OXYGEN (REVISED)

Note: This article, originally published in 1998, was updated in 2006 for the eBook edition.

Overview

Oxygen is the first element in Group 16 (VIA) of the periodic table. The periodic table is a chart that shows how chemical elements are related to each other. The elements in Group 16 are said to belong to the chalcogen family. Other elements in this group include sulfur, selenium, tellurium, and polonium. The name chalcogen comes from the Greek word chalkos, meaning "ore." The first two members of the family, oxygen and sulfur, are found in most ores.

Oxygen is by far the most abundant element in the Earth's crust. Nearly half of all the atoms in the earth are oxygen atoms. Oxygen also makes up about one-fifth of the Earth's atmosphere. Nearly 90 percent of the weight of the oceans is due to oxygen. In addition, oxygen is thought to be the third most abundant element in the universe and in the solar system.

The discovery of oxygen is usually credited to Swedish chemist Carl Wilhelm Scheele (1742-86) and English chemist Joseph Priestley (1733-1804). The two discovered oxygen at nearly the same time in 1774, working independently of each other.

SYMBOL
0

ATOMIC NUMBER
8

ATOMIC MASS
15.9994

FAMILY
Group 16 (VIA)
Chalcogen

PRONUNCIATION
OK-si-jun

Oxygen is necessary for the survival of all animal life on Earth. Animals breathe in oxygen and breathe out carbon dioxide.

One important use of oxygen is in medicine. People who have trouble breathing are given extra doses of oxygen. In many cases, this "extra oxygen" keeps people alive after they would otherwise have died.

But oxygen has many commercial uses also. The most important use is in the manufacture of metals. More than half of the oxygen produced in the United States is used for this purpose. Oxygen usually ranks third in the list of chemicals produced in the United States each year. In 1996, about 668 billion cubic feet of oxygen was manufactured in the United States. The gas is prepared almost entirely from liquid air.

Discovery and naming

What is air? Ancient peoples thought deeply about that question. And that should not be surprising. It is easy to see how essential air is to many processes. Objects cannot burn without air. Human life cannot survive without air. In fact, ancient peoples thought air must be an "element." But they used the word "element" differently than do modern scientists. To ancient people, an element was something that was very important and basic. Air fit that description, along with fire, water, and earth.

They often thought of air as an element in the modern sense—that it was as simple a material as could be found. Yet, some early scholars believed otherwise. For example, some Chinese scholars, as early as the eighth century a.d., thought of air as having two parts. They called these parts the yin and yang of air. The properties of the Chinese yin and yang can be compared to the properties of oxygen and nitrogen.

The first person in Western Europe to describe the "parts" of air was Italian artist and scientist Leonardo da Vinci (1452-1519). Leonardo pointed out that air is not entirely used up when something is burned in it. He said that air must consist, therefore, of two parts: one part that is consumed in burning and one part that is not.

For many years, Leonardo's ideas were not very popular among scholars. One problem was that early chemists did not have very good equipment. It was difficult for them to collect samples of air and then to study it.

In the early 1700s, chemists began to find out more about air, but in a somewhat roundabout way. For example, in 1771 and 1772, Scheele studied the effect of heat on a number of different compounds. In one experiment, he used silver carbonate (Ag₂CO₃), mercury carbonate (HgCO₃), and magnesium nitrate (Mg(NO₃)₂). When he heated these compounds, he found that a gas was produced. He then studied the properties of that gas. He found that flames burned brightly in the gas. He also found that animals could live when placed in the gas. Without knowing it, Scheele had discovered oxygen. (See sidebar on Scheele in the chlorine entry in Volume 1.)

About two years later, Priestley conducted similar experiments by heating mercury oxide (HgO) in a flame. The compound broke down, producing liquid mercury metal and a gas:

When Priestley tested the new gas, he found the same properties that Scheele had described.

Priestley even tried breathing the new gas he had produced. His description of that experience has now become famous:

Antoine-Laurent Lavoisier | French chemist

A ntoine-Laurent Lavoisier (1743-94) is often called the father of modern chemistry. He has been given that title for a number of reasons. The most important reason is the explanation he discovered for the process of combustion (burning).

Prior to Lavoisier's research, chemists thought that a burning object gave off a substance to the air. They called that substance phlogiston. When wood burned, for example, chemists said that phlogiston escaped from the wood to the air.

Lavoisier showed that this idea was incorrect. When something burns, it actually combines with oxygen in the air. Combustion, Lavoisier said, is really just oxidation (the process by which something combines with oxygen).

This discovery gave chemists a whole new way to look at chemical changes. The phlogiston theory gradually began to die out. Many of the ideas used in modern chemistry began to develop. No wonder Lavoisier is called the father of this revolution.

Lavoisier led an unusually interesting life. He was an avid chemist who carried out many experiments. But he also had a regular job as a tax collector. His job was to visit homes and businesses and collect taxes. This did not make him a very popular man!

Lavoisier also made some important enemies early in his life. One of these enemies was Jean-Paul Marat (1743-93). Marat thought of himself as a scientist and applied for membership in the French Academy of Scientists. Lavoisier voted against Marat's application. He said that Marat's research was not very good.

Less than a decade later, Lavoisier had reason to regret that decision. Marat had become a leader in the French Revolution (1774-1815). He accused Lavoisier of plotting against the revolution. He also said that Lavoisier was carrying out dangerous secret experiments.

These accusations were not true. But Marat was now a very powerful man. He was able to have Lavoisier convicted of the charges against him. On May 8, 1794, Lavoisier was beheaded and buried in an unmarked grave. Some people have said that Lavoisier's death was the worst single consequence of the French Revolution.

The feeling of it [the new gas, oxygen] to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards. Who can tell but that, in time, this pure air may become a fashionable article in luxury? Hitherto only two mice and myself have had the privilege of breathing it.

Some people think Scheele should get credit for discovering oxygen. He completed his experiments earlier than did Priestley. But his publisher was very slow in printing Scheele's reports. They actually came out after Priestley's reports. So most historians agree that Scheele and Priestly should share credit for discovering oxygen.

Neither Scheele nor Priestley fully understood the importance of their discovery. That step was taken by French chemist Antoine-Laurent Lavoisier (1743-94). Lavoisier was the first person to declare that the new gas was an element. He was also the first person to explain how oxygen is involved in burning. In addition, he suggested a name for the gas. That name, oxygen, comes from Greek words that mean "acidic" (oxy-) and "forming" (-gen). Lavoisier chose the name because he thought that all acids contain oxygen. Therefore, the new element was responsible for "forming acids." In this one respect, however, Lavoisier was wrong. All acids do not contain oxygen, although some do.

Physical properties

Oxygen is a colorless, odorless, tasteless gas. It changes from a gas to a liquid at a temperature of -182.96°C (-297.33°F). The liquid formed has a slightly bluish color to it. Liquid oxygen can then be solidified or frozen at a temperature of -218.4°C (-361.2°F). The density of oxygen is 1.429 grams per liter. By comparison, the density of air is about 1.29 grams per liter.

Oxygen exists in three allotropic forms. Allotropes are forms of an element with different physical and chemical properties. The three allotropes of oxygen are normal oxygen, or diatomic oxygen, or dioxygen; nascent, atomic, or monatomic oxygen; and ozone, or triatomic oxygen. The three allotropes differ from each other in a number of ways.

First, they differ on the simplest level of atoms and molecules. The oxygen that we are most familiar with in the atmosphere has two atoms in every molecule. Chemists show this by writing the formula as O₂. The small "2" means "two atoms per molecule."

By comparison, nascent oxygen has only one atom per molecule. The formula is simply O, or sometimes (O). The parentheses indicate that nascent oxygen does not exist very long under normal conditions. It has a tendency to form dioxygen:

That is, dioxygen is the normal condition of oxygen at room temperature.

