Oxygen Group Elements

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The oxygen group elements and their chemical symbols are oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po). They show great variation in their properties, from gas to solid, from life-sustaining to toxic, and from common to rare and radioactive.

Overview

Oxygen exists as a pure element in the air, as a binary compound in water, and as an element in countless minerals, especially in silicates. It is by far the most abundant element on Earth, and with its compounds makes up almost half the mass of the earth's crust. As far as can be inferred, it is the third most abundant element in the universe, after hydrogen and helium.

The oxygen group elements show as wide a range of behavior as perhaps any other group. To put the matter in perspective, consider these elements as they stand between groups V and VII:

Here, as throughout much of the periodic table, the lighter elements are far more familiar. They are also more abundant in nature.

Electronegativity increases as one moves toward fluorine at the upper right of the periodic table. It is the least metallic of all the elements. In contrast, electronegativity decreases so rapidly moving down and to the left of the chart that the elements at the bottom of groups V and VI are metals. Near them are three semimetals: arsenic, antimony, and tellurium. These three and selenium are also semiconductors. They resemble metals in appearance, but favor the nonmetals above them in chemical behavior.

The contrast in acid-base behavior among these fifteen elements is striking. Group VII contains elements that make three simple, strong acids: HCl, HBr, and HI. Some of the oxygen-containing acids of these elements, such as HClO₄, are also strong. At the lower left of the periodic table, bismuth forms an oxide that is weakly basic. Thus, group VI is in a part of the periodic table characterized by rapid change. With strong acids on the right and basic behavior on the lower left, telluric acid, TeOH₆, is caught in the middle, and is a weak acid.

All three of the groups begin with a gas and end with solids. Bromine (an element in group VII) is the only element that is liquid at room temperature. Both of the elements above bromine are gases.

The radii of the atoms increase greatly in going down group VI; polonium is about 2.3 times as large as oxygen; for tellurium, the ratio is almost 2.1:1. These size differences are roughly comparable with the ratios in the neighboring groups of the periodic table. As the atoms become larger, they are able to attach an increasing number of other atoms to themselves; thus, oxygen never attaches more than three atoms to itself, and when it does, they are hydrogens in acid solution. Nevertheless, sulfur, selenium, and tellurium can all attach up to six atoms to themselves. In group V, nitrogen can never attach more than four atoms to itself, but all the elements directly under it can accommodate six atoms. Elements of group VII behave similarly.

These trends are partly a result of the larger sizes of the heavier atoms.

The oxygen group elements are remarkable for their allotropes. Oxygen exists as the familiar O₂, but also as ozone, O₃. Ozone is made commercially by passing an electric discharge through oxygen; lightning also makes ozone. It is more reactive than O₂, and is thus used when combination with oxygen is required and the reaction cannot take place with ordinary O₂.

Airborne bacteria are killed by ozone. The largest natural concentration of ozone is found in the stratosphere, where it shields the earth's surface from harmful ultraviolet radiation. Liquid oxygen has a faint blue color; solid ozone is dark blue.

For solid sulfur, the commonest forms have six or eight atoms per molecule and are cyclic; that is, the atoms are attached to one another so as to make a closed ring. The various forms are a result of different methods of preparation. The most common is S₈, which has three crystalline forms--that is, three arrangements of atoms in a space-filling pattern. One of these, the most stable around room temperature, is commonly called rhombic sulfur because of its crystalline form. The second form of S₈, called monoclinic sulfur, is formed if rhombic sulfur is warmed to 95.5 degrees Celsius. If the heating is rapid, the solid melts instead at a temperature of 113 degrees Celsius. A second monoclinic form of S₈ is prepared in a solution of ethanol. Other solid forms with six, seven, nine, ten, and twelve atoms per molecule are known; all are more reactive than rhombic sulfur.

The behavior of liquid sulfur is also complicated. Whereas most liquids flow more easily with increasing temperature, sulfur does the opposite up to 200 degrees Celsius, while the sulfur molecules begin to break apart. The unattached atoms at the ends of the broken rings can bond to one another to make longer molecules called polymers. The specific heat also reaches its greatest value at about 200 degrees Celsius, then falls as the normal boiling point of 445 degrees Celsius is approached. This is highly unusual.

Gaseous sulfur is known in molecules of two, four, six, and eight atoms. As the temperature increases, the S₂ becomes the preferred molecule. The great variety of allotropic forms of sulfur results from its ability to form strong bonds to itself. Oxygen is much less able to do so, as it forms only two allotropes.

