Discovery, Symbol, Properties,
Uses, & Facts
Oxygen
Discovery, Symbol, Properties, Uses, & Facts
Element Properties
atomic number 8
atomic weight 15.9994
melting point −218.4 °C (−361.1 °F)
boiling point −183.0 °C (−297.4 °F)
density (1 atm, 0 °C) 1.429 g/litre
oxidation states −1, −2, +2 (in compounds with fluorine)
electron config. 1s22s22p4
History
Oxygen was discovered about 1772 by a Swedish chemist, Carl Wilhelm Scheele, who obtained it by heating potassium nitrate, mercuric oxide, and many other substances. An English chemist, Joseph Priestley, independently discovered oxygen in 1774 by the thermal decomposition of mercuric oxide and published his findings the same year, three years before Scheele published. In 1775–80, French chemist Antoine-Laurent Lavoisier, with remarkable insight, interpreted the role of oxygen in respiration as well as combustion, discarding the phlogiston theory, which had been accepted up to that time; he noted its tendency to form acids by combining with many different substances and accordingly named the element oxygen (oxygène) from the Greek words for “acid former.”
Occurrence and properties
At 46 percent of the mass, oxygen is the most plentiful element in Earth’s crust. The proportion of oxygen by volume in the atmosphere is 21 percent and by weight in seawater is 89 percent. In rocks, it is combined with metals and nonmetals in the form of oxides that are acidic (such as those of sulfur, carbon, aluminum, and phosphorus) or basic (such as those of calcium, magnesium, and iron) and as saltlike compounds that may be regarded as formed from the acidic and basic oxides, as sulfates, carbonates, silicates, aluminates, and phosphates. Plentiful as they are, these solid compounds are not useful as sources of oxygen, because separation of the element from its tight combinations with the metal atoms is too expensive.
Below −183 °C (−297 °F), oxygen is a pale blue liquid; it becomes solid at about −218 °C (−361 °F). Pure oxygen is 1.1 times heavier than air.
During respiration, animals and some bacteria take oxygen from the atmosphere and return to it carbon dioxide, whereas by photosynthesis, green plants assimilate carbon dioxide in the presence of sunlight and evolve free oxygen. Almost all the free oxygen in the atmosphere is due to photosynthesis. About 3 parts of oxygen by volume dissolve in 100 parts of fresh water at 20 °C (68 °F), slightly less in seawater. Dissolved oxygen is essential for the respiration of fish and other marine life.
Natural oxygen is a mixture of three stable isotopes: oxygen-16 (99.759 percent), oxygen-17 (0.037 percent), and oxygen-18 (0.204 percent). Several artificially prepared radioactive isotopes are known. The longest-lived, oxygen-15 (124-second half-life), has been used to study respiration in mammals.
Allotropy
Oxygen has two allotropic forms, diatomic (O2) and triatomic (O3, ozone). The properties of the diatomic form suggest that six electrons bond the atoms and two electrons remain unpaired, accounting for the paramagnetism of oxygen. The three atoms in the ozone molecule do not lie along a straight line.
Ozone may be produced from oxygen according to the equation:
The process, as written, is endothermic (energy must be provided to make it proceed); conversion of ozone back into diatomic oxygen is promoted by the presence of transition metals or their oxides. Pure oxygen is partly transformed into ozone by a silent electrical discharge; the reaction is also brought about by absorption of ultraviolet light of wavelengths around 250 nanometres (nm, the nanometre, equal to 10−9 metre); occurrence of this process in the upper atmosphere removes radiation that would be harmful to life on the surface of the Earth. The pungent odour of ozone is noticeable in confined areas in which there is sparking of electrical equipment, as in generator rooms. Ozone is light blue; its density is 1.658 times that of air, and it has a boiling point of −112 °C (−170 °F) at atmospheric pressure.