The third allotrope of oxygen, ozone, has three atoms in each molecule. The chemical formula is O₃. Like nascent oxygen, ozone does not exist for very long under normal conditions. It tends to break down and form dioxygen:

Ozone does occur in fairly large amounts under special conditions. For example, there is an unusually large amount of ozone in the Earth's upper atmosphere. That ozone layer is important to life on Earth. It shields out harmful radiation that comes from the Sun. Ozone is also sometimes found closer to the Earth's surface. I t is produced when gasoline is burned in cars and trucks. It is part of the condition known as air pollution. Ozone at ground level is not helpful to life, and may cause health problems for plants, humans, and other animals.

The physical properties of ozone are somewhat different from those of dioxygen. It has a slightly bluish color as both a gas and a liquid. It changes to a liquid at a temperature of -111.9°C (-169.4°F) and from a liquid to a solid at -193°C (-135°F) . The density is 2.144 grams per liter.

Chemical properties

Oxygen's most important chemical property is that it supports combustion. That is, it helps other objects to burn. The combustion (burning) of charcoal is an example. Charcoal is nearly pure carbon (C):

Oxygen also combines with elements at room temperature. Rusting is an example. Rusting is a process by which a metal combines with oxygen. When iron rusts, it combines with oxygen:

Oxygen also reacts with many compounds. Decay is an example. Decay is the process by which once-living material combines with oxygen. The products of decay are mainly carbon dioxide (CO₂) and water (H₂O):

(The chemical formula for "dead matter" is too complicated to use here.)

Oxygen itself does not burn. A lighted match in a container of pure oxygen burns much brighter, but the oxygen does not catch fire.

Occurrence in nature

Oxygen occurs mainly as an element in the atmosphere. It makes up 20.948 percent of the atmosphere. It also occurs in oceans, lakes, rivers, and ice caps in the form of water. Nearly 89 percent of the weight of water is oxygen. Oxygen is also the most abundant element in the Earth's crust. Its abundance is estimated at about 45 percent in the earth. That makes it almost twice as abundant as the next most common element, silicon.

Oxygen occurs in all kinds of minerals. Some common examples include the oxides, carbonates, nitrates, sulfates, and phosphates. Oxides are chemical compounds that contain oxygen and one other element. Calcium oxide, or lime or quicklime (CaO), is an example. Carbonates are compounds that contain oxygen, carbon, and at least one other element. Sodium carbonate, or soda, soda ash, or sal soda (Na₂CO₃), is an example. It is often found in detergents and cleaning products.

Nitrates, sulfates, and phosphates also contain oxygen and other elements. The other elements in these compounds are nitrogen, sulfur, or phosphorus plus one other element. Examples of these compounds are potassium nitrate, or saltpeter (KNO₃); magnesium sulfate, or Epsom salts (MgSO₄); and calcium phosphate (Ca₃(PO₄)₂).

Isotopes

There are three naturally occurring isotope of oxygen: oxygen-16, oxygen-17, and oxygen-18. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element's name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.

Five radioactive isotopes of oxygen are known also. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive.

None of the radioactive isotopes of oxygen has any commercial use.

Extraction

Oxygen is made from liquid air. Liquid air is made by cooling normal atmospheric air to very low temperatures. As the temperature drops, the gases contained in air turn into liquids. At -182.96°C (-297.33°F), oxygen changes from a gas into a liquid. At -195.79°C (-320.42°F), nitrogen changes from a gas into a liquid. And so on. Eventually, all the gases in air can be made to liquefy (change into a liquid).

But the reverse process also takes place. Suppose liquid air in a container warms up slowly. When its temperature reaches -195.79°C, liquid nitrogen changes back to a gas. A container can be put into place to catch the nitrogen as it boils off the liquid air. When the temperature reaches -182.96°C, oxygen changes from a liquid back to a gas. Another container can be put into place. The escaping oxygen can be collected. Oxygen with a purity of 99.995 percent can be made by this method. It is the only method by which oxygen is made for commercial purposes.

Uses

Many people are familiar with oxygen to help preserve lives. In some cases, people are not able to breathe on their own. Conditions such as emphysema damage the lungs. Oxygen cannot pass through the lungs into the blood stream. One way to treat this condition is to force oxygen into the lungs with a pump.

The same method is used to treat other medical conditions. For example, carbon monoxide poisoning occurs when carbon monoxide gas gets into the blood stream. Auto exhaust, poorly maintained oil furnaces, and wood fires produce carbon monoxide. The carbon monoxide replaces oxygen in the blood. Cells get carbon monoxide instead of oxygen. But they cannot use carbon monoxide, so they begin to die. Forcing oxygen into the blood can reverse some of the damage. In high enough amounts, it can force the carbon monoxide out of the blood and cells can recover.

Oxygen has other interesting uses. For example, it is used in rocket fuels. It is combined with hydrogen in the rocket engines. When hydrogen and oxygen combine, they give off very large amounts of energy. The energy is used to lift the rocket into space.

Metal production accounts for the greatest percentage of oxygen use. For example, oxygen is used to burn off carbon and other impurities that are in iron to make steel. A small amount of these impurities may be desirable in steel, but too much makes it brittle and unusable. The carbon and other impurities are burned off in steel-making by blasting oxygen through molten iron.

Two chemical changes that take place during steel-making are shown below:

The carbon dioxide escapes from the steel-making furnace as a gas. The silicon dioxide (SiO₂) forms slag. Slag is a crusty, metallic material that is scraped off after the steel is produced. Other impurities removed by a blast of oxygen are sulfur, phosphorus, manganese, and other metals.

Oxygen is also used in the production of such metals as copper, lead, and zinc. These metals occur in the earth in the form of sulfides, such as copper sulfide (CuS), lead sulfide (PbS), and zinc sulfide (ZnS). The first step in recovering these metals is to convert them to oxides:

The oxides are then heated with carbon to make the pure metals:

Another use of oxygen is in high-temperature torches. The oxy-acetylene torch, for example, produces heat by burning acetylene gas (C₂H₂) in pure oxygen. The torch can produce temperatures of 3,000°C (5,400°F) and cut through steel and other tough alloys.

Oxygen is also used in the chemical industry as a beginning material in making some very important compounds. Sometimes, the steps to get from oxygen to the final compound are lengthy. As an example, ethylene gas (C₂H₄) can be treated with oxygen to form ethylene oxide (CH₂CH₂O):

About 60 percent of ethylene oxide produced is made into ethylene glycol (CH₂CH₂(OH)₂). Ethylene glycol is used in antifreeze and as a starting point in making polyester fibers, film, plastic containers, bags, and packaging materials.

Compounds

Thousands of oxygen compounds have important commercial uses. Many of these compounds are discussed under other elements.

Coming soon: An oxygen bar near you!

Joseph Priestley was correct! He predicted that "taking a whiff" of oxygen might someday be a luxury for people.

For professional athletes, it already is a common practice. Football players who have run long yardage may inhale some oxygen. It helps them get their energy back.

But now, "taking a whiff" has become even more popular. Oxygen bars are open in Japan, the United States, and other parts of the world. People come to breathe pure oxygen for a few minutes. Of course, they pay a fee for the privilege.

Patrons say that pure oxygen makes them feel better and think more clearly. Others think it improves their looks. Owners of oxygen bars say that city air is often too polluted. They encourage people to "take a whiff" of pure oxygen for their health.

For more information about oxygen bars, contact the National Oxygen Bar Association, 104 North Willow Plaza, Broomfield, CO, 80020; telephone (303) 464-1744; or visit the following website: noba@o2bars.com.

