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H

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

Deuterium

>99,00

0,0115

Heavy Water

Hydrogen was discovered in 1766 by Henry Cavendish, who also independently discovered nitrogen. Its name derives from the Greek words hydro and genes, meaning “water-generator” or “water-creator.”

The most abundant element in the universe, hydrogen is a colorless, odorless, tasteless gas that is also flammable, burning in the air with a popping sound. It combines explosively with oxygen at ordinary temperatures in the presence of finely divided metals. Explosion of its mixture with chlorine is detonated by sunlight, heat or spark. It also combines explosively with halogens, and it autoignites at 574 ºC.

Hydrogen is lighter than air and is slightly soluble in water. It forms the largest number of chemical compounds — more than any other element, including carbon — and is a component of all mineral acids, ammonia, natural gases and hydrocarbons, as well as a vast number of organic compounds, from simple alcohols and aldehydes to complex proteins, carbohydrates and chlorophyll.

The first hydrogen-filled balloon was invented by Jacques Charles in 1783. German Count Ferdinand von Zeppelin later promoted the idea of rigid airships lifted by hydrogen that later were called Zeppelins, the first of which had its maiden flight in 1900. Hydrogen-lifted airships were used as observation platforms and bombers during World War I. Hydrogen has also been used as a coolant in turbogenerators, as battery fuel, for the processing of fossil fuels, in the production of ammonia, as a hydrogenating agent for food oils, and in the production of methanol. It is used as a tracer gas for minute leak detection, as a shielding gas in welding, and in cryogenic research.

Hydrogen
1
1.0079

He

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

3He

>99,80

0,000134

g

The first evidence of helium was observed in 1868 by French astronomer Pierre-Jules-César Janssen during a solar eclipse in Guntur, India. Later that same year, English astronomer Norman Lockyer observed a similar yellow line in the solar spectrum. He and English chemist Edward Frankland named the element with the Greek word helios, meaning “sun.” Scottish chemist Sir William Ramsay isolated helium on Earth for the first time in 1895.

Helium is the second most abundant element in the universe (although rare on Earth). It is a colorless, odorless gas that is also very slightly soluble in water, while being insoluble in ethanol. It is inert and monatomic in all standard conditions. Due to the small size of helium atoms, its diffusion rate through solids is three times that of air and around 65% that of hydrogen. No stable chemical compounds are known.

Helium-3 is used as a circulating medium in laboratory refrigerators to maintain constant temperatures below 3º K. Helium is also used as a lifting gas in buoyant airships and in most types of balloons, such as weather-, toy-, kite-type- and advertising balloons. Its lifting power is just slightly less than that of hydrogen. In nuclear physics, helium ions or alpha particles serve as projectiles in bombarding heavy nuclei to produce energy or to obtain artificial radioisotopes. Liquid helium is used in magnetic resonance imaging (MRI) equipment for the diagnosis of cancer and other soft-tissue diseases.

Helium
2
4.0026

Li

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

7Li

>99,95

92,50

Hydroxide Monohydrate

(7LiOH-H2O)

Lithium, named for the Greek word lithos (“stone”), was discovered in 1817 by Johan August Arfvedson during an analysis of petalite ore from the Swedish island of Utö.

A soft, silvery-white metal with a body-centered cubic structure, lithium has a heat capacity about the same as that of water. It ignites in air near its melting point and burns with a crimson-red flame and dense white fumes. It has a dangerous fire and explosion risk when exposed to water, nitrogen, acids or oxidizing agents. It is soluble in liquid ammonia, forming a blue solution.

Lithium has a high electrical conductivity and is used to make high-energy lithium batteries. It can be combined with lead, magnesium, aluminum or other metals for very useful alloys. Its most important application is in preparative chemistry as the starting material to prepare lithium hydride, amide, nitride, alkyls and aryls.

Lithium
3
6.941

C

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

13C

>99,00

1,10

g

Carbon has been known since ancient times, most familiarly as coal, charcoal, soot and diamond. Its name derives from the Latin word carbo, meaning “coal” or “charcoal.” Carbon exists in three allotropic forms — diamond, graphite and fullerenes — each differing distinctly from others in physical and chemical properties:

Diamond [7782-40-0] is one of the hardest substances known. It has a Mohs hardness of 10.0, a density of 3.513 g/cm3, and a melting point of about 3700 ºC. Carbon atoms in diamonds are arranged in cubic form, having stacking layers perpendicular to the diagonals of the cube. The diamond also occurs in hexagonal form, which is less stable than the cubic form. The hexagonal form of diamond is found in meteorites and can be synthesized.

Graphite [7440-44-0] is a black hexagonal crystal. The hexagonal layer has each carbon atom surrounded by three other carbon atoms. The C-C bond length is 1.415 Å. Because of the very weak van der Waal forces between the hexagonal layers, graphite is one of the softest solids, with a high lubricity and a density of 2.25 g/cm3. Graphite exhibits two manifestations: the stable hexagonal form that commonly occurs at ambient conditions, and a less stable rhombohedral form. Graphite can be converted to diamond under high temperatures (about 1400 ºC) and very high pressures (in the range of 4000-5000 atm) in the presence of a metal catalyst such as iron or nickel.

Fullerenes [99685-96-8] are polyhedral carbon allotropes consisting of large carbon molecules containing 60 to 120 carbon atoms. They are found in soot, charcoal, carbon black and many other carbonaceous matters, and they have high electrical conductivity and chemical reactivity.

Carbon is also produced and used in other forms — such as activated carbon, carbon black and coke — that have many commercial applications: purification of water and air, air analysis, waste treatment, removal of sulfur dioxide from stack gases, and decolorization of sugar. Carbon black includes several forms of artificially prepared carbon, such as furnace black, channel black, lamp black and animal charcoal, and is commonly used in typewriter ribbons, printing inks, carbon paper, paint pigments, and as an absorber for solar energy and UV radiation. Elemental carbon has many important applications: the diamond is a precious gem; graphite is used as an electrode and has numerous other applications; the isotope Carbon-14 is used in carbon dating; and the isotope Carbon-13 is used in tracer studies and nuclear magnetic resonance (NMR) imaging. 

Carbon
6
12.0107

N

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

15N

>96,00

0,37

KNO3

Nitrogen was discovered in 1772 by Daniel Rutherford. Its name derives from the Greek words nitron + genes, meaning “nitre” and “forming,” and the Latin word nitrum.

Nitrogen is a colorless, odorless, tasteless gas that is diamagnetic and converts to a colorless liquid. At ordinary temperatures, nitrogen is very stable and chemically inert to most substances, although elevated temperatures and pressures change this. Carbides of certain metals, such as cerium and uranium, react with nitrogen at very high temperatures, forming their nitrides; nitrogen also combines with alkali and alkaline earth elements at ordinary temperatures to form their nitrides. At low pressure and under electric discharge conditions, nitrogen produces a greenish-yellow glow, which continues to glow after the discharge. Such active nitrogen readily reacts with many unreactive elements in cold, such as mercury and sulfur, forming their nitrides.

Gaseous nitrogen has numerous uses in chemical, food, metal and electrical industries. It is needed in commercial production of ammonia and in the preparation of many nitrides. It is also the starting material in making cyanamide salts, cyanides and nitrogen oxides for producing nitric acid.

Nitrogen is also used in gas chromatography, as a carrier gas, to provide an inert atmosphere in chemical reactions, to prevent oxidation reactions, to reduce fire or explosion hazards, and to dilute a reacting gas. In the food industry it is used to prevent mold growth, spoilage from oxidation, and insect infestation. Other applications of nitrogen gas include pressurizing cable jackets, preventing carburization in welding and soldering, inflating balloons, agitating liquid baths, and cooling catalytic reactors in petroleum refining. Liquid nitrogen is used in the rapid freezing of food as well as its packaging, storage and transportation; for preserving blood, tissue and bone marrow; for cryopulverizing plastics, resins, waxes, spices and scrap rubber to achieve small particle size; and for deforming stainless steel to make high-strength wires for springs.

Nitrogen
7
14.0067

O

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

16O

99,99

99,76

Water

180

>97,00

0,20

Water

Oxygen was discovered in 1774 by Joseph Priestley and Carl Scheele, in England and Sweden. Its name derives from the Greek words oxy + genes, meaning “acid” (sharp) and “forming” or “creating.”

A colorless, odorless, tasteless gas, oxygen reacts with practically all elements, a number of inorganic salts, and all organics under various conditions, including elevated temperatures and pressures, or in the presence of a catalyst. While certain metals — such as sodium, potassium and calcium — react vigorously with oxygen at ordinary temperatures, most other metals react with it at elevated temperatures. All nonmetals except helium, neon and argon react with oxygen.

Oxygen is present in all living organisms and is vital for sustenance of life in the animal kingdom. Most oxygen manufactured today is consumed in refining iron in the steel industry, removing carbon, silicon, sulfur, phosphorous, manganese and other impurities from liquid iron. It is also used to enrich fuel-air flame in furnaces that produce copper and nickel from sulfide ores.

Oxygen also has major uses in the chemical industry for the oxidation of methane, ethylene and other hydrocarbons. It is used in all breathing masks and life-support devices, respirators and incubators. In medicine it is also administered under hypoxia, respiratory distress, impaired respiratory function, asthma attacks, for the treatment of cyanosis, and for poisoning by carbon monoxide and other toxicants. Oxygen is also used in fermentation, the bleaching of wood chips, odor control, and as a flame-enhancing agent.

Oxygen
8
15.9994

Ne

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

20Ne

99,995

90,48

g

21Ne

70,27-90,00

0,27

g

22Ne

99,95

9,25

g

Neon was discovered in 1898 by Sir William Ramsay and Morris Travers, who also discovered krypton and xenon. It takes its name from the Greek word neon, which means “new.”

Neon is a colorless, odorless, tasteless gas that forms face-centered cubic crystals. A zero-valent element with a highly stable octet configuration, neon is inert to practically all chemicals. Neon forms an unstable hydrate at low temperatures under high pressure. It does ionize, however, under high vacuum (as in an electric discharge tube), forming some ions.

The most important use of this gas is in “neon” lights and fluorescent signs for advertisements. Contained in glow discharge lamps or high-voltage discharge tubes at low pressure, neon emits red light. In the presence of mercury vapors, the color of the glow turns blue. Neon is also used in scintillation counters, neutron fission counters, proportional counters, and ionization chambers for detection of charged particles.

Neon
10
20.179

Mg

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

24Mg

99,75

78.99

o

25Mg

>98,30

10,00

o

26Mg

99,60

11,01

o

Magnesium was discovered in 1755 by Sir Humphrey Davy. It is a silvery-white metal, and its name originates from the Greek name Magnesia, a district of Thessaly. It occurs in all plants — its porphyrin complex, chlorophyll, is essential for photosynthesis. It is also an essential nutrient for humans.

