Group 6A — The Chalcogens

1A 2A 3A 4A 5A 6A 7A 8A
(1) (2) (13) (14) (15) (16) (17) (18)
3B 4B 5B 6B 7B 8B 1B 2B
(3) (4) (5) (6) (7) (8) (9) (10) (11) (12)
1 H He
2 Li Be B C N O F Ne
3 Na Mg Al Si P S Cl Ar
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 Cs Ba La   Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
7 Fr Ra Ac   Rf Db Sg Bh Hs Mt Ds Rg Uub Uuq
6   Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
7   Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

 

Group 6A (or VIA) of the periodic table are the chalcogens:  the nonmetals oxygen (O), sulfur (S), and selenium (Se), the metalloid tellurium (Te), and the metal polonium (Po).  The name "chalcogen" means "ore former," derived from the Greek words chalcos ("ore") and -gen ("formation").

The Group 6A elements have six valence electrons in their highest-energy orbitals (ns2np4).  This is only two electrons away from having a full octet of eight electrons, so many of these elements form anions having -2 charges:  oxide, O2-; sulfide, S2-, selenide, Se2-, etc.  In combination with other nonmetals, oxygen, sulfur, and selenium form compounds through covalent bonding.

 

Oxygen (O, Z=8).

Oxygen is a colorless, odorless, tasteless, and extremely reactive gas; in its elemental form, oxygen is found as the diatomic molecule O2; in the liquid phase, it is pale blue.  The name of the element is derived from the Latin words oxy and genes, meaning "acid forming" (see explanation below).  Oxygen is the most abundant element in the Earth's crust (47%), the second most abundant element in the atmosphere (21%), and the third most abundant element in the universe (for every atom of silicon, there are 22 atoms of oxygen, 3,100 of helium, and 40,000 of hydrogen).

Oxygen forms compounds with all of the other element except for the lighter noble gases.  It is found in many mineral ores, in the form of oxides (O2-), sulfates (SO42-), nitrates (NO3-), phosphates (PO43-), carbonates (CO32-), silica and quartz (SiO2) and the silicates, etc.  In the atmosphere, it is found in its elemental form, which is a diatomic molecule, O2, consisting of two atoms of oxygen joined by a double bond.  (Actually, the bonding in O2 is quite a bit more complicated than that, because the second pair of electrons in the "double bond" are not paired.)  When cooled to a temperature of -183C, oxygen forms a pale blue liquid, which is attracted to the poles of a very strong magnet.

Oxygen is essential for respiration; molecular oxygen taken up in the lungs is used to oxidize sugar molecules, resulting in the production of carbon dioxide, water, and a great deal of energy.  Photosynthesis in plants reverses this process; green plants and photosynthetic bacteria take in carbon dioxide, and use energy from sunlight to convert it into sugar molecules, releasing oxygen in the process.  Molecular oxygen was first released into the atmosphere by primitive photosynthetic bacteria starting about 3.5 to 3.8 billion years ago; when land plants evolved, the concentration of O2 in the atmosphere rose to about 20%.  Oxygen dissolves in water to the extent of about 1.2 milligrams per liter of water — that's not very much, but it is essential for aquatic life.

The burning of fuels in the presence of molecular oxygen is where we get a lot of our electricity, and how we power automobiles.  The hydrocarbons in petroleum and its refined products such as natural gas, gasoline, and diesel fuel, undergo combustion with oxygen to product carbon dioxide, water, and a great deal of energy.  The burning of coal to produce carbon dioxide also produces energy.

Complete combustion of hydrocarbons produces carbon dioxide; if there is insufficient oxygen present, carbon monoxide or elemental carbon can be produced.  (The yellow-orange color of burning wood comes from hot elemental carbon that is formed because the solid fuel does not mix very well with oxygen gas.)