Health effects

Nearly all organisms require oxygen—bacteria, plants, and animals. Humans, for example, can go weeks and even months without food. They can survive for many days without water. But they cannot survive more than a few minutes without oxygen.

Oxygen is used by the cells of animal bodies. It is used to "burn" chemicals and produce energy that cells need to stay alive. Without oxygen, cells begin to die in minutes.

Oxygen family

The oxygen family consists of the elements that make up group 16 on the periodic table: oxygen, sulfur, selenium, tellurium, and polonium. These elements all have six electrons in their outermost energy level, accounting for some common chemical properties among them. In another respect, the elements are quite different from each other. Oxygen is a gaseous nonmetal; sulfur and selenium are solid nonmetals; tellurium is a solid metalloid; and polonium is a solid metal.

Words to Know

Acid: Substances that, when dissolved in water, are capable of reacting with a base to form salts and release hydrogen ions.

Allotrope: One of two or more forms of an element.

Combustion: A form of oxidation that occurs so rapidly that noticeable heat and light are produced.

Cracking: The process by which large hydrocarbon molecules are broken down into smaller components.

Electrolysis: The process by which an electrical current causes a chemical change, usually the breakdown of some substance.

Isotopes: Two or more forms of the same element with the same number of protons but different numbers of neutrons in the atomic nucleus.

Lithosphere: The solid portion of Earth, especially the outer crustal region.

LOX: An abbreviation for liquid oxygen.

Metallurgy: The science and technology that deals with gaining metals from their ores and converting them into forms that have practical value.

Nascent oxygen: An allotrope of oxygen whose molecules each contain a single oxygen atom.

Ozone: An allotrope of oxygen that consists of three atoms per molecule.

Producer gas: A synthetic fuel that consists primarily of carbon monoxide and hydrogen gases.

Proteins: Large molecules that are essential to the structure and functioning of all living cells.

Radioactive decay: The predictable manner in which a population of atoms of a radioactive element spontaneously disintegrate over time.

Oxygen

Oxygen is a colorless, odorless, tasteless gas with a melting point of −218°C (−360°F) and a boiling point of −183°C (−297°F). It is the most abundant element in Earth's crust, making up about one-quarter of the atmosphere by weight, about one-half of the lithosphere (Earth's crust), and about 85 percent of the hydrosphere (the oceans, lakes, and other forms of water). It occurs both as a free element and in a large variety of compounds. In the atmosphere, it exists as elemental oxygen, sometimes known as dioxygen because it consists of diatomic molecules, O₂. In water it occurs as hydrogen oxide, H₂O, and in the lithosphere it occurs in compounds such as oxides, carbonates, sulfates, silicates, phosphates, and nitrates.

Oxygen also exists in two allotropic forms (physically or chemically different forms of the same substance): one atom per molecule (O) and three atoms per molecule (O₃). The former allotrope is known as monatomic, or nascent, oxygen and the latter as triatomic oxygen, or ozone. Under most circumstances in nature, the diatomic form of oxygen predominates. In the upper part of the stratosphere, however, solar energy causes the breakdown of the diatomic form into the monatomic form, which may then recombine with diatomic molecules to form ozone. The presence of ozone in Earth's atmosphere is critical for the survival of life on Earth since that allotrope has a tendency to absorb ultraviolet radiation that would otherwise be harmful or even fatal to both plant and animal life on the planet's surface.

Oxygen was discovered independently by Swedish chemist Carl Scheele (1742–1786) and English chemist Joseph Priestley (1733–1804) in the period between 1773 and 1774. The element was given its name in the late 1770s by French chemist Antoine Laurent Lavoisier (1743–1794). Its name comes from the French word for "acid-former," reflecting Lavoisier's incorrect belief that all acids contain oxygen.

Production. By far the most common method for producing oxygen commercially is by the fractional distillation of liquid air. A sample of air is first cooled to a very low temperature in the range of −200°C (−330°F). At this temperature, most gases that make up air become liquid. The liquid air is then allowed to evaporate. At a temperature of about −196°C (−320°F), nitrogen begins to boil off. When most of the nitrogen is gone, argon and neon also boil off, leaving an impure form of oxygen behind. The oxygen is impure because small amounts of krypton, xenon, and other gases may remain in the liquid form. In order to further purify the oxygen, the process of cooling, liquefying, and evaporation may be repeated.

Oxygen is commonly stored and transported in its liquid form, a form also known as LOX (for l iquid ox ygen). LOX containers look like very large vacuum bottles consisting of a double-walled container with a vacuum between the walls. The element can also be stored and transported less easily in gaseous form in steel-walled containers about 1.2 meters (4 feet) high and 23 centimeters (9 inches) in diameter. In many instances, oxygen is manufactured at the location where it will be used. The process of fractional distillation described earlier is sufficiently simple and inexpensive so that many industries can provide their own oxygen-production facilities.

Uses. Oxygen has so many commercial, industrial, and other uses that it consistently ranks among the top five chemicals in volume of production in the United States. In 1990, for example, about 18 billion kilograms (39 billion pounds) of the element were manufactured in the United States.

The uses to which oxygen is put can be classified into four major categories: metallurgy, rocketry, chemical synthesis, and medicine. In the processing of iron ore in a blast furnace, for example, oxygen is used to convert coke (carbon) to carbon monoxide. The carbon monoxide, in turn, reduces iron oxides to pure iron metal. Oxygen is then used in a second step of iron processing in the Bessemer converter, open hearth, or basic oxygen process method of converting "pig iron" to steel. In this step, the oxygen is used to react with the excess carbon, silicon, and metals remaining in the pig iron that must be removed in order to produce steel.

Another metallurgical application of oxygen is in torches used for welding and cutting. The two most common torches make use of the reaction between oxygen and hydrogen (the oxyhydrogen torch) or between oxygen and acetylene (the oxyacetylene torch). Both kinds of torch produce temperatures in the range of 3,000°C (5,400°F) or more and can, therefore, be used to cut through or weld the great majority of metallic materials.

In the form of LOX, oxygen is used widely as the oxidizing agent in many kinds of rockets and missiles. For example, the huge external fuel tank required to lift the space shuttle into space holds 550,000 liters (145,000 gallons) of liquid oxygen and 1,500,000 liters (390,000 gallons) of liquid hydrogen. When these two elements react in the shuttle's main engines, they provide a maximum thrust of 512,000 pounds.

The chemical industry uses vast amounts of oxygen every year in a variety of chemical synthesis (formation) reactions. One of the most important of these is the cracking of hydrocarbons by oxygen. Under most circumstances, heating a hydrocarbon with oxygen results in combustion, with carbon dioxide and water as the main products. However, if the rate at which oxygen is fed into a hydrocarbon mixture is carefully controlled, the hydrocarbon is "cracked," or broken apart to produce other products, such as acetylene, ethylene, and propylene.

Various types of synthetic fuels can also be manufactured with oxygen as one of the main reactants. Producer gas, as an example, is manufactured by passing oxygen at a controlled rate through a bed of hot coal or coke. The majority of carbon dioxide produced in this reaction is reduced to carbon monoxide so that the final product (the producer gas) consists primarily of carbon monoxide and hydrogen.

Perhaps the best-known medical application of oxygen is in oxygen therapy, where patients who are having trouble breathing are given doses of pure or nearly pure oxygen. Oxygen therapy is often used during surgical procedures, during childbirth, during recovery from heart attacks, and during treatment for infectious diseases. In each case, providing a person with pure oxygen reduces the stress on his or her heart and lungs, speeding the rate of recovery.

Pure oxygen or air enriched with oxygen may also be provided in environments where breathing may be difficult. Aircraft that fly at high altitudes, of course, are always provided with supplies of oxygen in case of any problems with the ship's normal air supply. Deep-sea divers also carry with them or have pumped to them supplies of air that are enriched with oxygen.