Magnesium is a silvery-white metal with a close-packed hexagonal structure. It is soluble in dilute acids and reacts very slowly with water at ordinary temperatures. At room temperature, it is not attacked by air; however, when heated, it burns with a dazzling white light, forming the oxide MgO and the nitride Mg3N2. At ordinary temperatures, magnesium is stable in alkalis, both dilute and concentrated; however, hot solutions of alkalis above 60 ºC attack the metal. Magnesium combines with halogens at elevated temperatures, forming halides; and with nitrogen phosphorus, sulfur, and selenium at elevated temperatures, forming their binary compounds. Magnesium exhibits single displacement reactions, replacing lower metals in electrochemical series from their salt solutions or melt. Magnesium also reduces nonmetallic oxides such as carbon dioxide, carbon monoxide, sulfur dioxide and nitrous oxide, burning at elevated temperatures.

According to anecdotal history, in 1618, a farmer in Epsom, England, attempted to feed his cows water from a well. Because of the bitter taste, the cows refused to drink the water. However, the farmer noticed that the same water seemed to heal scratches and rashes. The fame of “Epsom salts” spread. Eventually the compound magnesium sulfate was recognized.

Today, magnesium has many practical applications, including the production of titanium by Kroll process and obtaining uranium from its fluoride. Magnesium alloys with aluminum, zinc, copper, nickel, lead, zirconium and other metals are used in automobile parts, aircraft, missiles, space vehicles, ship hulls, underground pipelines, memory discs, machine tools, furniture, lawnmowers, ladders, toys and sporting goods. It is also used in making small and lightweight dry cell batteries. Chemical applications of magnesium include its uses as a reducing agent, in preparation of Grignard reagents for organic syntheses, and for purifying gases. Magnesium is also used in blasting compositions, explosive sensitizers, incendiaries, signal flares and pyrotechnics.

Magnesium
12
24.305

Si

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

28Si

99,90-99,99

92.23

e,o

29Si

>99,00

4.67

e,o

30Si

>99,00

3,10

e,o

The discovery of silicon is credited to Jöns Jakob Berzelius in 1824 in Sweden. Its name originates with the Latin word silicis, which means “flint.” It is the second most abundant element, exceeded only by oxygen, making up 25.7% of the earth's crust by weight.

Silicon exists in two allotropic modifications: Crystalline silicon is made up of grayish-black lustrous needle-like crystals, or octahedral platelets, with a cubic structure. Amorphous silicon is a brown powder. Elemental silicon is relatively stable in most substances at ordinary temperatures and shows similarity to other elements of its group. It exists as sand quartz, flint, amethyst, agate, opal, jasper and rock crystal.

Silicates and silica have many applications in numerous fields, including making cements and concretes for building materials, glasses and glassware, ceramics, pigments, adsorbents, paper boards, fillers, detergents, precious gems, catalysts and water-softeners. Silicones are used as lubricants and in making rubbers, plastics, electrical coatings, adhesives, paints and varnishes, and as water repellents for textiles, papers and concrete. Elemental silicon has a major application in computer chips. Silicon of hyperpurity — doped with trace elements, such as boron, phosphorus, arsenic or gallium — is one of the best semiconductors and is used in transistors, power rectifiers, diodes and solar cells. Hydrogenated amorphous silicon converts solar energy into electricity.

Inhalation of silica dusts or silicate mineral dusts can cause silicosis or other lung diseases.

Silicon
14
28.08555

S

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

32S

99,90

94,99

e

33S

>99,30

0.75

e

34S

>99,00

4,25

e

36S

55,00-99,20

0,01

e

Sulfur (also known as “sulphur”) has been known since ancient times and was referred to in Genesis as “brimstone.” Assyrian texts dated around 700-600 BC refer to it as the “product of the riverside,” where deposits could be found. Its name has origins in the Sanskrit word sulvere and the Latin word sulphurum, both meaning “sulfur.”

Sulfur, particularly in its S8 form, is insoluble in water but dissolves in carbon disulfide, anhydrous liquid ammonia and methylene iodide. It is moderately soluble in benzene, toluene, chloroform and acetone, its solubility increasing with temperature. Solid polymeric sulfur is practically insoluble in all solvents.

Sulfur exists in several allotropic forms: at ordinary temperatures it exists as thermodynamically stable alpha-cyclooctasulfur, which has two other modifications, the beta and the gamma forms. Alpha-cyclooctasulfur, or the alpha sulfur, is a yellow orthorhombic crystalline solid. It has a density of 2.07 g/cm3 at 20 ºC and is stable at ordinary temperatures. Beta-sulfur has pale yellow, opaque, needle-like crystals with a monoclinic structure that is brittle. It is stable between 94.5 ºC and 120 ºC and converts to an orthorhombic form on standing. It has a density of 1.96 g/cm3 and melts at 115.2 ºC. Gamma-sulfur, a pale yellow amorphous solid, is a second monoclinic form of cyclooctasulfur. It has a density of 1.92 g/cm3 and melts at 120 ºC. There are also various other forms of sulfur including cyclohexa-, cyclohepta-, cyclonona-, cyclodeca- and cyclododeca-sulfur.

Elemental sulfur is used for vulcanizing rubber, in making black gunpowder, as a soil conditioner, as a fungicide, preparing a number of metal sulfides, and producing carbon disulfide. It is also used in matches; for bleaching wood pulp, straw, silk and wool; and in the synthesis of many dyes. Pharmaceutical-grade precipitated and sublimed sulfurs are used as scabides and as antiseptics in lotions and ointments.

Sulfur
16
32.06

Cl

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

35Cl

99,00

75,78

sc

37Cl

98,00

24,22

sc

Chlorine was discovered in 1774 by Carl William Scheele. Its name derives from the Greek word chloros, meaning “pale green.”

A greenish-yellow gas with a suffocating odor, chlorine combines directly with nearly all other elements. It becomes a pale yellow crystal at -101.5 ºC. It is moderately soluble in water. It has known oxidation states from -1 to +7. It occurs as a diatomic molecule Cl2, containing a single covalent bond in which the Cl-Cl bond distance is 1.99 Å.

Chlorine is a respiratory irritant that was used in war as early as 1915. Today, much of the chlorine supply is used in the manufacture of chlorinated cleaning compounds, pulp bleaching, disinfectants and textile processing. It is also used in the production of safe drinking water all over the world.

Chlorine
17
35.452

Ar

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

36Ar

>99,00

0,34

g

38Ar

99,90

0,06

g

40Ar

99,99

99,60

g

Argon was discovered in 1894 by Sir William Ramsay and Lord Rayleigh. Its name derives from the Greek word argos, meaning “inactive.”  Today the chemical symbol for argon is Ar, but until 1957 its symbol was simply A.

Argon is a colorless and odorless gas that is present to a very small extent in the atmosphere. It is practically insoluble in water. It makes a good atmosphere for working with air-sensitive materials since it is heavier than air and less reactive than N2. It is an effective “blanket” for the production of titanium and other reactive elements. It provides a protective atmosphere for growing silicon and germanium crystals.

Commercial applications for argon include its use in electric light bulbs and in fluorescent tubes. It is also used as an inert gas shield for arc welding and cutting. 

Argon
18
39.948

K

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

39K

>99,90

93,26

Cl

40K

5,00-16,95

0,01

Cl

41K

>95,00

6,73

Cl

Potassium was discovered in 1807 by Sir Humphry Davy. Its name originates with the English word potash (“pot ashes”) and the Arabic word qali (meaning “alkali”). The origin of the symbol K is the Latin word kalium (“alkali”). It is an essential element needed for plant growth. Potassium deficiency has also been associated with several common animal ailments. It is in extracellular fluid in animals, at lower concentrations than sodium.

A soft, silvery metal, potassium has a body-centered cubic structure that rapidly oxidizes in moist air and imparts a crimson-red color to flame. It is soluble in liquid ammonia, aniline, mercury and sodium. It reacts violently with water and acids, reacts with alcohol, and dissolves in liquid ammonia and mercury. It also reacts with oxygen or air, forming three oxides: potassium monoxide, potassium peroxide and potassium superoxide. Reactions with halogens, fluorine, chlorine and bromine occur with explosive violence. Violent reactions can occur with many metal halides as well.

Potassium products have applications in the bleaching of textiles and straw, in the tanning of leather, and in the food industry — potassium is the main ingredient in baking powder, it improves dough strength and rise height, and it is a preservative in wine- and beer-making. Potassium nitrate is used in gunpowder and fertilizer. Potassium cyanide is used in mining and in organic synthesis. Potassium carbonate is used in glass, soap, fluorescent lamps, textile dyes and pigments. Potassium chlorate is added to matches and explosives. Potassium bromide was formerly used as a sedative and in photography.

Potassium
19
39.098

Ca

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

40Ca

99,99

96,941

c

42Ca

68,00-96,80

0,647

c

43Ca

62,20-90,00

0.135

c

44Ca

94,50-99,00

2.086

c

46Ca

15,90-24,80

0.004

c

48Ca

64,00-97,10

0.187

c

Calcium was discovered in 1808 by Sir Humphry Davy. Its name originates with the Latin word calx, meaning “lime.” Compounds such as lime were prepared by the Romans as early as the first century AD. Literature dating back to 975 AD indicates that plaster of Paris (calcium sulfate) is useful for setting bones. 

A moderately soft, bright silver-white, crystalline metal, calcium oxidizes in air to form adherent protective film and can be machined, extruded or drawn. It is soluble in acid and decomposes in water to liberate hydrogen. It has a brick-red color when introduced to flame in a flame test.

Calcium metal reacts with a number of nonmetallic elements, forming their corresponding binary compounds. While the reaction with fluorine occurs at ambient temperatures, other elements combine only at elevated temperatures in the range of 300-900 ºC. Calcium reacts vigorously with water at ordinary temperatures, with the evolution of hydrogen. It is a strong reducing agent and can reduce most metal oxides and halides into their metals at elevated temperatures. It can reduce all the lower electropositive metals.

Applications of calcium are mostly found in metallurgy. It is used to produce alloys with aluminum, lead, beryllium, copper, silicon and other metals. It is also used as a desulfurizer, decarburizer and deoxidizer for ferrous and nonferrous alloys; for removal of bismuth from lead; and as a reducing agent for zirconium, uranium, thorium and other metals.

Calcium
20
40.078

Ti

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

46Ti

97,00

8,25

o

47Ti

>95,00

7,44

o

48Ti

>96,00

73,72

o

49Ti

47,60-92,40

5,41

o

50Ti

83,00

5,18

o

Titanium was discovered in 1791 by Reverend William Gregor, who recognized the presence of the element in menachanite, a mineral named after Menaccan in Cornwall, England. It takes its name from the Titans, the sons of Gaia, the Earth goddess in Greek mythology.

A white lustrous metal that is ductile when free of oxygen, titanium is also a low-density, high-strength metal, as strong as steel but 45% lighter. It has two allotropic modifications. The alpha form has a close-packed hexagonal crystal structure, a density of 4.54 g/cm3 at 20 ºC, and is stable up to 882 ºC. It converts very slowly to a body-centered cubic beta form at 882 ºC. The density of the beta form is 4.40 g/cm3 at an estimated 900 ºC.