Ozone is a powerful oxidizing agent, capable of converting sulfur dioxide to sulfur trioxide, sulfides to sulfates, iodides to iodine (providing an analytical method for its estimation), and many organic compounds to oxygenated derivatives such as aldehydes and acids. The conversion by ozone of hydrocarbons from automotive exhaust gases to these acids and aldehydes contributes to the irritating nature of smog. Commercially, ozone has been used as a chemical reagent, as a disinfectant, in sewage treatment, water purification, and bleaching textiles.
Preparative methods
Production methods chosen for oxygen depend upon the quantity of the element desired. Laboratory procedures include the following:
1. Thermal decomposition of certain salts, such as potassium chlorate or potassium nitrate:
The decomposition of potassium chlorate is catalyzed by oxides of transition metals; manganese dioxide (pyrolusite, MnO2) is frequently used. The temperature necessary to effect the evolution of oxygen is reduced from 400 °C to 250 °C by the catalyst.
2. Thermal decomposition of oxides of heavy metals
Scheele and Priestley used mercury(II) oxide in their preparations of oxygen.
3. Thermal decomposition of metal peroxides or of hydrogen peroxide:
An early commercial procedure for isolating oxygen from the atmosphere or for manufacture of hydrogen peroxide depended on the formation of barium peroxide from the oxide as shown in the equations.
4. Electrolysis of water containing small proportions of salts or acids to allow conduction of the electric current:</span
Commercial production and use
When required in tonnage quantities, oxygen is prepared by the fractional distillation of liquid air. Of the main components of air, oxygen has the highest boiling point and therefore is less volatile than nitrogen and argon. The process takes advantage of the fact that when a compressed gas is allowed to expand, it cools. Major steps in the operation include the following: (1) Air is filtered to remove particulates; (2) moisture and carbon dioxide are removed by absorption in alkali; (3) the air is compressed and the heat of compression removed by ordinary cooling procedures; (4) the compressed and cooled air is passed into coils contained in a chamber; (5) a portion of the compressed air (at about 200 atmospheres pressure) is allowed to expand in the chamber, cooling the coils; (6) the expanded gas is returned to the compressor with multiple subsequent expansion and compression steps resulting finally in liquefaction of the compressed air at a temperature of −196 °C; (7) the liquid air is allowed to warm to distill first the light rare gases, then the nitrogen, leaving liquid oxygen. Multiple fractionations will produce a product pure enough (99.5 percent) for most industrial purposes.
The steel industry is the largest consumer of pure oxygen in “blowing” high carbon steel—that is, volatilizing carbon dioxide and other nonmetal impurities in a more rapid and more easily controlled process than if air were used. The treatment of sewage by oxygen holds promise for more efficient treatment of liquid effluents than other chemical processes. Incineration of wastes in closed systems using pure oxygen has become important. The so-called LOX of rocket oxidizer fuels is liquid oxygen; the consumption of LOX depends upon the activity of space programs. Pure oxygen is used in submarines and diving bells.
Commercial oxygen or oxygen-enriched air has replaced ordinary air in the chemical industry for the manufacture of such oxidation-controlled chemicals as acetylene, ethylene oxide, and methanol. Medical applications of oxygen include use in oxygen tents, inhalators, and pediatric incubators. Oxygen-enriched gaseous anesthetics ensure life support during general anesthesia. Oxygen is significant in a number of industries that use kilns.
Chemical properties and reactions
The large values of the electronegativity and the electron affinity of oxygen are typical of elements that show only nonmetallic behaviour. In all of its compounds, oxygen assumes a negative oxidation state as is expected from the two half-filled outer orbitals. When these orbitals are filled by electron transfer, the oxide ion O2− is created. In peroxides (species containing the ion O22−) it is assumed that each oxygen has a charge of −1. This property of accepting electrons by complete or partial transfer defines an oxidizing agent. When such an agent reacts with an electron-donating substance, its own oxidation state is lowered. The change (lowering), from the zero to the −2 state in the case of oxygen, is called a reduction. Oxygen may be thought of as the “original” oxidizing agent, the nomenclature used to describe oxidation and reduction being based upon this behaviour typical of oxygen.