Some water purification and sewage treatment plants use oxygen. The gas is pumped through water to increase the rate at which naturally occurring bacteria break down organic waste materials. A similar process has been found to reduce the rate at which eutrophication takes place in lakes and ponds and, in some cases, to actually reverse that process. (Eutrophication is the dissolving of nutrients in a body of water. Growth in aquatic plant life and a decrease in dissolved oxygen are the two main results of the process.)

Finally, oxygen is essential to all animal life on Earth. A person can survive a few days or weeks without water or food but no more than a few minutes without oxygen. In the absence of oxygen, energy-generating chemical reactions taking place within cells would come to an end, and a person would die.

Sulfur

Sulfur is a nonmetallic element that can exist in many allotropic forms (physically or chemically different forms of the same substance). The most familiar are called rhombic and monoclinic sulfur. Both are bright yellow solids with melting points of about 115°C (239°F). A third form is called plastic or amorphous sulfur. It is a brownish liquid produced when rhombic or monoclinic sulfur is melted.

Sulfur itself has no odor at all. It has a bad reputation in this regard, however, because some of its most common compounds have strong smells. Sulfur dioxide, one of these compounds, has a sharp, choking, suffocating effect on anyone who breathes it. The "fire and brimstone" of the Bible was one of the worst punishments that its authors could imagine. The brimstone in this expression referred to burning sulfur, or sulfur dioxide. The fact that sulfur comes from deep under the ground and that sulfur dioxide can be smelled in the fumes of volcanoes further fueled people's imaginations of what Hell must be like.

A second sulfur compound with a bad odor is hydrogen sulfide. The strong smell of rotten eggs is due to the presence of this compound.

Occurrence and preparation. Sulfur is the sixteenth most abundant element in Earth's crust. It occurs both as an element and in a variety of compounds. As an element it can be found in very large, underground mines, most commonly along the Gulf Coast of the United States and in Poland and Sicily. The sulfur is extracted from these mines by means of the Frasch process. In this process, superheated steam is pumped through the outermost of a set of three pipes. Compressed air is forced down the innermost pipe. The superheated steam causes the underground sulfur to melt, and the compressed air forces it upward, through the middle of the three pipes, to Earth's surface.

Sulfur is also widely distributed in the form of minerals and ores. Many of these are in the form of sulfates, including gypsum (calcium sulfate, CaSO₄), barite (barium sulfate, BaSO₄), and Epsom salts (magnesium sulfate, MgSO₄). Others are metal sulfides, including iron pyrites (iron sulfide, FeS₂), galena (lead sulfide, PbS), cinnabar (mercuric sulfide, HgS), stibnite (antimony sulfide, Sb₂S₃), and zinc blende (zinc sulfide, ZnS). The sulfur is recovered from these metal ores by heating them strongly in air, which converts the sulfur to sulfur dioxide and releases the pure metal. Then the sulfur dioxide can go directly into the manufacture of sulfuric acid, which is where more than 90 percent of the world's mined sulfur winds up.

Uses of sulfur and its compounds. Some sulfur is used directly as a fungicide and insecticide, in matches, fireworks, and gunpowder, and in the vulcanization of natural rubber (a treatment that gives rubber elasticity and strength). Most, however, is converted into a multitude of useful compounds.

Sulfuric acid is by far the most important of all sulfur compounds. Nearly 90 percent of all sulfur produced is converted first into sulfur dioxide and then into sulfuric acid. The acid consistently ranks number one among the chemicals produced in the United States. In 1990, more than 40 billion kilograms (89 billion pounds) of sulfuric acid were manufactured, more than 50 percent as much as the second most popular chemical (nitrogen gas). Sulfuric acid is used in the production of fertilizers, automobile batteries, petroleum products, pigments, iron and steel, and many other products.

The sulfur cycle. Like nitrogen, carbon, and phosphorus, sulfur passes through the gaseous, liquid, and solid parts of our planet in a series of continuous reactions known as the sulfur cycle. The main steps in the sulfur cycle are illustrated in the accompanying figure.

Sulfur is produced naturally as a result of volcanic eruptions and through emissions from hot springs. It enters the atmosphere primarily in the form of sulfur dioxide, then remains in the atmosphere in that form or, after reacting with water, in the form of sulfuric acid. Sulfur is carried back to Earth's surface as acid deposition when it rains or snows.

On Earth's surface, sulfur dioxide and sulfuric acid react with metals to form sulfates and sulfides. The element is also incorporated by plants in a form known as organic sulfur. Certain amino acids, the compounds from which proteins are made, contain sulfur. Organic sulfur from plants is eventually passed on to animals that eat those plants. It is, in turn, converted from plant proteins to animal proteins.

When plants and animals die, sulfur is returned to the soil where it is converted by microorganisms into hydrogen sulfide. Hydrogen sulfide gas is then returned to the atmosphere, where it is oxidized to sulfuric acid.

Human activities influence the sulfur cycle in a number of ways. For example, when coal and metallic ores are mined, sulfur and sulfides may be released and returned to the soil. Also, the combustion of coal, oil, and natural gas often releases sulfur dioxide to the atmosphere. This sulfur dioxide is added to the amount already present from natural sources, greatly increasing the amount of acid precipitation that falls to Earth's surface. Some people believe that acid precipitation (or acid rain) is responsible for the death of trees and other plants, the acidification of lakes that has hurt marine animals, damage to metal and stone structures, and other environmental harm.

Selenium, tellurium, and polonium

Selenium and tellurium are both relatively rare elements. They rank in the bottom ten percent of all elements in terms of abundance. They tend to occur in Earth's crust in association with ores of copper and other metals. Both are obtained as a by-product of the electrolytic refining of copper. During that process, they sink to the bottom of the electrolysis tank, where they can be removed from the sludge that develops.

Selenium occurs in a variety of allotropic forms (physically or chemically different forms of the same substance), the most common of which is a red powder that becomes black when exposed to air. The element's melting point is 217°C (423°F), and its boiling point is 685°C (1,265°F). Tellurium is a silvery-white solid that looks like a metal (although it is actually a metalloid). Its melting point is 450°C (842°F), and its boiling point is 990°C (1,814°F).

Selenium has an interesting role in living organisms. It is essential in very low concentrations for maintaining health in most animals. In fact, it is often added to animal feeds. In higher concentrations, however, the element has been found to have harmful effects on animals, causing deformed young and diseased adults.

The primary uses of selenium are in electronics and in the manufacture of colored glass. Photocopying machinery, solar cells, photocells, television picture tubes, and electronic rectifiers and relays (used to control the flow of electric current) all use selenium. Some of the most beautiful colored glasses, ranging from pale pink to brilliant reds, are made with compounds of selenium.

Small amounts of tellurium are also used in the production of colored glass. More than 90 percent of the element, however, goes to the production of alloys of iron and other metals.

Polonium has 27 isotopes, all of which are radioactive. It occurs naturally in uranium ores, where it is the final product in the long series of reactions by which uranium undergoes radioactive decay. It is one of the rarest elements on Earth, with an abundance of no more than about 3 × 10⁻¹⁰ parts per million. The discovery of polonium in 1898 by Polish-French chemist Marie Curie (1867–1934) is one of the most dramatic stories in the history of science. She processed tons of uranium ore in order to obtain a few milligrams of the new element, which she then named after her homeland of Poland. Polonium finds limited use in highly specialized power-generating devices, such as those used for space satellites and space probes.