Things are more flammable in pure oxygen than in the 21% oxygen of the normal atmosphere (see here for a demonstration).  Even iron, which does not burn in 21% oxygen, will burn in 100% oxygen (see here).  This led to a disaster in the early days of the American space program.  The capsules used in the Mercury, Gemini, and Apollo program operated in space with an atmosphere of pure oxygen at a pressure of 5 psi.  (Normal atmospheric pressure is about 14.7 psi.  The reduced pressure oxygen environment was used to eliminate the need to carry nitrogen tanks into space for the astronauts to breathe, in the belief that a two-component gas system would be more difficult to manage, since it would require the oxygen/nitrogen ratio to be calibrated precisely at all times to prevent the astronauts from suffocating.)  Since the spacecraft were to operate in a pure oxygen environment in space, they were tested on the ground in a pure oxygen environment.  There were no complications with these tests during the Mercury and Gemini programs of the early 1960s, but on January 27, 1967, a fire broke out during a routine test in the command module of what was to become the first mission of the Apollo program; the door opened inward, and rapidly become impossible to pull open against the pressure of the gases being generated in the fire.  Within 17 seconds, the astronauts Virgil Grissom, Ed White, and Roger Chaffee were killed.  As a result of the disaster, the Command Module was extensively redesigned to prevent such a tragedy from occurring again.

Oxygen is used in welding to generate the intense heat needed to cut and weld steel and other high-melting point metals.  It is used in hospitals to supply oxygen to patients who have difficulty breathing.  It is also used in rocket fuels:  the Saturn V rockets that launched the Apollo lunar missions used 209,000 gallons of kerosene and 334,500 gallons of liquid oxygen in its first stage (S-IC), and 260,000 gallons of liquid hydrogen and 83,000 gallons of liquid oxygen in its second stage (S-II), and 69,500 gallons of liquid hydrogen and 20,150 gallons of liquid oxygen in its third (S-IVB) stage; the Space Shuttle main engines use 385,000 gallons of liquid hydrogen and 143,000 gallons of liquid oxygen.

The discovery of oxygen is an extremely tangled story, partially because of questions of priority, and partially because of misunderstandings about the nature of combustion and the gas phase.  For thousands of years, air was considered to be an "element," and it was not recognized that air was actually a mixture of many different gases.  The nature of combustion was also hotly debated (pun intended); many scientists believed that flammable substances contained a material called phlogiston, which was released when a substance burned.  When nitrogen was discovered in 1772, it was referred to as "phlogisticated air," since an atmosphere of pure nitrogen (actually, nitrogen plus carbon dioxide) did not support combustion.  (It was thought that this "air" had absorbed the maximum amount of phlogiston.)  Oxygen was discovered by the Swedish chemist Carl Wilhelm Scheele in 1772, but his account of his experiment was not published until 1777.  The English chemist Joseph Priestley produced oxygen in 1774 by heating a sample of merucry(II) oxide, HgO, and collecting the oxygen gas it produced over water.  He called the gas "dephlogisticated air," since it supported combustion more vigorously that "normal" air, and therefore presumably was more capable of "pulling" phlogiston out of other substances.  The French chemist Antoine Lavoisier claimed to have produced oxygen in 1774, independently of Priestley, but Priestley had visited him a few months before and told him of his experiment.  Lavoisier did, however, correctly interpret the significance of Priestley's result:  that combustion is the not release of phlogiston from a substance, but the combination of the substance with oxygen in the air, to produce oxides (as well as heat and light).  Lavoisier believed that the new element was an essential component of all acids, and proposed that it be called "oxygen," from the Greek words oxy, "acid" and genes "forming."  (However, not all acids contain oxygen; for example, hydrochloric acid, HCl.)

Another well-known form of oxygen is ozone, O3.  Ozone is a powerful oxidizing agent, and is often used to kill bacteria during the purification of water.  At sea level, ozone in the atmosphere is a pollutant, produced by the action of sunlight on nitrogen oxides in car exhaust.  In the stratosphere, at an elevation of 10 to 50 km about the surface of the Earth, ozone is produced by the action of sunlight upon O2, which splits apart into atomic oxygen, O, and combines with another O2 molecule to form ozone, O3.  The ozone absorbs high-energy ultraviolet light, splitting apart into O2 and O, which can then recombine and absorb another photon of high-energy light.  This ozone layer forms a shield which protects living organisms on the Earth's surface from this damaging, high-energy light.  The release of CFCs (chlorofluorocarbons) into the atmosphere produces chlorine radicals which are damaging to the ozone layer; for this reason, these substances are being phased out.  (See the Molecules pages on dichlorodifluoromethane for more information.)