Titanium metal is highly resistant to corrosion. It is unaffected by atmospheric air, moisture and sea water, allowing many of its industrial applications. The metal burns incandescently in air at about 1200 ºC, forming titanium dioxide. The metal also burns on contact with liquid oxygen. It combines with nitrogen at about 800 ºC, forming the nitride and producing heat and light. Titanium reacts with all halogens at high temperatures. It is soluble in hot concentrated sulfuric acid, forming sulfate, and with hydrofluoric acid, forming fluoride.

Elemental titanium is found in plants, animals, eggs and milk. Its alloys have wide industrial applications, as they possess high tensile strength, are lightweight, and can withstand extreme temperatures. They are often used in the construction of aircraft and missiles. They have also been used in medical prostheses, orthopedic and dental implants, dental and endodontic instruments and files, dental implants, jewelry and mobile phones.

Titanium
22
47.867

V

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

50V

>55,00

0.25

o

51V

>99,90

99.75

o

Vanadium was discovered in 1801 by both Manuel del Rio and Nils Sefström. Its name derives from that of Vanadis, the goddess of beauty in Scandinavian mythology.

Vanadium is a bright, silvery-white, ductile solid with a body-centered cubic structure. It is insoluble in water, dilute sulfuric acid, hydrochloric acid and alkalis. It is resistant to corrosion but soluble in nitric, hydrofluoric and concentrated sulfuric acids. It can act as either a metal or a nonmetal, and it forms a variety of complex compounds and is nontoxic as metal.

Vanadium forms four oxides: the light gray monoxide, VO; the blue-black dioxide, VO2; the black sesquioxide, V2O3; and the orange-red pentoxide, V2O5. It also combines with chlorine on heating, producing three known chlorides: the green dichloride, VCl2; the pink trichloride, VCl3; and the brown-red tetrachloride, VCl4.

Among its industrial applications, vanadium is added to steel for high resistance to oxidation and to stabilize carbide. The foil is used for cladding titanium to steel, and a vanadium-gallium alloy is used in making superconductive magnets.

Vanadium
23
50.9415

Cr

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

50Cr

96,50-99,70

4.35

o,m

52Cr

>98,80

83.79

o,m

53Cr

92,80-97,70

9.50

o,m

54Cr

>99,00

2.36

o,m

Chromium was discovered in 1797 by Louis-Nicolas Vauquelin. It is named for the Greek word chroma, meaning "color," signifying the strong and varied colors of chromium compounds.

Chromium is a hard, brittle, blue-white metal, with a body-centered cubic crystal. It exists in active and passive forms, the latter giving rise to its corrosion resistance due to a thin surface oxide layer that passivates the metal when treated with oxidizing agents. The active form reacts readily with dilute acids to form chromous salts. It is soluble in acids (except nitric) and strong alkalis, but insoluble in water.

Elemental chromium reacts with anhydrous halogens, hydrogen fluoride and hydrogen chloride, forming the corresponding chromium halides. At elevated temperatures in the 600-700 ºC range, chromium reacts with hydrogen sulfide or sulfur vapor, forming chromium sulfides. Chromium metal reacts at 600-700 ºC with sulfur dioxide and caustic alkalis. It combines with phosphorus at 800 ºC. Reaction with ammonia at 850 ºC produces chromium nitride. Reaction with nitric oxide forms chromium nitride and chromium oxide.

The most important application of chromium is in the production of steel. High-carbon and other grades of ferro-chromium alloys are added to steel to improve mechanical properties, to increase hardening and to enhance corrosion resistance. Chromium also is added to cobalt- and nickel-based alloys for the same purposes. Refractory bricks are composed of chromium oxides and are used in roofs of open hearths, sidewalls of electric furnaces, vacuum apparati and copper converters. Chromium coatings are applied on the surfaces of other metals — for decorative purposes, to enhance resistance, and to lower the coefficient of friction. Radioactive Chromium-51 is used as a tracer in the diagnosis of blood volume.

Hexavalent chromium compounds have an irritating and corrosive effect on human tissue, resulting in ulcers and dermatitis on prolonged contact. Tolerance for chromium dust and fume is 0.5 mg/m3 of air. Hexavalent chromium is also a known carcinogen and is moderately toxic and corrosive to skin. Inhalation of Cr6+ dust or mist can cause perforation of the nasal septum, lung irritation, and congestion of the respiratory tract.

Chromium
24
51.996

Fe

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

54Fe

>94,30

5.85

m,o

56Fe

>99,70

91,75

m,o

57Fe

>95,00

2,12

m,o

58Fe

92,80-99,80

0,28

m,o

Iron has been known since prehistoric times. Genesis says that Tubal-Cain, seven generations from Adam, was “an instructor of every artificer in brass and iron.” Smelted iron artifacts have been identified from as early as 3000 BC. The name “iron” derives from the Anglo-Saxon word iron or iren, and the symbol Fe comes from the Latin word ferrum, meaning “iron.”

Iron is a soft, white, ductile metal, and the fourth most abundant element in earth’s crust. It is the only metal that can be tempered. Its body-centered cubic form is stable to 910 ºC; from 910 ºC to 1390 ºC it has a face-centered cubic form; and above 1390 ºC it returns to the body-centered form. Its mechanical properties are altered by impurities, especially carbon. Iron is highly reactive chemically, it is a strong reducing agent, and it oxidizes readily in moist air and reacts with steam when hot to yield hydrogen and iron oxides. It is attracted by magnets and rapidly loses its magnetism. It is ferromagnetic at ordinary temperatures but becomes paramagnetic when heated to its Curie point of 768 ºC.

Iron exhibits single-replacement reactions, precipitating less electropositive metals out of their salt solutions. Thus, solid iron can reduce many metals, such as copper, silver, gold, mercury, tin and nickel. Solid iron undergoes rusting by reacting with oxygen in the presence of water; in moist air, it rapidly converts to rust.

Iron occurs in every mammalian cell and is vital for life processes. It is bound to various proteins and is found in blood tissues. Industrial uses of iron as carbon steels are numerous, surpassing the uses of any other alloys (carbon steels are alloys of iron containing carbon in varying proportions). Non-steel iron alloys such as cast iron, wrought iron, nickel iron and silicon iron have many important applications as well. Another important application of iron is as an industrial catalyst: it is used in catalyst compositions in the Haber process for synthesis of ammonia, and in the Fischer-Tropsch process for producing synthetic gasoline.

Iron
26
55.847

Ni

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

58Ni

>99,48

68,08

m,o

60Ni

89,00-99,60

26,22

m,o

61Ni

91,00->99,00

1,14

m,o

62Ni

98,00-99,28

3,63

m,o

64Ni

91,00-99,32

0,93

m,o

Nickel was discovered in 1751 by Axel Fredrik Cronstedt. Its name is derived from the German word kupfernickel, meaning “Devil's copper,” “false copper” or “St. Nicholas's copper.”

A malleable, silvery-white lustrous metal, nickel has a face-centered cubic crystal structure. It is also ductile, ferromagnetic, and readily fabricated by hot and cold working. It takes high polish and demonstrates an excellent resistance to corrosion. It is insoluble in water as well as in ammonia solution, it is slightly soluble in dilute hydrochloric acid, and it is slightly soluble in dilute hydrochloric and sulfuric acids.

Because of nickel's slow rate of oxidation at room temperature, it is considered corrosion-resistant. Historically, this has led to its use for plating metals such as iron and brass; in chemical apparatus; and in certain alloys that retain a high silvery polish, such as German silver. About six percent of world nickel production is still used for corrosion-resistant pure-nickel plating. Nickel was once a common component of coins, but it has largely been replaced by cheaper iron for this purpose, especially since the metal is a skin allergen for some people. It was reintroduced into United Kingdom coinage in 2012 despite objections from dermatologists. Nickel is preeminently an alloy metal, and its chief use is in nickel steels and nickel cast irons, of which there are many varieties. It is also widely used in many other alloys, such as nickel brasses and bronzes, and alloys with copper, chromium, aluminium, lead, cobalt, silver and gold.

As a compound, nickel has a number of niche chemical manufacturing uses, such as a catalyst for hydrogenation. Enzymes of some microorganisms and plants contain nickel as an active site, which makes the metal an essential nutrient for them.

Nickel can be flammable and toxic as either a dust or a fume. It is also classified as a carcinogen.

Nickel
28
58.693

Cu

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

63Cu

99,90

69.17

m,o

65Cu

89,70-99,00

30.83

m,o

The discovery of copper dates from prehistoric times. There are reports that copper beads dating to 9000 BC have been found in Iraq. Techniques for refining copper from its ore were discovered around 5000 BC, and it was used in pottery in North Africa by about 4000 BC. Part of the reason it was used so early is that it is relatively easy to shape, although it is too soft to be used in most tools. When combined with other metals, the resulting alloys are harder than copper — bronze, for example, which is a mixture of copper and tin. Copper's name derives from the Latin word cuprum, meaning “the island of Cyprus.”

Copper has a distinctive reddish-brown color; it is ductile and has a face-centered cubic crystal. It is insoluble in water and dissolves in nitric acid and hot sulfuric acid; it is slightly soluble in hydrochloric acid and is soluble in ammonium hydroxide, ammonium carbonate and potassium cyanide solutions. Copper is more resistant to atmospheric corrosion than iron, forming a green layer of hydrated basic carbonate. Heating the metal in dry air or oxygen yields black copper(II) oxide, which on further heating at high temperatures converts to the red cuprous form. Copper(II) ion readily forms complexes with various ligands. It forms a deep blue solution in aqueous ammonia.

The metal, its compounds and its alloys have numerous applications in every sphere of life, making it one of the most important metals. Almost all coinages in the world are made out of copper or its alloys. The metal is an excellent conductor of electricity and heat and is used in electric wiring, switches and electrodes. Other applications include plumbing, piping, roofing, cooking utensils, construction materials and electroplated protective coatings. 

Although the toxicity of metallic copper is very low, many copper(II) salts may have varying degrees of toxicity. Inhalation of dusts, mists or fumes of the metal can cause nasal perforation, cough, dry throat, muscle ache, chills and “metal fever.” Copper in trace amounts is a nutritional requirement, however, used metabolically in plant and animal enzymes and other biological molecules.

Copper
29
63.546

Zn

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

64Zn

>99,00

48,60

m,o

66Zn

>98,70

27,90

m,o

67Zn

>89,60

4,10

m,o

68Zn

>98,20

18,80

m,o

70Zn

>95,00

0,60

m,o

Zinc was discovered in 1746 by Andreas Marggraf. It takes its name from the German word zink, which means “point” or “tine” and refers to the form of zinc crystals after smelting. Centuries earlier, zinc ores were used for making brass, a mixture of copper and zinc.

Zinc is a shining white metal with a bluish-gray luster (called “spelter”) and is brittle at room temperature. It is a good conductor of electricity, is soluble in acids and alkalis, and is insoluble in water. It has a hexagonal close-packed structure. Zinc is diamagnetic and is also highly electropositive, replacing less electropositive metals from their aqueous salt solutions or melts. The metal is attacked by mineral acids. Reactions with sulfuric and hydrochloric acids produce hydrogen.