Selenium has several allotropes; a gray form is most stable, and two others are red. One of the red allotropes resembles sulfur, in having eight atoms per molecule. Again, the form of the element depends on how it is prepared. Gray selenium is photoconductive.

Tellurium is a silvery solid with only one structure, the same as gray selenium.

Polonium has two allotropic forms; conversion of one to the other occurs around 100 degrees Celsius. The low-temperature form is structurally unique among all elements in that it has one atom at every corner of a cube and none anywhere else. This pattern is repeated throughout the crystal.

The stable elements of group VI have the following atomic masses in grams per mole: oxygen, 15.999; sulfur, 32.066; selenium, 78.96; tellurium, 127.60. The atomic masses of all these elements are averages of the masses of the stable isotopes found in nature. The fact that oxygen is so close to an integer and tellurium so far from one is understood by looking at the isotopes of these elements. For oxygen, nature provides only three; their masses in grams per mole are almost exactly the same as the total number of neutrons plus protons in each of these nuclei, that is, 16, 17, and 18. The natural abundance of the last two of these is so small, however, that the average mass of a large number of oxygen atoms is almost exactly 16, the mass of the lightest isotope. In sulfur, the isotopes and their natural abundances are: mass 32, 95 percent; mass 33, 0.8 percent; mass 34, 4.2 percent; and a trace at mass 36. In this case, the contribution of the mass 34 isotope is enough to pull the average mass of sulfur atoms noticeably above 32. In the case of selenium, five isotopes all contribute more than 7 percent to the total natural abundance, so that the nearness of the atomic mass to an integer is coincidental.

Tellurium has an even wider distribution of isotopes: three main isotopes contribute 19, 32, and 34 percent to the total, and the final atomic mass, the average of these and other numbers, is not close to an integer. Polonium is entirely radioactive; no stable isotopes exist. In these patterns, the group VI elements illustrate typical behavior among all elements: With some exceptions, the number of stable isotopes of an element increases at first going down a group, but not all the way to the bottom of the periodic table. Elements heavier than bismuth have no stable isotopes.

The chemistry of all elements depends primarily on the electrons that individual atoms are least able to hold. Such electrons are said to be in the valence shell of that atom, and only they can enter into chemical bonds. The elements of group VI all have electron configurations that end with four electrons in p orbitals. Such a configuration is written np to the power of 4, where n is the principal quantum number. This important common feature means that their chemical behavior is sometimes predictable: When heated, sulfur, selenium, and tellurium will all burn in air to give dioxides. They all form compounds similar to water, with two hydrogen atoms to one group VI atom. Nevertheless, H₂S, H₂Se, and H₂Te are all foul-smelling and toxic; chemistry is too complicated to permit casual predictions.

These elements also show striking differences in their chemical behavior, as is the case with groups just to their left in the periodic table. Oxygen is distinct from the others in many ways; the most obvious is that oxygen is a gas, while the others are solids at room temperature.

Also, oxygen readily forms compounds with almost all other elements, which is not the case for the other members of the group. Such differences can be explained by noting that the electronegativity decreases going down any group. Another difference is that, as the atoms become heavier, they become larger, and more neighboring atoms can be arranged around a large atom than around a small one. Thus, the atoms lower in a group will tend to form bonds to more atoms than will the atoms near the top of the same group.

The numbering of group VI elements is drawn from a method of numbering the columns of the periodic table that has been standard for many years. Yet, another numbering system may be implemented, in which case the column of elements beginning with oxygen will be called group 16.

Applications

The utility of the oxygen group elements in technology and general manufacturing varies from oxygen itself, which is among the most important elements, to tellurium, for which only a few specialized uses are known. Oxygen, obtained by fractional distillation of air, is the third most widely used industrial chemical in the United States. More than 15 million tons of oxygen are produced per year. Vast quantities are used in the making of steel. As one of the gases consumed in the oxyacetylene torch, it has many further uses in metal manufacturing. It is used in the production of other industrially important substances, such as sulfuric acid and ethylene glycol (antifreeze). Hospitals use oxygen in treating patients with breathing difficulties.

Oxygen-containing compounds are also enormously important; it is in thirty of the fifty most important industrial chemicals, which include every leading acid except hydrochloric and every leading base except ammonia.

