chemical element

Actinium (Ac)

Actinium (Ac), radioactive chemical element, in Group 3 (IIIb) of the periodic table, atomic number 89. Actinium was discovered (1899) by French chemist André-Louis Debierne in pitchblende residues left after French physicists Pierre and Marie Curie had extracted radium from them, and it was also discovered (1902) independently by German chemist Friedrich Oskar Giesel. Debierne named the element after the Greek word aktinos (“ray”). A ton of pitchblende ore contains about 0.15 mg of actinium. The rare silvery-white metal is highly radioactive, glowing blue in the dark.

The most common isotope of actinium is actinium-227; the others, natural and artificial, are too short-lived to accumulate in macroscopic quantity. Actinium-227, which is one of the decay products of uranium-235, has a 21.8-year half-life and in turn decays almost entirely to thorium-227, but about 1 percent decays to francium-223. This whole disintegration chain with its branches is called the actinium series.

Actinium-225 has a 10-day half-life, decaying by the emission of alpha particles. Its short-lived daughter isotopes emit only alpha and beta particles with no high-energy gamma rays. This isotope can thus deliver high-energy radiation to a tumour without greatly affecting the surrounding tissue. Complexes of actinium-225 have been studied for their use in nuclear medicine.

Aluminum (Al)

Aluminum (Al), also spelled aluminium, chemical element, a lightweight silvery white metal of main Group 13 (IIIa, or boron group) of the periodic table. Aluminum is the most abundant metallic element in Earth’s crust and the most widely used nonferrous metal. Because of its chemical activity, aluminum never occurs in the metallic form in nature, but its compounds are present to a greater or lesser extent in almost all rocks, vegetation, and animals. Aluminum is concentrated in the outer 16 km (10 miles) of Earth’s crust, of which it constitutes about 8 percent by weight; it is exceeded in amount only by oxygen and silicon. The name aluminum is derived from the Latin word alumen, used to describe potash alum, or aluminum potassium sulfate, KAl(SO₄)₂∙12H₂O.

Occurrence and history

Aluminum occurs in igneous rocks chiefly as aluminosilicates in feldspars, feldspathoids, and micas; in the soil derived from them as clay; and upon further weathering as bauxite and iron-rich laterite. Bauxite, a mixture of hydrated aluminum oxides, is the principal aluminum ore. Crystalline aluminum oxide (emery, corundum), which occurs in a few igneous rocks, is mined as a natural abrasive or in its finer varieties as rubies and sapphires. Aluminum is present in other gemstones, such as topaz, garnet, and chrysoberyl. Of the many other aluminum minerals, alunite and cryolite have some commercial importance.

Before 5000 bce people in Mesopotamia were making fine pottery from a clay that consisted largely of an aluminum compound, and almost 4,000 years ago Egyptians and Babylonians used aluminum compounds in various chemicals and medicines. Pliny refers to alumen, now known as alum, a compound of aluminum widely employed in the ancient and medieval world to fix dyes in textiles. In the latter half of the 18th century, chemists such as Antoine Lavoisier recognized alumina as the potential source of a metal.

Uses and properties

Aluminum is added in small amounts to certain metals to improve their properties for specific uses, as in aluminum bronzes and most magnesium-base alloys; or, for aluminum-base alloys, moderate amounts of other metals and silicon are added to aluminum. The metal and its alloys are used extensively for aircraft construction, building materials, consumer durables (refrigerators, air conditioners, cooking utensils), electrical conductors, and chemical and food-processing equipment.

Pure aluminum (99.996 percent) is quite soft and weak; commercial aluminum (99 to 99.6 percent pure) with small amounts of silicon and iron is hard and strong. Ductile and highly malleable, aluminum can be drawn into wire or rolled into thin foil. The metal is only about one-third as dense as iron or copper. Though chemically active, aluminum is nevertheless highly corrosion-resistant, because in air a hard, tough oxide film forms on its surface.

Aluminum is an excellent conductor of heat and electricity. Its thermal conductivity is about one-half that of copper; its electrical conductivity, about two-thirds. It crystallizes in the face-centred cubic structure. All natural aluminum is the stable isotope aluminum-27. Metallic aluminum and its oxide and hydroxide are nontoxic.