As described in the section on allotropy, oxygen forms the diatomic species, O2, under normal conditions and, as well, the triatomic species ozone, O3. There is some evidence for a very unstable tetratomic species, O4. In the molecular diatomic form there are two unpaired electrons that lie in antibonding orbitals. The paramagnetic behaviour of oxygen confirms the presence of such electrons.
The intense reactivity of ozone is sometimes explained by suggesting that one of the three oxygen atoms is in an “atomic” state; on reacting, this atom is dissociated from the O3 molecule, leaving molecular oxygen.
The molecular species, O2, is not especially reactive at normal (ambient) temperatures and pressures. The atomic species, O, is far more reactive. The energy of dissociation (O2 → 2O) is large at 117.2 kilocalories per mole.
Oxygen has an oxidation state of −2 in most of its compounds. It forms a large range of covalently bonded compounds, among which are oxides of nonmetals, such as water (H2O), sulfur dioxide (SO2), and carbon dioxide (CO2); organic compounds such as alcohols, aldehydes, and carboxylic acids; common acids such as sulfuric (H2SO4), carbonic (H2CO3), and nitric (HNO3); and corresponding salts, such as sodium sulfate (Na2SO4), sodium carbonate (Na2CO3), and sodium nitrate (NaNO3). Oxygen is present as the oxide ion, O2-, in the crystalline structure of solid metallic oxides such as calcium oxide, CaO. Metallic superoxides, such as potassium superoxide, KO2, contain the O2- ion, whereas metallic peroxides, such as barium peroxide, BaO2, contain the O22- ion.
The Origin of Oxygen in Earth's Atmosphere
The breathable air we enjoy today originated from tiny organisms, although the details remain lost in geologic time
It’s hard to keep oxygen molecules around, despite the fact that it’s the third-most abundant element in the universe, forged in the superhot, superdense core of stars. That’s because oxygen wants to react; it can form compounds with nearly every other element on the periodic table. So how did Earth end up with an atmosphere made up of roughly 21 percent of the stuff?
The answer is tiny organisms known as cyanobacteria, or blue-green algae. These microbes conduct photosynthesis: using sunshine, water and carbon dioxide to produce carbohydrates and, yes, oxygen. In fact, all the plants on Earth incorporate symbiotic cyanobacteria (known as chloroplasts) to do their photosynthesis for them down to this day.
For some untold eons prior to the evolution of these cyanobacteria, during the Archean eon, more primitive microbes lived the real old-fashioned way: anaerobically. These ancient organisms—and their “extremophile” descendants today—thrived in the absence of oxygen, relying on sulfate for their energy needs.
But roughly 2.45 billion years ago, the isotopic ratio of sulfur transformed, indicating that for the first time oxygen was becoming a significant component of Earth’s atmosphere, according to a 2000 paper in Science. At roughly the same time (and for eons thereafter), oxidized iron began to appear in ancient soils and bands of iron were deposited on the seafloor, a product of reactions with oxygen in the seawater.
“What it looks like is that oxygen was first produced somewhere around 2.7 billion to 2.8 billon years ago. It took up residence in atmosphere around 2.45 billion years ago,” says geochemist Dick Holland, a visiting scholar at the University of Pennsylvania. “It looks as if there’s a significant time interval between the appearance of oxygen-producing organisms and the actual oxygenation of the atmosphere.”
So a date and a culprit can be fixed for what scientists refer to as the Great Oxidation Event, but mysteries remain. What occurred 2.45 billion years ago that enabled cyanobacteria to take over? What were oxygen levels at that time? Why did it take another one billion years—dubbed the “boring billion” by scientists—for oxygen levels to rise high enough to enable the evolution of animals?
Most important, how did the amount of atmospheric oxygen reach its present level? “It’s not that easy why it should balance at 21 percent rather than 10 or 40 percent,” notes geoscientist James Kasting of Pennsylvania State University. “We don’t understand the modern oxygen control system that well.”