Oxygen

Background

Oxygen is one of the basic chemical elements. In its most common form, oxygen is a colorless gas found in air. It is one of the life-sustaining elements on Earth and is needed by all animals. Oxygen is also used in many industrial, commercial, medical, and scientific applications. It is used in blast furnaces to make steel, and is an important component in the production of many synthetic chemicals, including ammonia, alcohols, and various plastics. Oxygen and acetylene are combusted together to provide the very high temperatures needed for welding and metal cutting. When oxygen is cooled below -297° F (-183° C), it becomes a pale blue liquid that is used as a rocket fuel.

Oxygen is one of the most abundant chemical elements on Earth. About one-half of the earth's crust is made up of chemical compounds containing oxygen, and a fifth of our atmosphere is oxygen gas. The human body is about two-thirds oxygen. Although oxygen has been present since the beginning of scientific investigation, it wasn't discovered and recognized as a separate element until 1774 when Joseph Priestley of England isolated it by heating mercuric oxide in an inverted test tube with the focused rays of the sun. Priestley described his discovery to the French scientist Antoine Lavoisier, who experimented further and determined that it was one of the two main components of air. Lavoisier named the new gas oxygen using the Greek words oxys, meaning sour or acid, and genes, meaning producing or forming, because he believed it was an essential part of all acids.

In 1895, Karl Paul Gottfried von Linde of Germany and William Hampson of England independently developed a process for lowering the temperature of air until it liquefied. By carefully distillation of the liquid air, the various component gases could be boiled off one at a time and captured. This process quickly became the principal source of high quality oxygen, nitrogen, and argon.

In 1901, compressed oxygen gas was burned with acetylene gas in the first demonstration of oxy-acetylene welding. This technique became a common industrial method of welding and cutting metals.

The first use of liquid rocket propellants came in 1923 when Robert Goddard of the United States developed a rocket engine using gasoline as the fuel and liquid oxygen as the oxidizer. In 1926, he successfully flew a small liquid-fueled rocket a distance of 184 ft (56 m) at a speed of about 60 mph (97 kph).

After World War II, new technologies brought significant improvements to the air separation process used to produce oxygen. Production volumes and purity levels increased while costs decreased. In 1991, over 470 billion cubic feet (13.4 billion cubic meters) of oxygen were produced in the United States, making it the second-largest-volume industrial gas in use.

Worldwide the five largest oxygen-producing areas are Western Europe, Russia (formerly the USSR), the United States, Eastern Europe, and Japan.

Raw Mcatericals

Oxygen can be produced from a number of materials, using several different methods. The most common natural method is photo-synthesis, in which plants use sunlight convert carbon dioxide in the air into oxygen. This offsets the respiration process, in which animals convert oxygen in the air back into carbon dioxide.

The most common commercial method for producing oxygen is the separation of air using either a cryogenic distillation process or a vacuum swing adsorption process. Nitrogen and argon are also produced by separating them from air.

Oxygen can also be produced as the result of a chemical reaction in which oxygen is freed from a chemical compound and becomes a gas. This method is used to generate limited quantities of oxygen for life support on submarines, aircraft, and spacecraft.

Hydrogen and oxygen can be generated by passing an electric current through water and collecting the two gases as they bubble off. Hydrogen forms at the negative terminal and oxygen at the positive terminal. This method is called electrolysis and produces very pure hydrogen and oxygen. It uses a large amount of electrical energy, however, and is not economical for large-volume production.

The Manufacturing
Process

Most commercial oxygen is produced using a variation of the cryogenic distillation process originally developed in 1895. This process produces oxygen that is 99+% pure. More recently, the more energy-efficient vacuum swing adsorption process has been used for a limited number of applications that do not require oxygen with more than 90-93% purity.

Here are the steps used to produce commercial-grade oxygen from air using the cryogenic distillation process.

Pretreating

Because this process utilizes an extremely cold cryogenic section to separate the air, all impurities that might solidify—such as water vapor, carbon dioxide, and certain heavy hydrocarbons—must first be removed to prevent them from freezing and plugging the cryogenic piping.

This test tube is one of the most popular artifacts in Henry Ford Museum & Greenfield Village in Dearborn, Michigan. It is said to contain the last breath of Thomas Alva Edison, the great inventor. According to Edison's son Charles, a set of eight empty test tubes sat on the table next to Edison's deathbed in 1931. Immediately after Edison expired, his physician, put several of the tubes up to Edison's lips to catch the carbon dioxide from his deflating lungs. Then, the physician carefully sealed each tube with paraffin and gave the tubes to Charles Edison. Charles Edison knew that Henry Ford's idol was Thomas Edison and presented Ford with one of the tubes as a keepsake. The museum acquired the tube after the death of both Henry and Clara Ford.

There is some discussion among visitors just how much carbon dioxide and how much oxygen currently is contained in the tube. Some ask if anyone evacuated the tube of oxygen before putting the tube to Edison's mouth (very unlikely). If not, how much of Edison's breath could be in the tube? So, they say, it contains both carbon dioxide and oxygen? Nonetheless, it is an unconventional tribute to a great man by those sorry to see his light extinguished.

Nancy EV Bryk

• 1 The air is compressed to about 94 psi (650 kPa or 6.5 atm) in a multi-stage compressor. It then passes through a water-cooled aftercooler to condense any water vapor, and the condensed water is removed in a water separator.
• 2 The air passes through a molecular sieve adsorber. The adsorber contains zeolite and silica gel-type adsorbents, which trap the carbon dioxide, heavier hydrocarbons, and any remaining traces of water vapor. Periodically the adsorber is flushed clean to remove the trapped impurities. This usually requires two adsorbers operating in parallel, so that one can continue to process the air-flow while the other one is flushed.

Separating

Air is separated into its major components—nitrogen, oxygen, and argon—through a distillation process known as fractional distillation. Sometimes this name is shortened to fractionation, and the vertical structures used to perform this separation are called fractionating columns. In the fractional distillation process, the components are gradually separated in several stages. At each stage the level of concentration, or fraction, of each component is increased until the separation is complete.

Because all distillation processes work on the principle of boiling a liquid to separate one or more of the components, a cryogenic section is required to provide the very low temperatures needed to liquefy the gas components.

• 3 The pretreated air stream is split. A small portion of the air is diverted through a compressor, where its pressure is boosted. It is then cooled and allowed to expand to nearly atmospheric pressure. This expansion rapidly cools the air, which is injected into the cryogenic section to provide the required cold temperatures for operation.
• 4 The main stream of air passes through one side of a pair of plate fin heat exchangers operating in series, while very cold oxygen and nitrogen from the cryogenic section pass through the other side. The incoming air stream is cooled, while the oxygen and nitrogen are warmed. In some operations, the air may be cooled by passing it through an expansion valve instead of the second heat exchanger. In either case, the temperature of the air is lowered to the point where the oxygen, which has the highest boiling point, starts to liquefy.
• 5 The air stream—now part liquid and part gas—enters the base of the high-pressure fractionating column. As the air works its way up the column, it loses additional heat. The oxygen continues to liquefy, forming an oxygen-rich mixture in the bottom of the column, while most of the nitrogen and argon flow to the top as a vapor.
• 6 The liquid oxygen mixture, called crude liquid oxygen, is drawn out of the bottom of the lower fractionating column and is cooled further in the subcooler. Part of this stream is allowed to expand to nearly atmospheric pressure and is fed into the low-pressure fractionating column. As the crude liquid oxygen works its way down the column, most of the remaining nitrogen and argon separate, leaving 99.5% pure oxygen at the bottom of the column.
• 7 Meanwhile, the nitrogen/argon vapor from the top of the high-pressure column is cooled further in the subcooler. The mixed vapor is allowed to expand to nearly atmospheric pressure and is fed into the top of the low-pressure fractionating column. The nitrogen, which has the lowest boiling point, turns to gas first and flows out the top of the column as 99.995% pure nitrogen.
• 8 The argon, which has a boiling point between the oxygen and the nitrogen, remains a vapor and begins to sink as the nitrogen boils off. As the argon vapor reaches a point about two-thirds the way down the column, the argon concentration reaches its maximum of about 7-12% and is drawn off into a third fractionating column, where it is further recirculated and refined. The final product is a stream of crude argon containing 93-96% argon, 2-5% oxygen, and the balance nitrogen with traces of other gases.