Zinc is attacked by moist air at room temperatures. Dry air causes no reaction at ambient temperatures, but the metal combines with dry oxygen rapidly above 225 ºC. Zinc reacts with carbon dioxide in the presence of moisture at ordinary temperatures, forming a hydrated basic carbonate. The metal, on heating with dry halogen gases, yields zinc halides. However, in the presence of moisture, the reaction occurs rapidly at ambient temperatures. The metal dissolves in hot solutions of caustic alkalis to form zincates and to evolve hydrogen.

Some important applications of zinc include galvanizing steel, producing die castings, as a chemical addition to rubber and paints, in dry cells, in making electrodes, and as a reducing agent. Zinc forms numerous alloys, including brass, nickel silver, German silver, commercial bronze, soft solder, aluminum solder and spring brass. Zinc is also an essential nutrient element required for the growth of animals.

As an essential nutrient, zinc is not regarded as toxic. However, the metal fumes, oxide fumes and chloride fumes can produce adverse inhalation effects. Ingestion of soluble salts can cause nausea.

Zinc
30
65.39

Ga

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

69Ga

>99,60

60,11

m/o

71Ga

>99,80

39,89

m/o

Gallium was discovered in 1875 by Paul-Émile Lecoq de Boisbaudran. It takes its name from the Latin name Gallia, meaning “France,” and possibly also from the Latin word gallus, meaning “rooster” (and reflecting the discoverer's name, Lecoq). It is one of the few metals that can be liquid near room temperature, which makes it an effective feature of high-temperature thermometers. It can be undercooled to almost 0 °C without solidifying, and it is actually more dense as a liquid than as a solid.

Gallium is a gray orthogonal crystal or silvery liquid. It is soluble in acids and alkalis and slightly soluble in mercury. It reacts with most metals at high temperatures. Gallium exists in a liquid state in the widest temperature range. Chemical properties of gallium fall between those of aluminum and indium. It combines with phosphorus, arsenic and antimony, forming the corresponding binary compounds, which exhibit interesting semiconductor properties.

The most important use of gallium is as a doping agent for semiconductors, transistors and other solid-state devices. Some gallium compounds also have major applications in electroluminescent light emission, microwave generation and UV-activated powder phosphors. Another important use of gallium, in oxide form, involves the spectroscopic analysis of uranium oxide. Gallium also is used to make many low-melting alloys. Some other uses for gallium are in high-vacuum systems as a liquid sealant, as a heat-transfer medium, and to produce mirrors on glass surfaces.

Gallium
31
69.723

Ge

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

70Ge

95,30-97,60

20,38

m,o

72Ge

96,40-98,20

27,31

m,o

73Ge

>95,50

7,76

m,o

74Ge

>95,20

36,72

m,o

76Ge

>88,00

7,83

m,o

Germanium was discovered in 1886 by Clemens Winkler. Its name originates with the Latin name Germania, meaning “Germany.”

Germanium is a grayish-white cubic crystal. Elemental germanium can be prepared in extreme purification with a high degree of crystalline perfection, so as to yield highly-characterized surfaces. Its conductivity depends largely on added impurities. It is attacked by nitric acid and aqua regia, but it is stable in water, acids and alkalis in the absence of dissolved oxygen. It is insoluble in water, dilute acids and dilute alkalis. The chemical properties of germanium fall between those of silicon and tin. It forms both the divalent and tetravalent compounds, the oxidation state +4 being more stable than the oxidation state +2. The metal is stable in air and water at ambient temperatures. However, it reacts with oxygen at elevated temperatures, forming divalent and tetravalent oxides. 

The most important uses of germanium are in electronic industries. It is a semiconductor material exhibiting an exponential increase of conductivity with increasing temperature. Other applications include infrared detectors, microscopes, various optical instruments, as a phosphor in fluorescent lamps, as an alloying agent, and as a catalyst. The activity of some germanium compounds against certain bacteria makes them of interest as chemotherapeutic agents.

Germanium
32
72.61

Se

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

74Se

>99,90

0,89

e

76Se

>99,50

9,36

e

77Se

>99,20

7,63

e

78Se

>99,30

23,78

e

80Se

99,90

49,61

e

82Se

99,90

8,73

e

Selenium was discovered in 1817 by Jöns Jakob Berzelius. Its name originates with the Greek word selene, meaning “moon.” It exists in several allotropic forms:

Amorphous forms exhibit two colors, occurring as either a red powder with density 4.26 g/cm3 and hexagonal crystal structure, or a black vitreous solid with density 4.28 g/cm3. The red form converts to the black form on standing, and it melts at 60-80 ºC. It is insoluble in water but reacts with water at 50 ºC when freshly precipitated. It is soluble in sulfuric acid, benzene and carbon disulfide.

Crystalline selenium exhibits two monoclinic forms: an alpha form, constituting dark red transparent crystals with a density of 4.50 g/cm3, converts to a metastable beta form of hexagonal crystal structure when heated to about 170 ºC. They are both insoluble in water, soluble in sulfuric and nitric acids, and very slightly soluble in carbon disulfide. Also, both crystalline forms convert into gray metallic modification on heating.

The gray metallic form of selenium is its most stable modification. It constitutes lustrous-gray-to-black hexagonal crystals with a density of 4.18 g/cm3 at 20 ºC and a melting point of 217 ºC. It is soluble in sulfuric acid and chloroform, very slightly soluble in carbon disulfide, and insoluble in alcohol.

Electrically, selenium acts as a rectifier, and it has marked photoconductive and photovoltaic action (it converts radiant to electrical energy). Selenium forms binary alloys with silver, copper, zinc, lead and other elements. The chemical properties of selenium fall between those of sulfur and tellurium; thus, it reacts with oxygen similarly to sulfur, forming two oxides. The metal combines with halogens, forming their halides. It is not attacked by hydrochloric acid, nor does it react with dilute nitric and sulfuric acids.

Selenium has many industrial uses, particularly electronic and solid-state applications, which are attributed to its unique properties: it converts light directly to electricity; its electrical resistance decreases with increased illumination; and it is able to convert alternating current to direct current. It is used in photoelectric cells, in solar cells, and as a rectifier in radio and television sets. Although a toxic metal at moderate concentrations, selenium is also an essential nutritional element at trace concentrations. Some of its compounds, such as hydrogen selenide, are very toxic. Exposure to the metal fumes can cause severe irritation of eyes, nose and throat. The United States Environmental Protection Agency lists the metal as one of the priority pollutant metals in the environment.

Selenium
34
78.96

Br

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

79Br

>99,00

50,69

sb

81Br

>99,00

49,31

sb

Bromine was discovered in 1825 and 1826 by Carl Jacob Löwig and Antoine-Jérôme Balard, respectively. It takes its name from the Greek word bromos, meaning “stench.” Its compounds were used well before it was recognized as an element: a purple dye called “Tyrian purple” was produced using an organobromine compound as well as an excretion from a particular kind of mussel.

Bromine is a dark reddish-brown liquid with a strong, disagreeable odor. It is the only nonmetallic element that remains a liquid at ambient temperatures; it solidifies at -7.2 ºC. It is sparingly soluble in water and soluble in common organic solvents. Most reactions of bromine are similar to those of other halogens, and its reactivity falls between those of chlorine and iodine. It readily attacks a number of metals, including alkali and alkaline earth metals, palladium, platinum, aluminum, copper, antimony and tin, forming their bromides, which are normally vigorous-to-violent reactions. It oxidizes a number of substances, including metal carbides, carbonyls, hydrides and organic substances. Bromine combines with fluorine at room temperature, forming bromine trifluoride. The reaction produces a luminous flame.

Combination reactions occur with several nonmetals. In aqueous solutions, bromine hydrolyzes slightly, forming unstable hypobromous acid, which decomposes to hydrobromic acid and oxygen, causing the bleaching action of bromine water. The decomposition is accelerated by light.

Bromine is used in bleaching fibers and as a disinfectant for water purification. Other applications are in organic synthesis as an oxidizing or brominating agent; in the manufacture of ethylene dibromide, methyl bromide and other bromo compounds for dyes and pharmaceutical uses; as a fire retardant for plastic; and in chemical analysis. Ethylene dibromide is used in anti-knock fluids in motor fuels. Over 80% of the bromine produced is consumed in the manufacture of this compound.

Most reactions of bromine are highly exothermic, which can cause incandescence or a sudden increase in pressure and rupture of reaction flasks. Reactions of liquid bromine with most metals (or any metal in finely divided state, metal hydrides, carbonyls and nitrides) can be explosive. Many oxides and halides of nonmetals, such as nitrogen triiodide or phosphorus trioxide, react explosively or burst into flame in contact with liquid bromine. Bromine is moderately toxic by all routes of exposure. It is an irritant to the eye and respiratory tract. Inhalation can cause dizziness, headache, coughing and lacrimation. A short exposure (to 1000 ppm for 15 minutes) can be fatal to humans. Ingestion produces nausea, abdominal pain and diarrhea. The liquid is corrosive to skin.

Bromine
35
79.904

Kr

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

78Kr

99.90

0.35

g

80Kr

99.90

2.25

g

82Kr

99.90

11.60

g

83Kr

99.90

11.50

g

84Kr

99.90

57.00

g

86Kr

99.90

17.30

g

Krypton was discovered in 1898 by Sir William Ramsay and Morris W. Travers. Its name originates with the Greek word kryptos, meaning “hidden.”

Krypton is a colorless, odorless, tasteless gas. It liquefies at -153.22 ºC and solidifies at 157.36 ºC to a white crystalline substance with a face-centered cubic structure. It is slightly soluble in water.

Krypton is an inert gas element. Its closed-shell, stable octet electron configuration allows zero reactivity with practically any substance. Only a few types of compounds, complexes and clathrates have been synthesized, mostly with fluorine, the most electronegative element.

The commercial applications of krypton are fewer than those of helium or argon. Its principal use is in fluorescent lights. It is mixed with argon as a filling gas to enhance brightness in fluorescent tubes. Other applications are in flash tubes for high-speed photography and incandescent bulbs. Radioactive Krypton-85 is used as a tracer to monitor surface reactions. The unit of length “meter” was once defined in terms of the orange-red spectral line of Krypton-86.

Krypton
36
83.30

Rb

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

85Rb

>99,50

72,16

Cl

87Rb

>98,00

27,84

Cl,m

Rubidium was discovered spectroscopically in 1861 by Robert Bunsen and Gustav Kirchhoff. The origin of its name is the Latin word rubidius, meaning “dark red” or “deepest red,” referring to the element's bright red spectroscopic lines.

A soft, silvery white solid with body-centered cubic crystals, rubidium is ductile and very light and is easily oxidized in air. The liquid metal vaporizes, producing a blue vapor. It is soluble in acids and alcohol and reacts violently in water to form rubidium hydroxide. Rubidium is also a highly reactive metal, with most of its reactions similar to those of sodium or potassium. The metal ignites spontaneously in air, forming oxides, and is coated rapidly with a gray-blue oxide film. The reaction with dilute mineral acids can proceed with explosive violence, releasing hydrogen. Rubidium combines with hydrogen and nitrogen, forming the hydride RbH and the nitride Rb3N.