Aluminum is slowly attacked by most dilute acids and rapidly dissolves in concentrated hydrochloric acid. Concentrated nitric acid, however, can be shipped in aluminum tank cars because it renders the metal passive. Even very pure aluminum is vigorously attacked by alkalies such as sodium and potassium hydroxide to yield hydrogen and the aluminate ion. Because of its great affinity for oxygen, finely divided aluminum, if ignited, will burn in carbon monoxide or carbon dioxide with the formation of aluminum oxide and carbide, but, at temperatures up to red heat, aluminum is inert to sulfur.

Aluminum can be detected in concentrations as low as one part per million by means of emission spectroscopy. Aluminum can be quantitatively analyzed as the oxide (formula Al₂O₃) or as a derivative of the organic nitrogen compound 8-hydroxyquinoline. The derivative has the molecular formula Al(C₉H₆ON)_3.

Americium

Americium (Am), synthetic chemical element (atomic number 95) of the actinoid series of the periodic table. Unknown in nature, americium (as the isotope americium-241) was artificially produced from plutonium-239 (atomic number 94) in 1944 by American chemists Glenn T. Seaborg, Ralph A. James, Leon O. Morgan, and Albert Ghiorso in a nuclear reactor. It was the fourth transuranium element to be discovered (curium, atomic number 96, was discovered a few months previously). The element was named after the United States of America.

The metal is silvery white and tarnishes slowly in dry air at room temperature. The isotope americium-241 is the most important because of its availability. This isotope is produced by multiple neutron capture in nuclear reactors and has been isolated in kilogram amounts from plutonium and other actinoids in used nuclear fuel. Americium-241 has been used industrially in fluid-density gauges, thickness gauges, aircraft fuel gauges, and distance-sensing devices, all of which use its gamma radiation. The isotope’s alpha-particle emission is exploited in smoke detectors. All isotopes of americium are radioactive; the stablest isotope, americium-243, has proved more convenient for chemical investigations because of its longer half-life (7,370 years, compared with 433 years for americium-241).

Americium reacts with oxygen to form the dioxide AmO₂, with halogen elements to form compounds such as the tetrafluoride AmF₄ and all the trihalides, and with hydrogen to form the hydride AmH_2+x. Americium has four well-characterized oxidation states, from +3 to +6, in acidic aqueous solution with the following ionic species: Am³⁺, pink; Am⁴⁺, rose (very unstable); AmO₂ ⁺, yellow; and AmO₂²⁺, light tan. In the common +3 state, americium is very similar to the other actinoid and lanthanoid elements. There is some evidence that the ion Am²⁺ has been prepared in trace amounts; its existence suggests that americium is similar to its lanthanoid homologue, europium, which can be reduced to its +2 oxidation state. There is also evidence for heptavalent americium in strongly basic aqueous solution.

Antimony

Antimony (Sb), a metallic element belonging to the nitrogen group (Group 15 [Va] of the periodic table). Antimony exists in many allotropic forms (physically distinct conditions that result from different arrangements of the same atoms in molecules or crystals). Antimony is a lustrous, silvery, bluish white solid that is very brittle and has a flaky texture. It occurs chiefly as the gray sulfide mineral stibnite (Sb₂S₃).

History

The ancients were familiar with antimony both as a metal and in its sulfide form. Fragments of a Chaldean vase made of antimony have been estimated to date from about 4000 bc. The Old Testament tells of Queen Jezebel using the naturally occurring sulfide of antimony to beautify her eyes. Pliny, during the 1st century ad, wrote of seven different medicinal remedies using stibium or antimony sulfide. Early writings of Dioscorides, dating from about the same time, mention metallic antimony. Records of the 15th century show the use of the substance in alloys for type, bells, and mirrors. In 1615 Andreas Libavius, a German physician, described the preparation of metallic antimony by the direct reduction of the sulfide with iron; and a later chemistry textbook by Lémery, published in 1675, also describes methods of preparation of the element. In the same century, a book summarizing available knowledge of antimony and its compounds was purportedly written by a Basil Valentine, allegedly a Benedictine monk of the 15th century, whose name appears on chemical writings over a span of two centuries. The name antimony appears to be derived from the Latin antimonium, in a translation of a work by the alchemist Geber, but its real origin is uncertain.