Purifying

The oxygen at the bottom of the low-pressure column is about 99.5% pure. Newer cryogenic distillation units are designed to recover more of the argon from the low-pressure column, and this improves the oxygen purity to about 99.8%.

• 9 If higher purity is needed, one or more additional fractionating columns may be added in conjunction with the low-pressure column to further refine the oxygen product. In some cases, the oxygen may also be passed over a catalyst to oxidize any hydrocarbons. This process produces carbon dioxide and water vapor, which are then captured and removed.

Distributing

About 80-90% of the oxygen produced in the United States is distributed to the end users in gas pipelines from nearby air separation plants. In some parts of the country, an extensive network of pipelines serves many end users over an area of hundred of miles (kilometers). The gas is compressed to about 500 psi (3.4 MPa or 34 atm) and flows through pipes that are 4-12 in (10-30 cm) in diameter. Most of the remaining oxygen is distributed in insulated tank trailers or railroad tank cars as liquid oxygen.

• 10 If the oxygen is to be liquefied, this process is usually done within the low-pressure fractionating column of the air separation plant. Nitrogen from the top of the low-pressure column is compressed, cooled, and expanded to liquefy the nitrogen. This liquid nitrogen stream is then fed back into the low-pressure column to provide the additional cooling required to liquefy the oxygen as it sinks to the bottom of the column.
• 11 Because liquid oxygen has a high boiling point, it boils off rapidly and is rarely shipped farther than 500 mi (800 km). It is transported in large, insulated tanks. The tank body is constructed of two shells and the air is evacuated between the inner and outer shell to retard heat loss. The vacuum space is filled with a semisolid insulating material to further halt heat flow from the outside.

Quality Control

The Compressed Gas Association establishes grading standards for both gaseous oxygen and liquid oxygen based on the amount and type of impurities present. Gas grades are called Type I and range from A, which is 99.0% pure, to F, which is 99.995% pure. Liquid grades are called Type II and also range from A to F, although the types and amounts of allowable impurities in liquid grades are different than in gas grades. Type I Grade B and Grade C and Type II Grade C are 99.5% pure and are the most commonly produced grades of oxygen. They are used in steel making and in the manufacture of synthetic chemicals.

The operation of cryogenic distillation airseparation units is monitored by automatic instruments and often uses computer controls. As a result, their output is consistent in quality. Periodic sampling and analysis of the final product ensures that the standards of purity are being met.

The Future

In January 1998, the United States launched the Lunar Prospector satellite into orbit around the moon. Among its many tasks, this satellite will be scanning the surface of the moon for indications of water. Scientists hope that if sufficient quantities of water are found, it could be used to produce hydrogen and oxygen gases through electrolysis, using solar power to generate the electricity. The hydrogen could be used as a fuel, and the oxygen could be used to provide life support for lunar colonies. Another plan involves extracting oxygen from chemical compounds in the lunar soil using a solar-powered furnace for heat.

Where to Learn More

Books

Brady, George S., Henry R. Clauser, and John A. Vaccari. Materials Handbook, 14th Edition. McGraw-Hill, 1997.

Handbook of Compressed Gases, 3rd edition. Compressed Gas Association, Inc., Van Nostrand Reinhold Co., Inc., 1990.

Heiserman, David L. Exploring Chemical Elements and Their Compounds. TAB Books, 1992.

Kent, James A., editor. Riegel's Handbook of Industrial Chemistry, 9th edition. International Thomson Publishing, 1997.

Kroschwitz, Jacqueline I., executive editor, and Mary Howe-Grant, editor. Encyclopedia of Chemical Technology, 4th edition. John Wiley and Sons, Inc., 1993.

Stwertka, Albert. A Guide to the Elements. Oxford University Press, 1996.

Periodicals

Allen, J.B. "Making Oxygen on the Moon," Popular Science (August 1995): 23.

Other

Air Products and Chemicals, Inc. http://www.airproducts.com/gases/oxgen.html.

http://www.intercorr.com/periodic/8.htm (This website contains a summary of the history, sources, properties, and uses of oxygen.)

—Chris Cavette

Oxygen

Oxygen is essential to all aquatic animal life. Without it respiration is impossible. Because the distribution of oxygen throughout a body of water can vary widely, it is important to find the specific areas which hold optimal amounts to support fish life. Oxygen concentration is a prime factor in locating where specific fish species are found.

It’s self-evident that air includes an abundance of oxygen; however, water holds a much smaller amount. For example, one liter of air contains about 210 cubic centimeters of oxygen while one liter of water has only about 9 cubic centimeters. In air the lack of oxygen is rarely a problem, but in water its effect is more apparent. Water harbors a tenuous margin of safety because its oxygen content is small and varies widely. Insufficient oxygen supplies will result in animal death.

Two physical factors affecting water’s oxygen content are altitude (atmospheric pressure) and temperature. Oxygen solubility in water increases with decreases in both temperature and altitude. Conversely, oxygen solubility decreases with rises in temperature and altitude.

Atmospheric pressure is inversely proportional to altitude. That is, a rise in altitude lowers the atmospheric pressure, and a decrease in altitude raises the atmospheric pressure. On the other hand, atmospheric pressure is directly proportional to oxygen water solubility. That is, a rise in atmospheric pressure increases the water’s oxygen content while a decrease in atmospheric pressure lowers the water’s oxygen content.

Oxygen diffusion occurs at the water’s surface, and its mixing throughout happens during a lake’s spring and fall turnover times. The oxygen diffusion process is very slow. Both wind and wave actions mix oxygen at the surface, while plant photosynthesis generates oxygen when plants are present. In deep water, where light cannot penetrate, oxygen cannot be produced by plants. Hence, a lake’s hypolimnion only receives its limited oxygen by spring and fall mixing turnovers. This is why a lake’s greatest depths can become so oxygen poor. The stratification of water layers seals the hypolimnion from the surface air supply and this seal is only interrupted at spring and fall.

An increase in water temperature lowers the amount of oxygen it can contain. During hot conditions a lake’s shallows may become too warm to hold sufficient oxygen, and the animal life forms must migrate to cooler regions which have more oxygen. Inlet streams and underwater springs can supply the needed oxygenated

water. The deeper shaded areas just above the thermocline can be cool enough to hold sufficient oxygen. Fish concentrate in these limited areas when such conditions prevail.

Since a lake depends upon wave action, photosynthesis, and currents to mix and produce its oxygen, areas of adequate and inadequate oxygenated water exist simultaneously. This results in areas where fish can and cannot live. It is a precarious balancing act which includes oxygen content, temperature, light intensity, protective cover, and available food supply that dictates just where fish can be found.

The altitude’s influence on an alpine lake can be dramatic. For example, in my youth I spent considerable time fishing alpine lakes above ten thousand feet in altitude. During the dog days of summer my success dramatically fell. I was puzzled because these alpine lakes contained colder water than their valley counterparts. The alpine lakes’ depths were void of fish, and my success was limited to the lakes’ shallows. Earlier in the season the angling was good in the alpine lakes’ depths. Now in midsummer the fishing was poor. The valley lakes fished best in their depths while the alpine lakes fished best in their shallows. I neglected to take into account the difference in altitude and atmospheric pressure between the alpine and valley lakes. The high altitude lakes’ oxygen was less because there was less oxygen available when the spring and fall turnovers occurred. This limited oxygen supply was quickly depleted in the alpine lakes’ depths, making it so fish couldn’t live there. The lower elevation lakes took in higher oxygen content during the spring and fall turnover times. This confined the alpine lakes’ fish to the shallows while the valley lakes’ fish were in the lakes’ depths. I once thought that all big fish lived in the depths and only small ones were in the shallows. This notion is untrue.