Rubidium metal and its salts have very few commercial applications. They are used in research involving magnetohydrodynamics and thermoionic experiments. Rubidium is also used in photocells and in vacuum tubes. The beta-emitter isotope Rubidium-87 is used to determine the age of some rocks and minerals. Radioisotopes of rubidium have been used as radioactive tracers to trace the flow of blood in the body. Rubidium salts are used in pharmaceuticals as soporifics and sedatives, as well as for treating epilepsy.

Rubidium
37
85.468

Sr

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

84Sr

68,70-82,00

0.56

c

86Sr

>96,00

9.86

c

87Sr

>90,00

7,00

c

88Sr

99,90

87.58

c

Strontium was discovered in 1790 by Adair Crawford. It is named for the village of Strontian in Scotland.

A silvery-white metal when freshly cut, strontium rapidly turns yellow on exposure to air, forming a thin oxide coating. It has a face-centered cubic structure, is malleable and ductile, and is chemically similar to calcium. Soluble in alcohol, ethanol and acids, strontium is also a reactive metal. In its finely-divided form, the metal is pyrophoric and ignites in air to form both the oxide SrO and the peroxide SrO2. Similarly, when heated with chlorine gas or bromine vapor, it burns brightly, forming its halides, SrCl2 or SrBr2. Strontium is also a reducing agent: it reduces oxides and halides of metals at elevated temperatures to metallic form. Spontaneously flammable in powder form, strontium ignites when heated above its melting point. It reacts with water to evolve hydrogen.

Elemental strontium has only minor uses, since most applications also involve calcium and barium. Strontium's alloys are used as “getters” for vacuum tubes. Strontium is incorporated in glass to make picture tubes for color televisions. Its compounds are used in tracer bullets and in fireworks to produce red signal flares. Strontium titanate is a gemstone. The radioactive isotope Strontium-90, with a half-life of 29 years, is a product of nuclear fission and a high-energy beta emitter. This isotope is a lightweight nuclear-electric power source in space vehicles and remote weather stations.

Strontium's radioactive isotopes emit high-energy beta radiation, causing damage — including cancer — to bone marrow and blood-forming organs.

Strontium
38
87.62

Zr

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

90Zr

>96,80

51,45

o,m

91Zr

>88,00

11,22

o,m

92Zr

>91,00

17,15

o,m

94Zr

>98,00

17,38

o,m

96Zr

>86,00

2.80

o,m

Zirconium was discovered in 1789 by Martin Heinrich Klaproth. Its name originates with the Arabic word zargun, meaning “gold color.”

Zirconium may exist as a hard, lustrous, silvery-gray, crystalline scale or as a bluish-black amorphous powder. It is corrosion-resistant. It starts as a close-packed hexagonal lattice and transforms to a body-centered cubic structure at 865 ºC. It is soluble in hot, very concentrated acids and aqua regia and is insoluble in water and cold acids. It exhibits quadrivalency in most of its compounds, although divalent and trivalent compounds also exist.

Solid metal zirconium is stable in air at ordinary temperatures, but it reacts slowly at 200 ºC. The reaction is more rapid at high temperatures. Reactions with hydrogen occur at temperatures of 300-1000 ºC, forming ZrH2, a brittle dihydride. Zirconium combines with halogens at high temperatures, forming tetrahalides. Reactions occur in the range of 200-400 ºC. Solid tetrahalides sublime above 300 ºC. Although stable to most acids, the metal is attacked by concentrated hydrochloric and sulfuric acids under boiling conditions or in the presence of aqua regia or hydrofluoric acid. The metal is stable in organic acids under all conditions.

The most important applications of zirconium involve its alloy, Zircaloy, which offers excellent mechanical and heat-transfer properties and great resistance to corrosion and chemical attack. Other uses are as an ingredient of explosive mixtures; as a “getter” in vacuum tubes; in making flash bulb, flash powder and lamp filaments; in rayon spinnerets; and in surgical appliances.

A suspected carcinogen, zirconium is not permitted in cosmetics (per the United States Food and Drug Administration). It is flammable and explosive in the form of dust, powder, borings or shavings.

Zirconium
40
91.224

Mo

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

92Mo

75,00-98,70

14,77

m,o

94Mo

>98,00

9,23

m,o

95Mo

>94,30

15,90

m,o

96Mo

>95,00

16,68

m,o

97Mo

>96,60

9,56

m,o

98Mo

>98,40

24,19

m,o

100Mo

90,00-99,86

9,67

m,o

Molybdenum was discovered in 1781 by Carl William Scheele. Its name originates with the Greek word molybdos, meaning “lead.” It does not occur free in nature, and it is a necessary trace element in plant nutrition.

Molybdenum is a silvery-white metal or grayish-black powder with a cubic crystalline structure. It has high strength at very high temperatures and oxidizes rapidly above 1000 °F in air at sea level, but it is stable in an upper atmosphere. It is insoluble in hydrochloric or hydrofluoric acids, ammonia, sodium hydroxide, water or dilute sulfuric acid, and soluble in hot concentrated sulfuric or nitric acids. At ordinary temperatures, molybdenum metal is fairly stable to air, water and most mineral acids. The metal shows high resistance to HCl, H2SO4, HF and H3PO4, and most other mineral acids in the absence of any oxidizing agent. However, the metal is attacked by nitric acid and aqua regia.

Molybdenum is very stable to oxygen at ambient temperatures; however, when heated in air or oxygen to red heat, the metal readily converts to trioxide. When bromine vapor is passed over molybdenum metal at 600-700 ºC in an atmosphere of nitrogen, the product is trimeric molybdenum dibromide. The metal has very little affinity for hydrogen, even in a finely divided state. When heated with carbon monoxide at temperatures above 1000 ºC, no carbonyl is obtained, but a thin film of molybdenum carbide forms over the metal. The metal is used as the starting material to prepare many of its salts.

The largest quantities of molybdenum produced are consumed in the steel industry. The metal is incorporated to impart high resistance and hardness to the steel, as well as to improve its mechanical properties. In the chemical industry, molybdenum compounds are used widely in coloring agents and solid lubricants. Molybdenum compounds are also used as catalysts in many oxidation-reduction reactions and in petroleum refining for production of high-octane gasoline. 

Molybdenum
42
95.94

Ru

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

96Ru

>96,00

12,70

m

100Ru

>97,00

12,60

m

101Ru

>98,00

17.00

m

102Ru

>98,00

31.60

m

104Ru

>99,50

18,70

m

Ruthenium was discovered in 1844 by Karl Karlovich Klaus. It takes its name from the Latin name Ruthenia, meaning “Russia.”

Ruthenium is a hard, silvery-white solid with a hexagonal close-packed crystal structure. It is insoluble in water, acids and aqua regia, and it is attacked by alkaline oxidants and by fused alkalis. When heated in air to 500-700 ºC, ruthenium converts to its dioxide, a black crystalline solid of rutile structure. A trioxide of ruthenium is also known and is formed when the metal is heated above 1000 ºC. Above 1100 ºC the metal loses weight because the trioxide partially volatizes.

Halogens react with ruthenium metal at elevated temperatures. The metal is attacked by chlorine water, bromine water or alcoholic solution of iodine at ambient temperatures. When finely divided metal is heated with carbon monoxide under 200 atm pressure, it converts to pentacarbonyl, a colorless liquid that decomposes to diruthenium nonacarbonyl, a yellow crystalline solid. Ruthenium reacts with cyclopentadiene in ether to form a sandwich complex, a yellow crystalline compound, bis(cyslopentadiene) ruthenium(0), also known as ruthenocene.

Ruthenium alloyed to platinum, palladium, titanium or molybdenum has many applications. Because ruthenium is an effective hardening element for platinum and palladium, such alloys have high resistance to corrosion and oxidation and are used to make electrical contacts for resistance to severe wear. Ruthenium-palladium alloys are used in jewelry, decorations and dental work. The addition of 0.1% ruthenium markedly improves the corrosion resistance of titanium. Ruthenium alloys also make tips for fountain pen nibs, instrument pivots and electrical goods.

Ruthenium
44
101.07

Pd

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

102Pd

>96,00

1.02

m

104Pd

86,00-98,00

11.14

m

105Pd

94,00-98,00

22.33

m

106Pd

95,00-98,00

27.33

m

108Pd

>99,00

26.46

m

110Pd

>99,00

11.72

m

Palladium was discovered in 1803 by William Hyde Wollaston. It is named for the asteroid Pallas, which was discovered at about the same time, as well as for the Greek name Pallas, goddess of wisdom.

A silver-white, ductile metal, palladium has a face-centered cubic crystalline structure and does not tarnish in air. It is the least noble (most reactive) of the platinum group and can absorb up to 800 times its own volume of hydrogen. Upon doing so, the metal swells, becoming brittle and cracked. Such absorption of hydrogen decreases the electrical conductivity of the metal. Also, such absorption activates molecular hydrogen, dissociating it to atomic hydrogen. It is attacked by hot, concentrated nitric acid and boiling sulfuric acid. It is soluble in aqua regia, fused alkalis, hot nitric acid and boiling sulfuric acid, and insoluble in organic acids and water. It is a good electrical conductor. It is nontoxic and noncombustible, except as dust. The metal forms mostly bivalent compounds, although a small number of tetravalent and fewer trivalent compounds are known. Palladium exhibits a strong tendency to form complexes, and it dissolves more oxygen in its molten state than in solid form. Palladium reacts with fluorine and chlorine at 500 ºC, forming its halides. Hydrochloric acid has no effect on the metal.

One of the most important applications of palladium is to catalyze hydrogenation, dehydrogenation and petroleum cracking. Such reactions are widely employed in organic syntheses and petroleum refining. Palladium and platinum are installed in catalytic converters in automobiles to cut down the emission of unsaturated hydrocarbon gases.

Palladium
46
106.42

Ag

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

107Ag

>99,00

51,84

m

109Ag

>99,00

48,16

m

Silver has been known since ancient times. It is mentioned in Genesis, and slag dumps in Asia Minor and on islands in the Aegean Sea indicate that humans learned to separate silver from lead as early as 3000 BC. It takes its name from the Anglo-Saxon word siolfur, meaning “silver” — the origin of the symbol Ag is the Latin word argentum, also meaning “silver.”

A white metal with brilliant metallic luster and face-centered cubic crystals, silver also has the highest electrical and thermal conductivity of all metals. It resists oxidation but tarnishes in air through reaction with atmospheric sulfur compounds, as well as with mercury. It is soluble in nitric acid and alkali cyanide solutions, and insoluble in water and alkalis. At ordinary temperatures, silver is not affected by either dry or moist air. At a temperature just above its melting point, silver absorbs a large volume of oxygen, greater than ten times its own volume. Such oxygen absorption, however, drops dramatically below its melting point; just before solidification, the absorbed oxygen is ejected violently. Silver also absorbs hydrogen above 800 ºC. Exposure of pure silver at about 810 ºC alternately to both hydrogen and oxygen gases embrittles the metal. Silver reacts with halogens at elevated temperatures, forming halides. It is attacked by nitric acid at all concentrations. It dissolves very slowly in hot concentrated sulfuric acid, forming silver sulfate. It is attacked by ozone, hydrogen peroxide, chromic acid, ferric sulfate and permanganate solutions.