Sulfur is less widely used as an element than is oxygen. Among its important uses is the hardening of rubber. When heated with rubber, it acts to join neighboring molecules through crosslinks. For many years, sulfuric acid has been the industrial chemical produced in greatest quantity in the United States; economists regard its production as one of the basic measures of industrial advancement in any country. This results from the fact that sulfuric acid is used to make many other leading products. Among these are fertilizer, detergents, explosives, paint, paper, and refined petroleum. Sulfur dioxide is used to sterilize dried fruit; it is especially toxic to fungi. Sulfur hexafluoride, a heavy gas, is used as an insulator in electrical equipment. A sulfur-containing group of atoms is prominent in the sulfa drugs. A hot alkaline solution of sodium sulfate dissolves the wood from which paper is made.

Sulfur is found as the free element in several major underground deposits in the Northern Hemisphere, notably along the Gulf Coast in the United States. Sulfur's low melting point, 112 degrees Celsius, led to the development of the Frasch process for its extraction. Three concentric pipes are required; superheated steam is forced down the outermost pipe, which melts the sulfur. Compressed air is then forced down the innermost pipe and forces the sulfur to the surface through the pipe in between, where it remains liquid because of the high temperature of the steam pipe. Geochemists suggest that these deposits of elemental sulfur resulted from the action of bacteria on mineral sulfates. Sulfur is also found near volcanoes and hot springs. Other common forms of it are found in mineral sulfides of antimony, iron, lead, mercury, and zinc, and as the sulfate of magnesium (Epsom salt) and calcium (gypsum). It is cheaper than the metals in its common sulfides and thus is a frequent by-product of metal production.

Selenium has far fewer uses than either oxygen or sulfur. Its main industrial use is a consequence of its photoconductivity: When light falls on selenium, its normally poor electrical conductivity greatly increases. This is exploited in xerography. A photocopier contains a belt of selenium. This belt is given an electrical charge; then it is exposed to the image of a printed page.

The bright areas on the belt lose their charge as they become photoconductive. A black powder is attracted to the charged part of the belt, and this forms an image of the printed page. A blank sheet of paper is warmed and brought into brief contact with the belt. The toner transfers to the paper, and the heat causes the toner to bond to the paper, transferring the original image.

Selenium also gives a bright red color to glass, and is used to make stoplight lenses. Curiously, the same element in low concentrations decolorizes glass.

Tellurium is a silvery-white solid that gives desirable properties to a few special-purpose alloys. Lead containing a small amount of it is stronger, harder, and more resistant to the corrosive action of sulfuric acid. Tellurium is also important in blasting caps.

Selenium and tellurium are much less abundant than oxygen and sulfur. In this, they follow a general trend among all elements, according to which the heavier ones are less abundant. Selenium and tellurium are usually prepared as by-products of the refinement of copper. Polonium is a natural radioactive mineral present in trace amounts in pitchblende, the main source of uranium.

Polonium is intensely radioactive, emitting α particles five thousand times as fast as an equal mass of radium. Its only known use is as a source of α particles in nuclear reactions. A block of the material becomes hot because of the absorption of its own radiation, and this makes it potentially useful as a heat source for the production of electricity in satellites.

Its radiation has no known medical uses.

In their importance to life, oxygen group elements range from essential to useless. It is widely known that oxygen is fundamental to life, both as the gas and in water. Sulfur is found in cysteine and methionine, two amino acids, and as a result is present in many proteins. Selenium is required in the human diet as a trace element at concentrations around 100 parts per billion; this amount is so small that it was not recognized for many years.

Selenium, however, is toxic at higher concentrations. Some other elements and compounds show this property. Locoweed contains enough selenium to be harmful to animals that graze on it. Selenium hydride, H₂Se, and tellurium hydride, H₂Te, are both irritating gases with disagreeable odors. Taken into human bodies, they give a garlic odor to the breath.

Tellurium and polonium are not needed by living systems as far as is known, and the radioactivity of the latter rules out any conceivable benefit from it.

Oxygen group elements or their simple compounds vary greatly in their toxic effects.

Ozone begins to attack human tissue at concentrations above a few parts per million. It is a powerful oxidizing agent. Sulfur is a generally unreactive yellow solid, but hydrogen sulfide (H₂S), sulfur dioxide (SO₂), and especially sulfur trioxide (SO₃) are all dangerous gases. The toxicity of sulfur trioxide is quickly appreciated by noting that it reacts with water to form sulfuric acid. Thus, any moist tissue such as an eye surface quickly becomes a sulfuric acid solution. Hydrogen sulfide has the smell associated with rotten eggs and is released by the action of bacteria that act on organic matter in the absence of air.