One of the most important compounds of oxygen is water, H2O, which makes up nearly 75% of the Earth's surface.  Water freezes at 0C to form solid ice, which is less dense than liquid water.  (This is unusual for the solid form of a liquid substance, and one reason why if you're sailing in a ship in the North Atlantic, it's a good idea to keep a lookout for icebergs.)  Water boils at 100C, which again is unusual; most compounds of that low a molecular weight (water weighs 18.02 g/mol) are gases at room temperature.  These "anomalous" properties of water are in part a result of the large differences in electronegativity between oxygen and hydrogen — the oxygen-hydrogen bond is extremely polar, and water molecules attract each other much more strongly than most other small molecules do, as a results of these hydrogen bonds.  (See the entry on hydrogen for more on hydrogen bonds.)

 

Sulfur (S, Z=16).

Sulfur is a yellow nonmetal, and is found in a variety of forms, ranging from a yellow powder to more crystalline structures.  The name is derived from the ancient names for the element, either the Sanskrit word sulvere, the Latin word sulfurium, or the Arabic word sufra.  It is found in the Earth's crust at a concentration of 260 ppm, making it the 17th most abundant element.  Sulfur can be mined in its elemental form near volcanoes and hot springs, and is mined from salt domes along the Gulf of Mexico, Poland, and Russia (where it was produced by the action of bacteria on sulfate-containing minerals).  It is also present in many ores, such as stibnite [antimony sulfide, Sb2S3], galena [lead(II) sulfide, PbS], cinnabar [mercury(II) sulfide, HgS], sphalerite [zinc sulfide, ZnS], pyrite [iron(II) sulfide, FeS], anhydrite and gypsum [calcium sulfate, CaSO4], epsomite [magnesium sulfate, MgSO4], alunite [potassium aluminum sulfate, KAl(SO4)2], and barite [barium sulfate, BaSO4], among others.

Sulfur has been known in its elemental form since ancient times (often under the name "brimstone").  It is a yellow nonmetal, and is found in a number of allotropic forms, including orthorhombic sulfur, monoclinic sulfur, etc.  It forms rings containing anywhere from 4 to 20 atoms of sulfur; S8 is the most common form.  When heated to high temperatures, these rings break open, and join to form long chains; the resulting material is extremely viscous, and forms a rubbery solid called "plastic sulfur."  Sulfur is used in the manufacture of sulfuric acid, in making vulcanized rubber, in gunpowder and fireworks, etc.  Sulfur is present in proteins, in the form of the amino acids cysteine and methionine; on average 900 milligrams of sulfur are consumed every day in this manner.

Probably the most important compound of sulfur is sulfuric acid, H2SO4, which is the industrial chemical produced in the largest amounts (165 million tons in 2001).  Sulfuric acid is used in the production of phosphates for fertilizers, the removal of rust from iron, the production of explosives, paints, paper, detergents, dyes, in lead-acid car batteries, and many other uses.  Sulfuric acid has a high affinity for water, and is used as a dehydrating agent.  This can be easily demonstrated:  applying concentrated sulfuric acid to paper causes the paper to become black and charred, as if it had been burned; sulfuric acid also removes water from sugar, leaving behind a solid mass of carbon.  A common laboratory mistake is to mix sulfuric acid and water by adding water to the concentrated sulfuric acid:  the mixing of sulfuric acid and water can produce enough heat to boil the water, splattering the water and acid all over the incautious chemist.  It is a standing rule in all chemistry labs that, when diluting acids, always add the acid to the water, never the water to the acid.

Since sulfur is found in two amino acids (cysteine and methionine), some sulfur is present in fossil fuels such as coal.  When the coal is burned, the sulfur is oxidized to form sulfur dioxide, SO2, and sulfur trioxide, SO3; these gases react with moisture in the air, producing sulfurous acid, H2SO3, and sulfuric acid, H2SO4, respectively.  This leads to the formation of acid rain, which is a serious environmental pollutant in some areas.