Silver and its alloys and compounds have numerous applications. As a precious metal, silver is used in jewelry. One of its alloys, sterling silver, containing 92.5 weight % silver and 7.5 weight % copper, is a jewelry base that is used in tableware and decorative pieces as well. The metal and its copper alloys are also used in coins. Silver-copper crazing alloys and solders have applications in automotive radiators, heat exchangers, electrical contacts, steam tubes, coins and musical instruments. Some other uses of silver metal include electrodes, catalysts, mirrors and dental amalgam. Silver is used as a catalyst in oxidation reductions involving conversions of alcohol to aldehydes, ethylene to ethylene oxide, and ethylene glycol to glyoxal. Many silver compounds — such as silver nitrate, silver chloride and silver oxides — have wide commercial applications: their most important uses are in photography and batteries.

All water-soluble silver salts are toxic, and ingestion can cause severe poisoning. Silver is listed by the United States Environmental Protection Agency as one of the priority pollutant metals in the environment.

Silver
47
107.87

Cd

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

106Cd

81,00-99,00

1.25

m,o

108Cd

66,00-99,00

0.89

m,o

110Cd

95,00-99,00

12,49

m,o

111Cd

95,00-97,50

12.80

m,o

112Cd

88,00-98,00

24,13

m,o

113Cd

93,00-96,00

12,22

m,o

114Cd

99.00

28.73

m,o

116Cd

99,00

7.49

m,o

Cadmium was discovered in 1817 by Friedrich Strohmeyer. Its name originates with the Latin word cadmia (meaning "calamine" or "zinc carbonate") as well as the Greek word kadmeia, with the same meaning.

Cadmium is a bluish-white lustrous soft metal with a close-packed hexagonal system. It is insoluble in water. The metal is slowly oxidized in moist air at ordinary temperatures, forming a protective coating of cadmium oxide. The element combines with many nonmetals upon heating, forming its binary salts. It combines with halogens when heated, forming the corresponding halides. The metal is attacked by mineral acids. Reactions with hot dilute nitric acid give various oxides of nitrogen and hydrogen. Aqueous solutions of alkali hydroxides do not attack cadmium.

Cadmium replaces elements that are less electropositive in the activity series from their salt solutions. It can displace a number of metals that are less active, such as copper, lead, silver, mercury, tin and antimony from their aqueous salt solutions. It is used for electroplating, to impart a protective coating on iron and steel. It provides resistance against caustic alkalis.

A major application is in the nickel-cadmium storage battery, where it enhances long service life and a wide operating range. Cadmium alloys also find wide applications in bearing metals, solders, fusible metals, electrical conductors, power transmission wires and jewelry. Cadmium electrodes are used in photoelectric cells, cadmium vapor lamps and selenium rectifiers. Graphite impregnated with cadmium is used in electrical controller switches, oil-less bearings and busing lines. Cadmium rods are used in nuclear reactors to absorb low-energy neutrons.

Cadmium
48
112.41

In

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

113In

>93,00

4,29

m

115In

99.99

95,71

m

Indium was discovered in 1863 by Ferdinand Reich and Theodore Richter. It is named for the indigo line in its atomic spectrum.

A silver-white, lustrous, soft metal, indium is highly malleable and ductile, with a face-centered tetragonal crystalline structure. It is soluble in acids and insoluble in alkalis. It is corrosion-resistant at room temperatures but oxidizes readily at higher temperatures. It is nontoxic and becomes a superconductor at -269.8 ºC.

Indium is stable in air at ambient temperatures. At red heat, it oxidizes to indium trioxide. Three other oxides of indium are known: the suboxide, the monoxide, and the sesquioxide — a mixture of the trioxide and the monoxide. Chemical properties of indium are similar to those of aluminum. When heated with chlorine at 200 ºC, indium becomes a dichloride. However, in the presence of excess chlorine, indium trichloride is formed. Similar reactions with other halogens. Indium dissolves in mineral acids. Concentration or evaporation of the solution produces corresponding salts. With sulfuric acid, it forms indium trisulfate and indium hydrogen sulfate. The metal combines with sulfur and phosphorus on heating, forming the sulfide and phosphide salts. Metalloid elements, such as arsenic, antimony, selenium, and tellurium, also combine with indium at elevated temperatures, forming their respective binary salts. Indium combines with several metals, such as sodium, potassium, magnesium, iron, palladium, platinum, lanthanum and cerium, forming semiconductor-type intermetallic compounds.

A major use of indium metal is in the production of bearings for automobile and aircraft engines. The addition of indium improves strength and hardness of bearings as well as their resistance to corrosion and fatigue. Electroplated coatings of indium are applied onto aluminum for electrical wiring and as indium oxide coatings in sodium-vapor lamps. In the semiconductor industry, indium is used as a doping agent to obtain p-type germanium. Other applications are in glass-to-metal seals, in electroluminescent panels, as conductive coatings on glasses and ceramics, and in nuclear reactor control rods.

Indium
49
114.82

Sn

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

112Sn

99,60

0,97

m

114Sn

71,00-87,10

0,65

m

115Sn

51,00-69,00

0,34

m

116Sn

96,60-99,50

14,54

m

117Sn

87,00-97,50

7,68

m

118Sn

99,50

24,22

m

119Sn

90,50-97,40

8,59

m

120Sn

98,80-99,70

32,59

m

122Sn

93,40-99,30

4,63

m

124Sn

99,90

5,79

m

Tin was known in ancient times and is mentioned in the Old Testament. Early metal workers found it too soft for most purposes, but when mixed with copper, it produced bronze, which was suitable for many handcrafted items, especially tools and weapons. Tin takes its name from the Anglo-Saxon word tin, and the symbol Sn derives from the Latin word stannum, meaning “tin.”

Tin is a silvery-white metal at ordinary temperatures, slowly changing to gray below 13.2 ºC; it is soft, malleable and somewhat ductile. Tin has two allotropic forms: white tin, the beta form, which has a tetragonal structure and a density of 7.28 g/cm3, its color slowly changing from white to gray when cooled below 13.2 º C, converts to gray tin, the alpha form, with a density of 5.75 g/cm3. The presence of small amounts of antimony or bismuth prevents this transformation from white to gray tin.

Tin is insoluble in water; soluble in HCl, H2SO4, aqua regia and alkalis; and slightly soluble in dilute nitric acid. At ordinary temperatures, tin is stable in air, actually forming a very thin protective oxide film. In powder form, and especially in the presence of moisture, it oxidizes. When heated with oxygen, it forms tin oxide. Tin reacts with all halogens to form their halides. Tin is attacked by concentrated acids; with dilute acids, the reaction may be slow or very slow. The metal readily reacts with hot concentrated hydrochloric acid and aqua regia, but more slowly with cold dilute hydrochloric acid. The reaction also is slow with hot dilute sulfuric acid, which dissolves the metal, particularly in the presence of an oxidizing agent. The reaction with nitric acid is also generally slow. Hot concentrated acid converts the metal to an insoluble hydrate tin oxide. Reactions are more rapid with moist sulfur dioxide or sulfurous acid, chlorosulfonic or pyrosulfuric acids. Organic acids — such as acetic, oxalic and citric acids — react slowly with the metal, particularly in the presence of air or an oxidizing agent. The metal is stable in dilute solutions of ammonia or sodium carbonate. Tin dissolves in solutions of oxidizing salts such as potassium chlorate or potassium persulfate. The metal does not react with neutral salts in aqueous solutions; however, it reacts slowly with neutral salts in air. The metal does not combine directly with hydrogen, nitrogen or ammonia gas.

Today, tin is used for plating steel to make “tin” cans for preserving food. Also, tin coats other metals to prevent corrosion. An important application of tin is to produce "float glass" (made by floating molten glass on molten tin) for windows. A number of tin alloys have wide industrial applications and include bronze, solder, Babbit metal, White metal, type metal, fusible metal and phosphor bronze. A tin-niobium alloy that is superconducting at low temperatures is used in the construction of super magnets. Tin is also used in wrapping foil and collapsible tubes.

All organic tin compounds are toxic.

Tin
50
118.69

Sb

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

121Sb

>99,00

57.30

m/o

123Sb

>98,00

42.7

m

123Sb

94,00-98,00

42.7

o

Antimony was recognized in compounds by the ancients and was known as a metal at the beginning of the 17th century, possibly much earlier. Its name derives from the Greek words anti + monos, meaning “not alone,” and its symbol, Sb, comes from the Latin word stibium (“antimony”). Its most important mineral is stibnite, which formed the basis of black eye makeup in biblical times.

The French chemist Nicolas Lémery conducted early studies on antimony chemistry. It is a silvery-white, brittle, metallic element with a crystal system-hexagonal structure. It exists in two unstable allotropic forms – a yellow modification and a dark-gray lustrous amorphous powder – both of which revert to crystalline form.

Antimony is stable in dry air, not readily attacked by moisture, but slowly oxidizes in moist air. Oxidation may result under controlled conditions, forming tri-, tetra- and pentaoxides. Antimony reacts with sulfur, combining in all proportions, forming tri- and pentasulfides. It is oxidized by nitric acid, forming a gelatinous precipitate of hydrated antimony pentoxide. It does not react with cold dilute sulfuric acid; however, reaction occurs in hot concentrated acid: an oxysulfate of indefinite composition and low acid-solubility is formed. It reacts with hydrofluoric acid to form soluble antimony trifluoride and pentafluoride.

Antimony alloys have many commercial applications. The metal makes its alloys hard and stiff and imparts resistance to corrosion; such alloys are used in battery grids and parts, tank linings, pipes and pumps. The lead plates in lead storage batteries comprise 94% lead and 6% antimony. Babbit metal — an alloy of antimony, tin, and copper — is used to make antifriction machine bearings. Alloys made from very high-purity-grade antimony with indium, gallium and bismuth, are used as infrared detectors, diodes, Hall effect devices and thermoelectric coolers.

Antimony
51
121.75

Te

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

120Te

99,40-99,90

0,09

m,o

122Te

99,90

2,55

m,o

123Te

99,90

0,89

m,o

124Te

99,90

4,72

m,o

125Te

99,00-99,90

7,07

m,o

126Te

96,60-99,90

18,84

m,o

128Te

98,80-99,90

31,74

m,o

130Te

>99,70

34,08

m,o

Tellurium was discovered in 1783 by Franz-Joseph Müller Freiherr von Reichenstein, an Austrian mineralogist and mining engineer. Its name originates with the Latin word tellus, meaning “earth.”

