fact has a number of important consequences. One of these consequences has to do with the use of metals as structural materials. Metals that are used in construction, such as iron (principally as steel) and aluminum, form highly stable oxides. For example, the oxidation of aluminum has a significant tendency to occur. However, in spite of this tendency, the reaction occurs so slowly at room temperature that it can be said for most practical purposes not to occur at all, and for this reason aluminum is an appropriate and widely used structural material. The slowness of this reaction is due in part to the stability of the oxygen-oxygen bond and in part because of a very thin, protective layer of oxide that forms on the surface of the aluminum. The oxidation of iron is a complex process involving impurities in the iron, as well as water and carbon dioxide. This oxidative destruction, or rusting, of iron and steel–which are among our most important structural materials–is extremely costly to modern societies.

Biological Oxidation

Another important aspect of the rates of oxygen reactions concerns the rate of reaction with organic materials. Such oxidation reactions are, ultimately, the sources of energy for the higher plants and animals, are responsible for the cleansing of streams of biodegradable wastes, and are responsible for the natural decomposition of organic material. The rates of reactions in this category are selectively controlled by enzymes in the organisms that facilitate the reactions. Thus waste products and dead plants and animals decompose (are oxidized) principally through the agency of microorganisms, and energy-bearing foods are metabolized (oxidized) by means of biological processes.


There is a marked difference between the rates of reactions with oxygen at room temperature and the rates at elevated temperatures. Many substances that do not react rapidly with oxygen in air at temperatures below 100 deg C will do so at 1000 deg C with a strong evolution of heat (exothermically). For example, coal and petroleum can be stored indefinitely at the temperatures encountered under normal climatic conditions, but they readily oxidize, exothermically, at elevated temperatures.

The most common compounds of oxygen are those in which the element exhibits a valence of two. This fact is associated with the electronic structure of atomic oxygen; this atom requires two additional electrons to fill its outermost energy level. Examples of divalent oxides are numerous among well-known substances such as water; carbon dioxide; aluminum oxide; silicon dioxide; the silicates, calcium carbonate or limestone; and sulfur dioxide. Oxygen is also known to have other valences, such as in the PEROXIDES, of which hydrogen peroxide is an example.

The direct reaction of oxygen with another element frequently follows the pattern discussed above; that is, it does not occur rapidly or at all at room temperature but is strongly exothermic, and once oxidation is initiated the evolved heat raises the temperature of the reactants such that the reaction is self-sustaining. Examples of such reactions are with the elements magnesium, carbon, and hydrogen. Magnesium and carbon burn in air once the reaction is initiated, and a hydrogen-oxygen mixture can react explosively when the reaction is initiated by a flame or spark. The explosion of a hydrogen-oxygen mixture is an extremely fast reaction and occurs because of the formation of atomic oxygen in the exploding mixture.


Pure oxygen is used extensively in technological processes. It is used in the welding, cutting, and forming of metals, as in oxyacetylene welding, in which oxygen reacts with acetylene to form an extremely hot flame. Oxygen is added to the inlet air (3 to 5%) in modern blast furnaces to increase the temperature in the furnace; it is also used in the basic oxygen converter for steel production, in the manufacture of chemicals, and for rocket propulsion.

Oxygen is also used in the partial combustion of methane (natural gas) or coal (taken here to be carbon) to make mixtures of carbon monoxide and hydrogen called synthesis gas, which is in turn used for the manufacture of methanol. Processes in which combustible liquids are produced from coal will become increasingly important as petroleum resources become further depleted.


Oxygen is conveniently produced in the laboratory by heating mercuric oxide or potassium chlorate to moderately high temperatures. The production from mercuric oxide is the method that was employed by Joseph Priestley, and the production from the potassium chlorate method commonly used by students in today’s laboratories. Oxygen is liberated when solid potassium chlorate is heated to 400 deg C or, when manganese dioxide is added as a CATALYST, to 200 deg C. The liberated oxygen can be collected by water displacement because of the low solubility of oxygen in water.

Oxygen can also be produced in the laboratory by the electrolysis of water, a process that reverses the violent hydrogen-oxygen reaction discussed previously. When a current is passed through water the liquid is decomposed at the electrodes. This method is also used to produce oxygen on a commercial scale when a high-purity product is desired.