Context

Among the elements in the oxygen group, sulfur was known in antiquity as brimstone, literally "burning stone." An early reference to it is in Genesis 19:24. Oxygen, the second of the group to be isolated, was prepared independently by Joseph Priestley and Karl Wilhelm Scheele in the 1770's. Priestley, though he probably did not think of oxygen as an element in the modern sense of the term, firmly established its importance to life by experiments with mice. He also observed the greater rate of burning of a flame in pure oxygen and speculated that human lives might also pass more quickly if we were to breathe pure oxygen. He declared that "the air which nature has provided for us is as good as we deserve."

Selenium was first isolated in 1817 by the Swedish chemist Jons Jakob Berzelius, who stimulated the development of modern chemistry by writing element symbols as they are known today in place of the alchemists' strange symbols. Tellurium was discovered by Franz Muller in 1782 but was not isolated until 1798. Marie Curie discovered polonium in 1898, her first radioactive element, which she named after her native Poland.

Bibliography

Bailar, John C., Jr., et al. CHEMISTRY. San Diego: Harcourt Brace Jovanovich, 1989. The authors, senior inorganic chemists, present much descriptive information in chapters 16, 28, and 29.

Chang, Raymond, and Wayne Tikkanen. THE TOP FIFTY INDUSTRIAL CHEMICALS. New York: Random House, 1988. The book is based on the analysis of the chemical industry published each year by CHEMICAL AND ENGINEERING NEWS. The fifty short chapters typically include background, industrial preparation, properties, uses, production data, and price. A brief appendix is included on the petroleum industry. Useful for browsing or reference.

Chemical Rubber Company. CRC HANDBOOK OF CHEMISTRY AND PHYSICS. Boca Raton, Fla.: CRC Press, 1991. Annual editions are a valuable source of information on the elements, their properties, their histories, and their uses.

Ebbing, Darrell D. GENERAL CHEMISTRY. 3d ed. Boston: Houghton Mifflin, 1990. A well-organized presentation with beautiful illustrations. Chapter 22 is useful.

Jaffe, Bernard. CRUCIBLES: THE STORY OF CHEMISTRY. 4th ed. New York: Dover, 1976. Chapter 4 discusses Priestley and his experiments on oxygen. The most popular history of chemistry, from alchemy to nuclear fission.

Kotz, John C., and Keith F. Purcell. CHEMISTRY AND CHEMICAL REACTIVITY. Philadelphia: Saunders College Publishing, 1987. The table of properties of the group VI elements on page 837 is particularly useful.

McQuarrie, Donald A., and Peter A. Rock. GENERAL CHEMISTRY. 3d ed. New York: W. H. Freeman, 1991. Page 926 and section 26-2 deal with the oxygen group elements. The discussion of sulfur and accompanying illustrations are exceptional.

Rochow, Eugene G. MODERN DESCRIPTIVE CHEMISTRY. Philadelphia: W. B. Saunders, 1977. An unusual, highly readable book with an industrial orientation that distinguishes it from many other titles.

Windholz, Martha, ed. THE MERCK INDEX. 10th ed. Rahway, N.J.: Merck, 1983. A detailed reference work of chemicals, drugs, and biological substances with ten thousand entries, plus many useful tables. Information for a typical compound includes synonyms, formula, molecular structure, literature citation to the synthesis, basic physical and chemical properties, therapeutic uses, and toxicity.

Elements in Groups V, VI, and VII

Principal terms

Acids and Bases

The Periodic Table and the Atomic Shell Model

ALLOTROPES: different forms of an element in the same state, caused by different molecular or crystalline structure

ATOM: the smallest and simplest unit of an element that can exist

COVALENT BOND: a bond formed by the sharing of electrons by atoms

ELECTRON: the smallest unit of electrical charge; an atom consists of a nucleus and the electrons around it

ELECTRON CONFIGURATION: the arrangement of electrons in a given atom according to their energies and the volumes they occupy around the atomic nuclei

ELECTRONEGATIVITY: the attraction an atom has for a pair of electrons it is sharing with a neighboring atom in a covalent bond

ISOTOPES: atoms of the same element that differ in the number of neutrons in the nucleus

MOLECULE: two or more atoms that are held together strongly enough to be identified and studied as a unit