Many compounds containing sulfur are bad news for the nose.  Hydrogen sulfide, H2S, is produced during the breakdown of organic matter by bacteria in the absence of oxygen; it is found is swamps, natural gas, volcanic gases, and many other sources.  It is detectable by humans at concentrations of 4.7 parts per billion or higher, and smells like rotten eggs.  Organic compounds containing sulfur are called thiols (also known as mercaptans for their ability to encapsulate mercury); thiols are responsible for the odor of skunks, the odor of cut onions and garlic, the smell of natural gas (which comes from ethanethiol, CH3CH2SH, which added to natural gas to make gas leaks detectable), bad breath (halitosis), and other similar odors.

Gunpowder was first made in China in about the year 950.  It is composed of charcoal, sulfur, and potassium nitrate.  (This mixture also comes in handy if you're being chased by a Gorn.  If you have to ask, you probably don't want to know.)  Sulfur was also probably present in Greek fire, along with tar, and antimony sulfide (see entry for antimony).  It is also present in mustard gas, and other poison gases and nerve agents.

 

Selenium (Se, Z=34).

Selenium is a soft metalloid, found in several allotropic forms.  The name of the element is derived from the Greek word for the Moon, selene, since tellurium had been named for the Earth.  It is found in the Earth's crust at a concentration of 50 ppb, making it the 67th most abundant element.  Selenium ores are rare, and it is usually found as selenide (Se2-) along with sulfide (S2-) in ores of copper, lead, zinc, and other metals.

Selenium conducts electricity much better in the presence of light than it does in the darkness; for this reason, it is used in photoelectric cells, light meters in cameras, solar cells, photocopiers, and other light-sensitive devices.  Selenium is essential in the diet in small amounts, but at levels of more than 5 mg, it can be toxic.  High levels of selenium in the body can produce foul breath and body odor, as a result of the production of volatile dimethylselenide, (CH3)2Se.

In areas where the soil is particularly rich in selenium (such as the Great Plains in the central United States), there are some plants that are particularly efficient at pulling selenium from the soil.  These plants include milkvetch (genus Astragalus) and several species of the genus Oxytropis.  When cattle eat these plants, the develop some of the symptoms of selenium poisoning, including erratic behavior, aggression, nervousness, lethargy, and loss of balance ("blind staggers").  These plants are often referred to as "loco weed" (from the Spanish word loco, "crazy").

 

Tellurium (Te, Z=52).

Tellurium is a metalloid; it has a characteristic metallic luster, but can be pulverized to form a gray powder.  The name of the element is derived from the Latin word for the Earth, tellus.  It is found in the Earth's crust at a concentration of 5 ppb, making it the 72nd most abundant element.  It is found in the ores calaverite and krennerite [gold telluride, AuTe2], sylvanite [silver gold telluride, (Ag,Au)Te2], petzite [Ag3AuTe2],and tellurite [TeO2], although it is usually obtained as a byproduct of the refining of copper and lead.

Tellurium is used to make alloys with stainless steel and copper to increase their machinability, and is sometimes used to vulcanize rubber.  It is also used in blasting caps.  Tellurium exposure can result in both bad breath and a foul body odor, as a result of the production of volatile dimethyltelluride, (CH3)2Te.

 

Polonium (Po, Z=84).

Polonium is a silvery, radioactive metal discovered by Marie Curie in 1898 (along with radium); she named the element after Poland, her native country.  Although the discovery of polonium was to some extent eclipsed by the discovery of radium, polonium is more radioactive than radium by a factor of about 5000.  It is found naturally in uranium ore in extremely small amounts; one ton of uranium ore contains about 100 micrograms (0.0001 g) of polonium, making it one of the ten least abundant naturally-occurring elements.  For industrial use (primarily to dissipate static electricity), it is produced from the irradiation of bismuth-209 with neutrons to produce bismuth-210, which undergoes beta-decay to produce polonium-210.  It is radioactive, and is regarded as one of the deadliest substances known.  Polonium is produced in the radioactive decay of radon gas; when the radon is inhaled, it can undergo alpha-decay to produce polonium, which is a solid, and remains trapped within the lungs.

 

 

References

John Emsley, The Elements, 3rd edition.  Oxford:  Clarendon Press, 1998.

John Emsley, Nature's Building Blocks:  An A-Z Guide to the Elements.  Oxford:  Oxford University Press, 2001.

David L. Heiserman, Exploring Chemical Elements and their Compounds.  New York:  TAB Books, 1992.