The more economical, and therefore preferred, method for the commercial production of oxygen is the liquefaction and distillation of air. The air is cooled until it liquefies, principally by being made to do work in a rotating expansion turbine, and the resulting liquid air is fractionated by a complex distillation process. The gaseous oxygen produced in this fashion is shipped in pressurized cylinders or, as is often the case when large amounts are involved, through pipelines to nearby industrial plants.


Most organisms depend on oxygen to sustain their biological processes. The great majority of living organisms fall into two categories. In the first category are the higher plants and the photosynthetic bacteria. These organisms utilize light energy through PHOTOSYNTHESIS to combine carbon dioxide and water (or, infrequently, other inorganic substances in place of water) into more complex materials characterized as CARBOHYDRATES while at the same time releasing oxygen into the atmosphere. In the second category are the higher animals, most microorganisms, and photosynthetic cells that live in the dark. All these second-category organisms use complex series of enzyme-catalyzed OXIDATION AND REDUCTION reactions using materials such as glucose as the fuel and oxygen as the terminal oxidizing agent (see METABOLISM). The end products of metabolism in these organisms are carbon dioxide and water, which are returned to the atmosphere.

The net result of these complementary functions is the oxygen cycle, in which the photosynthetic organisms, using solar energy, synthesize carbohydrates from water and carbon dioxide and give off oxygen as a by-product, while the aerobic organisms oxidize ingested organic materials, using up oxygen and giving off carbon dioxide and water through a complex series of metabolic processes. It has been estimated that 3.5 X (10 to the power of 11) tons of carbon dioxide are cycled annually via these processes.

Thus, in the vertebrates–and in humans in particular–oxygen is necessary to sustain metabolism and thus life. Air is inhaled and oxygen in the air is exchanged in the lungs between the atmosphere and the hemoglobin in the blood. The blood carries the oxygen, complexed with hemoglobin, to all parts of the body in which metabolic processes occur. It also carries carbon dioxide back to the lungs, where the carbon dioxide is exchanged with the atmosphere and exhaled.

If the oxygen concentration were to drop to about half its value in the atmosphere, humans could no longer survive. For this reason an important component of the life-support systems of divers and astronauts is a source of oxygen gas. Similarly, persons ill with respiratory diseases that interfere with normal respiration, such as pneumonia and emphysema, are often kept in OXYGEN TENTS and HYPERBARIC CHAMBERS, the latter administering high-pressure oxygen, may be used to treat a variety of ailments.


In the realm of inorganic chemistry there is a very large number of oxygen-containing compounds. There are very few elements for which no oxides are known, and there are several metallic elements (such as titanium, vanadium, and praseodymium) for which a wide variety of solid oxides exist. The solid oxides of the metallic elements can generally be synthesized by the direct reaction of the elements at high temperatures. In many cases such reactions will result in the formation of a single oxide of the metal in its most oxidized form. Typical examples are the metallic oxides of sodium, calcium, lanthanum, titanium, vanadium, and tungsten.

In the cases of elements capable of forming reduced oxides, in particular the early transition metals, the reduced oxides can be formed by heating the highest oxide, formed as above, to very high temperatures (1,500 K or higher) either in an inert container or in the presence of the metallic element. The reduced oxides that result exhibit a variation in the extent and importance of direct metal-metal bonding in the compounds, and this variation gives rise to a variety of electrical and magnetic properties. The more metal-rich of these oxides are metallic conductors and tend to be nonstoichiometric; that is, they are observed to exist over a range of compositions all possessing the same underlying structure. A number of these titanium oxides exhibit more than one crystal structure (polymorphism). The most oxidized compound, titanium, is widely used in the RUTILE form as a white pigment in paints.

Ternary oxides, consisting of two metallic elements plus oxygen, are of great interest to solid-state scientists. For example, compounds such as the SPINELS and the PEROVSKITES are studied extensively because of their interesting magnetic and electrical properties. Examples of important ternary oxides are the magnetic FERRITES, whose magnetic properties can be tailored, making them useful in computer memory units. The ferrites are prepared by firing compacted mixtures of iron oxide and one or more metal oxides (such as those of nickel, copper, zinc, magnesium, and manganese).

Also of importance in inorganic chemistry are the oxides of the nonmetals. Most of the nonmetals are known to form a wide variety of compounds with oxygen. The nitrogen oxides are undesirable by-products of high-temperature combustion in air (as in an internal combustion engine) and can cause serious environmental pollution.

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