Tellurium is a silvery white, lustrous solid. It is soluble in sulfuric acid, nitric acid, potassium hydroxide, potassium cyanide solution, caustic potash, and solutions of alkali metal cyanides. It is insoluble in water, carbon disulfide, benzene and hydrochloric acid. It burns in air with a greenish-blue flame; the combustion product is dioxide, the most stable oxide of the metal. Tellurium also forms other oxides: monoxide, trioxide and pentoxide. It combines with halogens, forming halides at different oxidation states. It also forms a black dichloride and a brown dibromide, usually by its reaction with dichlorodifluoromethane and trifluorobromomethane. It forms binary tellurides with several metals: the reaction is carried out by heating tellurium with a metal in stoichiometric amounts in the absence of air in an evacuated ampoule. Tellurium reacts with halides of several metals, when heated in a stream of hydrogen, to produce metal tellurides.

Small amounts of tellurium are added to stainless steel and copper to improve their machinability. It enhances the strength and hardness of lead and protects lead from the corrosive action of sulfuric acid. Tellurium also is a strong chilling agent in iron castings: it controls the chill and imparts a tough abrasion resistance to the surface. Tellurium is a curing agent for natural and synthetic rubber: it improves mechanical properties of the rubber, imparting resistance to heat and abrasion. It is a coloring agent in glass, ceramics and enamels. Traces of tellurium incorporated into platinum catalysts make the catalytic hydrogenation of nitric oxide favorable to the formation of hydroxylamine. A major application of tellurium is in semiconductor research. Tellurides of lead and bismuth are used in thermoelectric devices for power generation and refrigeration.

Human exposure to tellurium causes “garlic breath” due to dimethyl telluride, which persists for a considerable period after exposure. The toxic effects of tellurium are nausea, giddiness, headache, metallic taste, and dryness in the throat.

Tellurium
52
127.60

Xe

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

124Xe

99,90

0,10

g

126Xe

99,90

0.09

g

128Xe

99,90

1,91

g

129Xe

99,90

26,40

g

130Xe

99,90

4,10

g

131Xe

99,90

21,20

g

132Xe

99,90

26,90

g

134Xe

99,90

10,40

g

136Xe

99,90

8,90

g

Xenon was discovered in 1898 by Sir William Ramsay and Morris W. Travers. Its name derives from the Greek word xenos, meaning “stranger.”

Xenon is a filling gas for light bulbs in high-intensity lamps and in flash lamps for photography. It forms a beautiful blue glow under vacuum in electric discharge tubes. It can also be used as an anesthetic gas in surgery. Radioactive xenon is used as a biological tracer.

Xenon
54
131.29

Ba

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

130Ba

33,00

0.11

c

132Ba

12,80-40,70

0.10

c

134Ba

88,00

2.42

c

135Ba

>94,00

6,59

c

136Ba

61,00-95,40

7.85

c

137Ba

91,70

11.23

c

138Ba

99,80

71.70

c

Barium was discovered in 1808 by Sir Humphry Davy. Its name derives from the Greek word barys, meaning “heavy.”

A silvery-white, soft, ductile and somewhat malleable metal, barium gives off a green color in flame. It is extremely reactive and readily reacts with water, ammonia, halogens, oxygen and most acids. Barium metal reacts exothermically with oxygen at ambient temperatures, forming barium oxide; the reaction is especially violent when the metal is present in powder form. Barium also reacts violently with water, forming barium hydroxide and liberating hydrogen. It reacts violently with dilute acids, evolving hydrogen.

Barium is a strong reducing agent. It reduces oxidizing agents, reacting violently. It also combines with several metals — including aluminum, zinc, lead and tin — forming a wide range of intermetallic compounds and alloys.

The most important use of barium is as a scavenger in electronic tubes. The metal, often in powder form or as an alloy with aluminum, is employed to remove the last traces of gases from vacuum and television picture tubes. Alloys of barium have numerous applications, including battery performance and deoxidizing alloys to lower the oxygen content. Thin films of barium are used as lubricants on the rotors of anodes in vacuum x-ray tubes, as well as on alloys used for spark plugs. A few radioactive isotopes of barium find applications in nuclear reactions and spectrometry.

Finely divided barium powder is pyrophoric. It can explode in contact with air or oxidizing gases. It is likely to explode when mixed and stirred with halogenated hydrocarbon solvents. All barium salts, especially the water- and acid-soluble compounds, are highly toxic. Barium ion is a stimulant to the heart muscle and can cause death through ventricular fibrillation of the heart. Intake of a few grams of barium salt can be lethal to humans. The insoluble salts such as barium sulfate, however, have little toxic action.

Barium
56
137.327

Hf

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

174Hf

13,46

0,162

o

176Hf

84,60

5,260

o

177Hf

85,40

18,606

o

178Hf

92,40

27,297

o

179Hf

75,00

13,629

o

180Hf

>94,60

35,100

o

Hafnium was discovered in 1923 by Dirk Coster and George Charles von Hevesy. Its name derives from the Latin name Hafnia, meaning “Copenhagen.”

Hafnium is generally similar to zirconium. It has gray crystals, good corrosion resistance and high strength. It occurs as a close-packed hexagonal alpha form and as a body-centered cubic beta modification. It has a magnetic susceptibility of 0.42 x 10-6 emu/g at 25 ºC. It is insoluble in water, dilute mineral acids and nitric acid at all concentrations, and is soluble in hydrofluoric acid, concentrated sulfuric acid and aqua regia. The metal in bulk form does not react with most reagents at ordinary temperatures; however, the powdered metal or hafnium sponge may readily burn in air after being ignited with a spark. When heated to 360 ºC under water pressure, the metal is oxidized to hafnium oxide, forming a thin, protective surface oxide layer. Reaction with hydrofluoric acid at ordinary temperatures yields hafnium tetrafluoride. In finely divided form, hafnium is pyrophoric, igniting in air spontaneously; however, bulk metal reacts slowly in oxygen or air above 400 ºC. Reaction with hydrogen occurs around 700 ºC.

Hafnium is used in control rods for nuclear reactors. It has high resistance to radiation, as well as very high corrosion resistance. Another major application is in alloys with other refractory metals, such as tungsten, niobium and tantalum.

Hafnium
72
178.49

Ta

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

180Ta

0.26-5,70

0.012

m

181Ta

99.99

99.988

m

Tantalum was discovered in 1802 by Anders Ekeberg. Its name derives from the Greek name Tantalos, meaning “father of Niobe” (tantalum is closely related to niobium in the periodic table).

A gray, heavy, very hard metal, tantalum is also malleable and ductile and has a body-centered cubic lattice structure. It is soluble in fused alkalis and hydrofluoric acid, and insoluble in water, alcohol and all acids, except hydrofluoric and fuming sulfuric acids. Its aqueous solution chemistry is that of its pentavalent ion, Ta5+. The metal is attacked by hydrofluoric acid below 150 ºC. It also is dissolved by hot fuming sulfuric acid. It reacts with fluorine and chlorine on heating. The metal is immune to dilute aqueous alkalis but is attacked slowly by concentrated fused alkalis. Tantalum forms alloys with several metals: tantalum and its alloys have high melting points, high strength and ductility, and show excellent resistance to chemical attack. Tantalum carbide graphite composite is one of the hardest substances ever made; it has a melting point of 6700 ºC.

The pure metal is ductile and can be drawn into fine wire, which is used as a filament for evaporating aluminum and other metals. Tantalum and its alloys are used to build reactors, vessels, and crucibles for preparing and carrying out reactions involving many reactive intermediates. The metal and its alloys also are used to construct furnace parts, electrolytic capacitors, aircraft and missile parts, chemical process equipment, and nuclear reactors. Since it is nonreactive to body fluids and a nonirritant to body tissues, tantalum is used in making surgical appliances. Plate and sheet tantalum are applied in bone repair; foil and wire in nerve repair; and plate, gauge and sheet in the repair of abdominal muscle. Tantalum oxide is used to produce optical glasses of high refractive index. The oxide film on the metal makes it a rectifier for converting alternating current to direct current.

Tantalum
73
180.95

W

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

180W

>91,00

0,12

m,o

182W

91,50-99,10

26,50

m,o

183W

>99,00

14,31

m,o

184W

95,00-99,90

30,64

m,o

186W

99,79-99,90

28,43

m,o

Tungsten was discovered in 1783 by Fausto and Juan José de Elhuyar. Its name derives from the Swedish phrase tung sten, meaning “heavy stone.” The origin of the symbol W is wolfram, named for the tungsten mineral wolframite.

Tungsten is a grayish-white metal with a body-centered cubic crystalline structure. It is not found in its native form. It has a high electrical conductivity. It oxidizes in air at 400° C, the oxidation rate increasing rapidly with temperature or in the presence of an oxidizing agent such as potassium nitrate, potassium chlorate or lead dioxide. It is insoluble in water and practically insoluble in most acids and alkalis. It dissolves slowly in hot concentrated nitric acid, in saturated aqueous solution of sodium chlorate and basic solution of potassium ferricyanide. It is also solubilized by fusion with sodium hydroxide or sodium carbonate in the presence of potassium nitrate, followed by treatment with water. It is soluble in a mixture of nitric acid and hydrofluoric acid and is corroded by seawater. Compounds of lower oxidation states show alkaline properties and are less stable than those produced in higher oxidation states. Tungsten metal is not affected by aqueous alkalis at room temperature. Although it exhibits a high degree of resistance to most chemicals, it is readily oxidized by a number of oxidizing agents. Tungsten reacts with oxygen at high temperatures, and the finely-divided powder is pyrophoric. It reacts with all halogens and with ammonia at elevated temperatures, forming tungstic nitrides and amides.

Tungsten forms a number of compounds with nonmetals and light metalloid elements. Many are important refractory materials in commerce. In industry, tungsten is a very important metal with wide applications, due to its many outstanding physical properties. Among all the metals, it has the highest melting point and the lowest vapor pressure. At high temperatures it also has the highest tensile strength. The metal has an excellent resistance to corrosion and attack by mineral acids, as well as a thermal expansion comparable to that of borosilicate glass. It is used extensively in alloy steel to impart high strength and hardness.

Heavy metal alloys with nickel, copper and iron, produced by powder metallurgy, can be made machineable and moderately ductile for applications as high-density materials. They are used extensively in the tool and die industry for drilling and cutting tools, sand-blasting nozzles, armor-piercing bullets, and studs to increase tire traction. Among nonferrous tungsten alloys, its alloys with copper and silver are used as electrical contacts and switches and with molybdenum in aerospace components. Unalloyed tungsten has several major applications: in electric lamp filaments for light bulbs, as electrodes in arc-welding, in heating elements for high-temperature furnaces, in electron and television tubes, in glass-to-metal seals and in solar energy devices. In its finely divided form, tungsten is highly flammable and may ignite spontaneously.

Tungsten
74
183.85

Re

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

185Re

93,00-96,80

37,40

m

187Re

>99,00

62,60

m

Rhenium was discovered in 1925 by Walter Noddack, Ida Tacke and Otto Berg. Its name derives from the Greek name Rhenus, meaning “River Rhine.”