The current in a river mixes oxygen much better than limited currents and waves in a lake. Also, this thorough mixing causes little variation in both, oxygen and temperature differences. During high temperature times, the whitewater river sections mix higher oxygen content. In times of hot spells fish may migrate to the rapids for survival.

Look for them in the pocket areas downstream from rapids. Also areas close to adjacent inlet streams and underwater springs can harbor more favorable conditions. For example, in warm thermal rivers such as Yellowstone Park’s Firehole River, fish migrate to the mouths of cooler tributary streams during warm seasons.

In conclusion, the oxygen content of water is a prime factor in determining the location of fish. An adequate oxygen supply is essential to sustain fish life.

oxygen is the most common of all chemical elements on earth, being found in water, air, and most mineral and organic substances, including most compounds in the human body. It combines with almost all other elements, and is so reactive that it was given the Greek name ‘oxygen’, meaning acid-forming. However, most of the compounds it forms are not acids. Its chemical reactions usually form heat (as in the animal body) and sometimes light (as in candles).

It has always been known that animals cannot live without air, but in 1674 Mayow showed that only one part of the air, about one-fifth, is essential for life, and named it ‘vital air’. A hundred years later Priestley isolated this part, oxygen; Lavoisier purified oxygen and its properties began to be studied.

Atmospheric air contains 21% oxygen, at a pressure of about 150 mm Hg varying with barometric pressure and to a small extent with humidity. It enters the lungs during breathing and is absorbed into the blood passively by diffusion, combining with haemoglobin and being carried in the bloodstream to all parts of the body. There it is used to metabolize or ‘burn’ foodstuffs in the cells, especially fats and carbohydrates, providing heat and creating new chemical compounds, water, and the waste product carbon dioxide. Tissues and organs vary in the length of time they can survive without oxygen, according to their provision for anaerobic metabolism. The brain cannot survive without oxygen; the cessation of breathing will cause unconsciousness in a few minutes, and death soon afterwards. Other tissues such as skeletal muscle can continue to work for a limited time, when glycogen stores are broken down without oxygen to provide energy; lactic acid is a by-product that leaks into the blood and makes it more acid, but can be recycled into carbohydrate stores in the liver.

In quiet breathing at rest we absorb about 0.2–0.3 litres/min of oxygen (depending on body size), but in vigorous exercise this can go up to over 2 litres/min. This increase is accomplished by increased breathing (which supplies oxygen to the lungs at a greater rate), increased cardiac output and flow of blood to the muscles, and greater extraction of oxygen from the blood by the muscles. If the oxygen supply to the muscles is inadequate then the anaerobic threshold is passed and anaerobic metabolism takes place, with production of lactic acid. After the exercise additional oxygen is needed to convert the lactic acid back to glycogen, and breathing remains enhanced while the oxygen debt is repaid.

The supply of oxygen to the body depends not on the percentage in the air breathed but on its tension or pressure. At high altitude, say 5000 metres above sea level, the percentage of oxygen is still 21%, but because atmospheric pressure is halved, the oxygen pressure is half that at sea level — 75 mm Hg rather than 150 mm Hg. A person may as a result suffer from hypoxia — a lack of oxygen.

High oxygen pressures can be harmful and cause oxygen poisoning, including lung damage and brain dysfunction. In nature high oxygen pressures only exist in deep water diving, and mankind has not had to evolve ways of combating them. Once scientists had purified oxygen it became possible to administer it to patients; this has life-saving possibilities, but care has to be taken not to exceed the toxic level.

Some compounds rich in oxygen, such as the pollutant ozone (itself a molecular form of oxygen), and hydrogen peroxide, can react with cells to produce strongly reactive forms of oxygen. Superoxide anions and unstable oxygen free radicals (such as hydroxyl and hydroperoxy radicals) can be toxic to cells, by way of excess lipid peroxidation. These are implicated, for example, in damage following the restoration of blood flow (reperfusion) after the blockage which causes heart attacks or strokes, and in a variety of other disease processes. However the body does have inherent enzymatic defences against free radical accumulation, and there are antioxidants, such as uric acid, ascorbate, and glutathione, which can inactivate them. Free radicals are likely to contribute also to the ageing process: the very substance by which we live may itself limit our lifespan. Thus oxygen, like most good things, can be dangerous in excess.

Mankind evolved to live close to sea level. Climbing mountains (causing hypoxia) and deep-sea diving (causing nitrogen narcosis or oxygen poisoning) can both be dangerous, in the absence of the right precautions.

John Widdicombe

oxygen gaseous chemical element; symbol O; at. no. 8; interval in which at. wt. ranges 15.99903–15.99977; m.p. -218.4°C; b.p. -182.962°C; density 1.429 grams per liter at STP; valence -2. The existence and properties of oxygen had been noted by many scientists before the announcement of its isolation by Priestley in 1774. Scheele had also succeeded in preparing oxygen from a number of substances, but publication of his findings was delayed until after that of Priestley's. As a result, Priestley and Scheele are credited with the discovery of the element independently. The fact that the gas is a component of the atmosphere was finally and definitely established by Lavoisier a few years later. In 1929, W. F. Giaque and H. L. Johnston announced the discovery of two isotopes of oxygen, of mass numbers 17 and 18.

Properties and Compounds

Oxygen is a colorless, odorless, tasteless gas; it is the first member of Group 16 of the periodic table . It is denser than air and only slightly soluble in water. A poor conductor of heat and electricity, oxygen supports combustion but does not burn. Normal atmospheric oxygen is a diatomic gas (O ₂ ) with molecular weight 31.9988. Ozone is a highly reactive triatomic (O ₃ ) allotrope of oxygen (see allotropy ). When cooled below its boiling point oxygen becomes a pale blue liquid; when cooled still further the liquid solidifies, retaining its color. Oxygen is paramagnetic in its solid, liquid, and gaseous forms. Although eight isotopes of oxygen are known, atmospheric oxygen is a mixture of the three isotopes with mass numbers 16, 17, and 18.

Oxygen is extremely active chemically, forming compounds with almost all of the elements except the inert gases. Oxygen unites directly with a number of other elements to form oxides . It is a constituent of many acids and of hydroxides, carbohydrates, proteins, fats and oils, alcohols, cellulose, and numerous other compounds such as the carbonates, chlorates, nitrates and nitrites, phosphates and phosphites, and sulphates and sulphites.

The common reaction in which it unites with another substance is called oxidation (see oxidation and reduction ). The burning of substances in air is rapid oxidation or combustion . The respiration of animals and plants is a form of oxidation essential to the liberation of the energy stored in such food materials as carbohydrates and fats. The rusting of iron and the corrosion of many metals results from the action of the oxygen in the air.

Natural Occurrence and Preparation

Oxygen is the most abundant element on earth, constituting about half of the total material of its surface. Most of this oxygen is combined in the form of silicates, oxides and water. It makes up about 90% of water, two thirds of the human body and one fifth by volume of air. It is found in the sun, and has a role in the stellar carbon cycle (see nucleosynthesis ). Oxygen is prepared for commercial use by the liquefaction and fractional distillation of air and more expensively by the electrolysis of water; it is stored and transported under high pressure in steel cylinders. It can also be obtained by heating certain of its compounds, such as barium peroxide, potassium chlorate, and the red oxide of mercury.