Occurrence and distribution

Antimony is about one-fifth as abundant as arsenic, contributing on the average about one gram to every ton of the Earth’s crust. Its cosmic abundance is estimated as about one atom to every 5,000,000 atoms of silicon. Small deposits of native metal have been found, but most antimony occurs in the form of more than 100 different minerals. The most important of these is stibnite, Sb₂S₃. Small stibnite deposits are found in Algeria, Bolivia, China, Mexico, Peru, South Africa, and in parts of the Balkan Peninsula. Some economic value also attaches to kermesite (2Sb₂S₃ · Sb₂O₃), argentiferous tetrahedrite [(Cu,Fe)₁₂Sb₄S₁₃], livingstonite (HgSb₄S₇), and jamesonite (Pb₄FeSb₆S₁₄). Small amounts are also recoverable from the production of copper and lead. About half of all the antimony produced is reclaimed from scrap lead alloy from old batteries, to which antimony had been added to provide hardness.

Two stable isotopes, nearly equal in abundance, occur in nature. One has mass 121 and the other mass 123. Radioactive isotopes of masses 120, 122, 124, 125, 126, 127, 129, and 132 have been prepared.

Commercial production and uses

High-grade or enriched stibnite reacts directly with scrap iron in the molten state, liberating antimony metal. The metal can also be obtained by conversion of stibnite to the oxide, followed by reduction with carbon. Sodium sulfide solutions are effective leaching agents for the concentration of stibnite from ores. Electrolysis of these solutions produces antimony. After further purification of the crude antimony, the metal, called regulus, is cast into cakes.

About half of this antimony is used metallurgically, principally in alloys. Because some antimony alloys expand on solidifying (a rare characteristic that they share with water), they are particularly valuable as castings and type metal; the expansion of the alloy forces the metal to fill the small crevices of casting molds. Moreover, the presence of antimony in type metal, which also includes lead and small amounts of tin, increases the hardness of the type and gives it a sharp definition. Even when added in minor quantities, antimony imparts strength and hardness to other metals, particularly lead, with which it forms alloys used in plates of automobile storage batteries, in bullets, in coverings for cables, and in chemical equipment such as tanks, pipes, and pumps. Combined with tin and lead, antimony forms antifriction alloys called babbitt metals that are used as components of machine bearings. With tin, antimony forms such alloys as britannia metal and pewter, used for utensils. Antimony is also used as an alloy in solder. Highly purified antimony is used in semiconductor technology to prepare the intermetallic compounds indium, aluminum, and gallium antimonide for diodes and infrared detectors.

Antimony compounds (especially the trioxide) are widely used as flame retardants in paints, plastics, rubber, and textiles. Several other antimony compounds are used as paint pigments; tartar emetic (an organic salt of antimony) is used in the textile industry to aid in binding certain dyes to fabrics and in medicine as an expectorant and a nauseant.

Biological and physiological significance

Antimony and a number of its compounds are highly toxic. In fact, the use of antimony compounds for medicinal purposes was temporarily outlawed several centuries ago because of the number of fatalities they had caused. A hydrated potassium antimonyl tartrate called “tartar emetic” is currently used in medicine as an expectorant, diaphoretic, and emetic. The maximum tolerable concentration of antimony dust in air is about the same as for arsenic, 0.5 milligrams per cubic metre.

Argon

argon (Ar), chemical element, inert gas of Group 18 (noble gases) of the periodic table, terrestrially the most abundant and industrially the most frequently used of the noble gases. Colourless, odourless, and tasteless, argon gas was isolated (1894) from air by the British scientists Lord Rayleigh and Sir William Ramsay. Henry Cavendish, while investigating atmospheric nitrogen (“phlogisticated air”), had concluded in 1785 that not more than ¹/₁₂₀ part of the nitrogen might be some inert constituent. His work was forgotten until Lord Rayleigh, more than a century later, found that nitrogen prepared by removing oxygen from air is always about 0.5 percent more dense than nitrogen derived from chemical sources such as ammonia. The heavier gas remaining after both oxygen and nitrogen had been removed from air was the first of the noble gases to be discovered on Earth and was named after the Greek word argos, “lazy,” because of its chemical inertness. (Helium had been spectroscopically detected in the Sun in 1868.)