Climate, volcanism, plate tectonics all played a key role in regulating the oxygen level during various time periods. Yet no one has come up with a rock-solid test to determine the precise oxygen content of the atmosphere at any given time from the geologic record. But one thing is clear—the origins of oxygen in Earth’s atmosphere derive from one thing: life.
Discovery, Symbol, Properties, Uses & Facts
Air (National Geographic Society) is the invisible mixture of gases that surrounds the Earth. Air contains important substances, such as oxygen and nitrogen, that most species need to survive. Human beings, of course, are one of those species. Sometimes, the word “atmosphere” is used instead of the word “air.”
Standard Dry Air is the composition of gases that make up air at sea level. It is a standard scientific unit of measurement. Standard Dry Air is made up of nitrogen, oxygen, argon, carbon dioxide, neon, helium, krypton, hydrogen, and xenon. It does not include water vapor because the amount of vapor changes based on humidity and temperature. Because air masses are constantly moving, Standard Dry Air is not accurate everywhere at once.
Nitrogen and oxygen make up about 99 percent of Earth’s air. People and other animals need oxygen to live. Carbon dioxide, a gas that plants depend on, makes up less than .04 percent.
Plants and animals each produce the gases that the other needs to live. Plants need carbon dioxide—people and other animals exhale carbon dioxide as a waste product. People and other animals need oxygen—plants produce oxygen during an important process called photosynthesis, which turns the sun’s energy into nutrients.
Water vapor in the air is sometimes visible as clouds. Water enters the atmosphere through the water cycle. The water cycle also brings molecules in the air into oceans, lakes, and rivers.
Some gases in the air come from volcanic eruptions. Volcanic eruptions eject gases from the interior of the Earth. The most common gas emitted by volcanoes is water vapor. Other gases, such as carbon monoxide and sulfur dioxide, are toxic to most organisms. A few organisms, however, thrive on these gases. At the bottom of the ocean are bacteria that do not need oxygen or sunlight to survive. In other words, they do not need air. These strange organisms create their own nutrients using hydrogen sulfide, not carbon dioxide. The hydrogen sulfide comes from cracks, or vents, in the Earth’s crust.
The air is different as you move higher and higher into the atmosphere. The air gets “thinner” as elevation climbs because there are fewer air molecules up there. Mountain climbers often have to use canisters of oxygen as they climb above 3,800 meters (12,500 feet) because there is not enough oxygen in the atmosphere for most people to breathe. High mountains such as Mount Everest (8,848 meters, or 29,035 feet), in Nepal and China, are littered with empty oxygen canisters that climbers discard when they are used up.
High in the stratosphere, a layer of the Earth’s atmosphere, is a special air molecule called ozone. Ozone is made up of three atoms of oxygen. The massive collection of these molecules is called the ozone layer. The ozone layer blocks harmful ultraviolet, or UV, rays so the sun’s powerful radiation does less damage to living things on Earth.
Unfortunately, air pollution has a negative effect on the air we breathe. Air pollution happens when harmful byproducts, like exhaust from cars, enter the air. These pollutants can clog the atmosphere with smog, a combination of smoke and fog. They can also create toxic clouds of dust. Other air pollutants, such as methane and excess amounts of carbon dioxide, can upset the balance of molecules in the air, contributing to global warming.
Non-Atmospheric Air
Compressed air is air kept at a consistent pressure, such as air pressure at sea level. Airplanes are usually pressurized at ground level so passengers can breathe without canisters of air.
Often, compressed air is kept at a pressure higher than normal air pressure. Scuba divers use compressed air to breathe under water. The canisters of air allow divers to inhale through a tube and exhale into the water.
Pneumatics is the science and work of pressurized air and other gases. The uses for pneumatics are far-ranging. Air brakes in automobiles, trucks, and trains use compressed air to slow wheel rotation and stop the vehicles. Pipe organs use compressed air at different pressures to create different musical notes.