Uses

Oxygen is of great importance in the chemical and the iron and steel industries. Its major use is in steel production, for example in the Bessemer process . The oxyacetylene torch is another important industrial application. Oxygen is utilized in medicine in the treatment of respiratory diseases and is mixed with other gases for respiration in submarines, high-flying aircraft, and spacecraft. Liquid oxygen is used as an oxidizer in the fuel systems of large rockets.

Oxygen was formerly the official standard for the atomic weights of elements. The chemists used natural oxygen, a mixture of three isotopes, to which the value of 16 was assigned while the physicists assigned the value of 16 specifically to the oxygen isotope 16. In 1961 carbon-12 replaced oxygen as the standard.

Oxygen

Oxygen is the simplest group VIA element and is, under normal atmospheric conditions, usually found as a colorless, odorless, and tasteless gas. Oxygen has an atomic number of 8 and an atomic mass of 16.0 amu. The liquid and solid forms, which are strongly paramagnetic, are a pale blue color. Oxygen has a boiling point of −297°F (−182.8°C) and a melting point of −368.7°F (−222.6°C).

Oxygen is the third most abundant element found in the Sun , after hydrogen and helium, and plays an important role in the carbon-nitrogen cycle. Oxygen composes 21% of Earth's atmosphere by volume and is vital to the existence of carbon-based life forms.

Although English chemist Joseph Priestley (1733–1804) is generally credited with the discovery of oxygen in 1774, many science historians contend that Swedish chemist Carl Scheele (1742–1786) probably discovered oxygen a few years prior to Priestly. French chemist Antoine Lavoisier's (1743–1794) contributions to the study of the important reactions, combustion and oxidation, were spurred by the discovery of oxygen. Lavoisier noticed that something was absorbed when combustion took place and that it was obtained from the surrounding air. Lavoisier noted that the increase in the weight of the substance burned was equal to the decrease in the weight of the air used. His studies lead to Lavoisier's oxidation theory, which eventually superseded the phlogistonists' theory (i.e., that every combustible substance was thought to contain a phlogiston, or inherent principal of fire, liberated through burning, along with a residue) that was widely accepted at that time. Lavoisier eventually named the gas he studied oxygen from the Greek oxys meaning acid or sharp, and geinomial meaning forming. Lavoisier named the gas oxygen because he noted that the burned materials were converted into acids.

Although oxygen has nine isotopes, natural oxygen is a mixture of only three of these. The most abundant isotope, oxygen-18, is stable and available commercially. The most common use for commercial oxygen gas is in enrichment of steel blast furnaces and for medical purposes. Large quantities are also used in making synthetic ammonia gas, methanol and ethylene oxide. Oxygen is also consumed in oxy-acetylene welding. Most commercial oxygen is produced in air separation plants. It is estimated that the United States consumes 20 million tons of oxygen in commercial use per year and the demand is expected to increase dramatically.

When oxygen is exposed to ultraviolet light, as from the Sun, or an electrical discharge, as from lightening, ozone (O₃) is formed. Although ozone is toxic to breathe, the 0.12 in (3 mm) thick layer of ozone in the earth's atmosphere provides a shield from harmful ultraviolet rays from the Sun. The ozone layer has recently been the subject of intense scientific interest to determine whether, and to what extent, it may be deteriorating, mainly from pollutants in the atmosphere. Unlike pure oxygen gas, ozone has a bluish color and its liquid and solid forms are bluish black to violet-black.

See also Atmospheric chemistry; Atmospheric composition and structure; Global warming; Greenhouse gases and greenhouse effect; Ozone layer and hole dynamics; Ozone layer depletion

Oxygen

melting point: −218.4°C
boiling point: −182.96°C
density: 1.429 g/L
most common ions: OH⁻, OH₂⁻, O²⁻

Joseph Priestley and Carl Scheele (each working independently) are credited with the isolation and "discovery" in 1774 of the element oxygen. A few years later Antoine Lavoisier showed that oxygen is a component of the atmosphere. Oxygen is the most abundant element on Earth, constituting about half of the total material of its surface (47 percent by weight of the lithosphere and 89 percent by weight of the ocean) and about 21 percent by volume of the air. Under ordinary conditions (STP) on Earth, oxygen is a colorless, odorless, tasteless gas that is only slightly soluble in water. Oxygen has a pale blue color in the liquid and the solid phases. Ordinary oxygen gas (O₂) exists as diatomic molecules. It also exists in another allotropic form, the triatomic molecule ozone (O₃). Although eight isotopes of oxygen are known, atmospheric oxygen is a mixture of only three: those having mass numbers 16, 17, and 18.

Oxygen is very reactive. Its reaction with another substance to form an oxide is called oxidation . It is a constituent of a number of compound groups, such as acids, hydroxides, carbonates, chlorates, nitrates and nitrites, and phosphates and phosphites—as well as carbohydrates, proteins, fats, and oils. The respiration of animals and plants is actually a form of oxidation, essential to the production of energy within these organisms. The burning of substances in air is a rapid form of oxidation called combustion . In the eighteenth century the idea of combustion replaced the idea (phlogiston theory) that a colorless, odorless, tasteless, and weightless substance named phlogiston was given off during the burning of a substance.

see also Lavoisier, Antoine; Priestley, Joseph; Scheele, Carl.

Ágúst Kvaran

Bibliography

Lane, Nick (2002). Oxygen: The Molecule that Made the World. Oxford: Oxford University Press.

oxygen (symbol O) Common gaseous element that is necessary for the respiration of plants and animals and for combustion. Oxygen is the most abundant element in the Earth's crust (49.2% by weight), and is a constituent of water and many rocks. It is also present in the atmosphere (23.14% by weight). Oxygen was discovered in 1772 by English chemist Joseph Priestley and independently that year by Swedish chemist Karl Scheele. It can be obtained by the electrolysis of water or fractional distillation of liquid air. It is used in apparatus for breathing (oxygen masks) and resuscitation (oxygen tents); liquid oxygen is used in rocket fuels. It is chemically reactive, and forms compounds with most elements (especially by oxidation). Properties: at.no. 8; r.a.m. 15.9994; r.d. 1.429; m.p. −218.4°C (−361.1°F); b.p. −182.96 °C; (−297.3°F); most common isotope O¹⁶ (99.759%). See also oxidation-reduction; ozone

oxygen (oks-i-jĕn) n. an odourless colourless gas that makes up one-fifth of the atmosphere. Oxygen is essential to most forms of life; in humans, it is absorbed into the blood from air breathed into the lungs. Oxygen is administered therapeutically in various conditions in which the tissues are unable to obtain an adequate supply through the lungs. Symbol: O. o. deficit a physiological condition that exists in cells during periods of temporary oxygen shortage. o. tent see tent.

ox·y·gen / ˈäksəjən/ • n. a colorless, odorless reactive gas, the chemical element of atomic number 8 and the life-supporting component of the air. Oxygen forms about 20 percent of the earth's atmosphere, and is the most abundant element in the earth's crust, mainly in the form of oxides, silicates, and carbonates. (Symbol: O) DERIVATIVES: ox·yg·e·nous / äkˈsijənəs/ adj.

oxygen Symbol O. A colourless odourless gaseous element. It is the most abundant element in the earth's crust (49.2% by weight) and is present in the atmosphere (28% by volume), mainly as dioxygen (O₂), with much smaller amounts of ozone (O₃). Atmospheric dioxygen is of vital importance for all organisms that carry out aerobic respiration.

oxygen XVIII. — F. oxygène, intended to mean ‘acidifying principle’ (acid-producer), f. Gr. oxús sharp, acid + -gène -GEN.

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