In cosmic abundance, argon ranks approximately 12th among the chemical elements. Argon constitutes 1.288 percent of the atmosphere by weight and 0.934 percent by volume and is found occluded in rocks. Although the stable isotopes argon-36 and argon-38 make up all but a trace of this element in the universe, the third stable isotope, argon-40, makes up 99.60 percent of the argon found on Earth. (Argon-36 and argon-38 make up 0.34 and 0.06 percent of Earth’s argon, respectively.) A major portion of terrestrial argon has been produced, since the Earth’s formation, in potassium-containing minerals by decay of the rare, naturally radioactive isotope potassium-40. The gas slowly leaks into the atmosphere from the rocks in which it is still being formed. The production of argon-40 from potassium-40 decay is utilized as a means of determining Earth’s age (potassium-argon dating).

Argon is isolated on a large scale by the fractional distillation of liquid air. It is used in gas-filled electric light bulbs, radio tubes, and Geiger counters. It also is widely utilized as an inert atmosphere for arc-welding metals, such as aluminum and stainless steel; for the production and fabrication of metals, such as titanium, zirconium, and uranium; and for growing crystals of semiconductors, such as silicon and germanium.

Arsenic

Arsenic (As), a chemical element in the nitrogen group (Group 15 [Va] of the periodic table), existing in both gray and yellow crystalline forms.

History

Arsenic was known in the form of certain of its compounds long before it was clearly recognized as a chemical element. In the 4th century bce Aristotle wrote of a substance called sandarache, now believed to have been the mineral realgar, a sulfide of arsenic. Then, in the 1st century ce, the writers Pliny the Elder and Pedanius Dioscorides both described auripigmentum, a substance now thought to have been the dyestuff orpiment, As₂S₃. By the 11th century ce three species of “arsenic” were recognized: white (As₄O₆), yellow (As₂S₃), and red (As₄S₄). The element itself possibly was first observed in the 13th century by Albertus Magnus, who noted the appearance of a metal-like substance when arsenicum, another name for As₂S₃, was heated with soap. It is not certain, however, that this natural scientist and scholar actually observed the free element. The first clearly authentic report of the free substance was made in 1649 by Johann Schroeder, a German pharmacist, who prepared arsenic by heating its oxide with charcoal. Later, Nicolas Lémery, a French physician and chemist, observed the formation of arsenic when heating a mixture of the oxide, soap, and potash. By the 18th century, arsenic was well known as a unique semimetal.

Occurrence and distribution

The abundance of arsenic in the Earth’s crust is about five grams per ton; the cosmic abundance is estimated as about four atoms per million atoms of silicon. The element is widely distributed. A small amount exists in the native state, in 90–98 percent purity, generally in association with such metals as antimony and silver. Most, however, is combined in more than 150 different minerals, as sulfides, arsenides, sulfoarsenides, and arsenites. Mispickel, or arsenopyrite, FeAsS, is among the most common of arsenic-bearing minerals; others are realgar, As₄S₄; orpiment, As₂S₃; loellingite, FeAs₂; and enargite, Cu₃AsS₄. Arsenic oxide is also common. Most commercial arsenic is recovered as a by-product of the smelting of copper, lead, cobalt, and gold ores.

Astatine

astatine (At), radioactive chemical element and the heaviest member of the halogen elements, or Group 17 (VIIa) of the periodic table. Astatine, which has no stable isotopes, was first synthetically produced (1940) at the University of California by American physicists Dale R. Corson, Kenneth R. MacKenzie, and Emilio Segrè, who bombarded bismuth with accelerated alpha particles (helium nuclei) to yield astatine-211 and neutrons. Naturally occurring astatine isotopes have subsequently been found in minute amounts in the three natural radioactive decay series, in which they occur by minor branching (astatine-218 in the uranium series, astatine-216 in the thorium series, and astatine-215 and astatine-219 in the actinium series). Thirty-two isotopes are known; astatine-210, with a half-life of 8.1 hours, is the longest lived. Because astatine has no stable or long-lived isotopes, it was given its name from the Greek word astatos, meaning “unstable.”

Production and use

The only practical way of obtaining astatine is by synthesizing it through nuclear reactions.