A silver-white solid or gray-to-black powder, rhenium has a hexagonal crystal system, which it retains all the way to its melting point. It has the widest range of valences of any element. It is insoluble in water, practically insoluble in hydrochloric acid, soluble in dilute nitric acid and hydrogen peroxide, and slightly soluble in sulfuric acid. It is not attacked by sea water, hydrochloric acid, cold sulfuric or hydrofluoric acids. It is attacked by strong oxidizing agents (nitric and sulfuric acids). In compact or massive form, it is stable at ordinary temperatures. Rhenium reacts with all halogens, including iodine, to yield halides in several valence states. However, oxidizing acids — such as nitric acid or hot sulfuric acid — vigorously react with the metal, forming perrhenic acid. Rhenium combines with phosphorus, arsenic, silicon, selenium and tellurium at elevated temperatures, forming binary compounds. The metal, however, is stable in hydrogen and nitrogen at high temperatures.

Rhenium is used in tungsten- and molybdenum-based alloys. It is used for filaments for ion gauges in mass spectrometers. Rhenium-tungsten alloys are used in thermocouples to measure temperatures up to 2200 ºC. Rhenium wire is used in flash bulbs for photography. Rhenium compounds also are used as catalysts in hydrogenation and hydrofracking reactions in petroleum refining.

Rhenium
75
186.21

Os

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

184Os

>96,90

0.02

m

186Os

>99,00

1,59

m

187Os

>99,00

1,96

m

188Os

>94,00

13,24

m

189Os

>99,00

16,15

m

190Os

>99,00

26,26

m

192Os

>99,00

40,78

m

Osmium was discovered in 1803 by Smithson Tennant. Its name derives from the Greek word osme, meaning “smell.”

A bluish-white, lustrous, brittle and fairly hard metal of the platinum group, osmium has a close-packed hexagonal system. On heating in air, it gives off the poisonous fume of osmium tetroxide. It has the highest specific gravity and melting point of the platinum metals. It is metallurgically unworkable and has a magnetic susceptibility of 0.052 x 10-6 cm3/g. It is insoluble in water, HCl, H2SO4 and ammonia; slightly soluble in nitric acid and aqua regia; and solubilized by fusion with caustic soda, sodium peroxide, potassium chlorate and the mass dissolved in water. In its finely divided form, it reacts slowly with oxygen or air at ambient temperatures to form osmium tetroxide. The bulk metal is stable in oxygen at ordinary temperatures but reacts at 200 ºC, forming osmium tetroxide. Osmium is stable in mineral acids, even under boiling conditions. The metal, however, is attacked by fused alkalis such as caustic soda and caustic potash, particularly in the presence of an oxidizing agent such as sodium peroxide, sodium hypochlorite or sodium nitrite, forming osmates.

The commercial applications of osmium are limited and considerably fewer than those of other platinum group metals. Its alloys are very hard and are used to make tips of fountain-pen nibs, phonograph needles and pivots. The metal also exhibits effective catalytic properties in hydrogenation and other organic reactions. Such catalytic applications, however, are limited; osmium fails to replace other noble metals, particularly palladium and platinum, which cost less and are more effective as catalysts.

Osmium
76
190.2

Ir

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

191Ir

>93,00

37,30

m

193Ir

62,70

98,40

m

Iridium was discovered in 1803 by French chemists Victor Collet-Descotils, Antoine François Comte de Fourcroy and Louis Nicolas Vauquelin, and British scientist Smithson Tennant. It is named for Iris, the Greek winged goddess of the rainbow and messenger of the Olympic gods.

A silver-white metal with low ductility and close-packed cubic crystals, iridium does not tarnish in air. On heating strongly, a slightly volatile oxide is formed. It is insoluble in acid and slowly soluble in aqua regia and in fused alkalis. It is the most corrosion-resistant element and is highly resistant to chemical attack at ordinary temperatures. At elevated temperatures (near 600 ºC), iridium metal combines with oxygen to form a coating of iridium dioxide. Similarly, the metal reacts with halogens only at elevated temperatures. Iridium forms alloys with several metals, mostly platinum group metals. Iridium does not react with concentrated acids or with molten alkalis.

The most important use of iridium is as an alloying metal for platinum and palladium. Iridium enhances the resistance of platinum to chemical attack and corrosion, as well as its enhancing hardness and tensile strength. Such alloys are used for jewelry, decorative purposes, electrical contacts, thermocouples, crucibles, electrodes, hypodermic needles and medical accessories. The radioisotope Iridium-192 is used in the examination of ferrous welds and in other radiographic applications.

Iridium
77
192.22

Pt

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

190Pt

4,00

0,014

m

192Pt

56,00

0,782

m

194Pt

>96,00

32,967

m

195Pt

>96,00

33,832

m

196Pt

>97,00

25,242

m

198Pt

>91,00

7,163

m

Platinum was discovered in 1735 by Antonio de Ulloa. Its name derives from the Spanish word platina, meaning “silver.”

Platinum is a silvery white, ductile, lustrous metal with face-centered cubic crystals. When heated, platinum absorbs large volumes of hydrogen. It is also a strong complexing agent. It has a vapor pressure of 0.00014 torr at its melting point. It has a magnetic susceptibility of 9.0 x 10-6 cm3/g. As a catalyst it is abnormally sensitive to poisons. It is insoluble in mineral and organic acids, and soluble in aqua regia. It is attacked by fused alkalis. It does not corrode or tarnish. At ordinary temperatures platinum is inert to practically all substances except aqua regia and, to a small extent, chlorine water. It reacts with oxygen only at elevated temperatures. Finely divided metal forms platinum oxide at about 500 ºC. Fused alkalis, particularly potassium and barium hydroxides, are corrosive to platinum. Platinum can be alloyed with many elements at elevated temperatures.

Platinum metal and its alloys have numerous applications. It is used extensively as a precious metal in the manufacture of jewelry. Other important applications include construction of laboratory crucibles and high-temperature electric furnaces, in instruments as thermocouple elements, as wire for electrical contacts, as electrodes, in dentistry, in cigarette lighters, and for coating missile and jet engine parts. Platinum is also used extensively as a catalyst in hydrogenation, dehydrogenation, oxidation, isomerization, carbonylation and hydrocracking. It is also used in organic synthesis and petroleum refining. An important application is in the catalytic oxidation of ammonia in Ostwald’s process in the manufacture of nitric acid. Platinum is installed in the catalytic converters in automobile engines for pollution control.

Platinum
78
195.09

Hg

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

196Hg

51,00-95,00

0.15

m,o

198Hg

94,00-99,00

9,97

m,o

199Hg

>90,00

16,87

m,o

200Hg

>95.0

23,10

m,o

202Hg

>99.00

29,86

m,o

204Hg

>96,00

6,87

m,o

Mercury was known and used in ancient civilizations. It has been found in Egyptian tombs dated from 1500 BC; the Greeks used it in ointments; and the Romans used it in cosmetics. It is named for the planet Mercury, and its symbol, Hg, originates with the Latin word hydrargyrum, meaning “liquid silver.”

Mercury is an extremely heavy liquid, the only metal and one of only two elements that naturally occur as a liquid in ambient temperatures. It does not wet glass, and it forms tiny globules. It is insoluble in hydrochloric acid, water, alcohol and ether; soluble in sulfuric acid upon boiling; and readily soluble in nitric acid and lipids. Mercury is stable in dry air or oxygen at ordinary temperatures; however, in the presence of moisture, oxygen slowly attacks the metal, forming red mercury(II) oxide. When the metal is heated in air or oxygen (to about 350 ºC), it is gradually converted to its oxide. Mercury readily combines with halogens at ordinary temperatures, forming mercury(II) halides. Its metal forms both mercury(I) and mercury(II) salts. Dilute sulfuric acid has no effect on the metal; nor does air-free hydrochloric acid. Water has no effect on mercury; nor does molecular hydrogen. However, atomic hydrogen readily combines with mercury vapors when exposed to radiation from a mercury arc, forming hydride.

Some of the most important uses of mercury are in the electrical and electrolytic industries, including batteries and cells in portable radios, microphones, cameras, hearing aids, watches, smoke alarms, wiring and switching devices, mercury vapor lamps, fluorescent tubes, electrical discharge tubes, and mercury electrodes in electrolytic cells. Mercury cathodes are employed in the electrolysis of sodium chloride to produce caustic soda and chlorine. Another major use is in thermometers, manometers, barometers and other pressure-sensing devices. Mercury is also used as a catalyst in making urethane foams and vinyl chloride monomers; elemental mercury and its compounds have long been used as fungicides in paints and in agriculture; and its compounds are used in medicines, pigments and analytical reagents.

Elemental mercury and all of its compounds are highly toxic by all routes of exposure. The element has significant vapor pressure at ambient temperatures that can produce a severe inhalation hazard. Symptoms from short exposure to high concentrations of mercury vapors include bronchitis, coughing, chest pain, respiratory distress, salivation, diarrhea, tremors, insomnia and depression. Mercury can cause damage to the kidneys, liver, lungs and brain. Organomercury compounds and inorganic salt solutions can be absorbed into the body through skin contact, causing severe poisoning. It accumulates as Hg2+ in the brain and kidneys. The United States Environmental Protection Agency has classified mercury as one of the priority pollutant metals.

Mercury
80
200.59

Tl

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

203Tl

>99,00

29,52

m,o

205Tl

99,90

70,48

m,o

Thallium was discovered in 1861 by Sir William Crookes. Its name derives from the Greek word thallos, meaning “green twig” or “green shoot.”

A bluish-white, lead-like solid, thallium has tetragonal crystals. It oxidizes in air at room temperature. It is soluble in nitric and sulfuric acid, insoluble in water (but readily forms soluble compounds when exposed to air or water), and slightly soluble in hydrochloric acid. It reacts with water containing oxygen to form thallous hydroxide, a relatively strong base, absorbing carbon dioxide and attacking glass. It burns in fluorine with incandescence. It reacts with other halogens to form halides. It also combines with several elements, forming binary compounds.

Thallium and its compounds have limited applications, including insecticides and rodenticides. Thallium-mercury alloys are used for switches and closures for use at sub-zero temperatures; another use is in low-melting glasses for electronic encapsulation. Thallium sulfide is used in photocells.

Thallium and its compounds (particularly its soluble salts) can cause serious or fatal poisoning from accidental ingestion or external application. Acute symptoms include nausea, vomiting, diarrhea, weakness, pain in extremities, convulsions and coma. Chronic effects include weakness, pain in extremities, and rapid loss of hair. Thallium and its compounds are listed under federal toxics regulations. It is listed by the United States Environmental Protection Agency as a priority pollutant metal in the environment.

Thallium
81
204.38

Pb

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

204Pb

99,90

1,40

o,m

206Pb

>99,00

24,10

o,m

207Pb

72,00->99,00

22,10

o,m

208Pb

97,80->99,50

52,40

o,m

Lead has been known and used throughout history — ancient alchemists believed it was the world's oldest metal and devoted a great deal of time attempting to transmute it into gold. Its name originates with the Anglo-Saxon word lead, and its symbol — Pb — with the Latin word plumbum, meaning “liquid silver.”

Lead is a heavy, very soft, malleable, ductile solid with face-centered cubic crystals. It is soluble in dilute nitric acid and insoluble in water (but dissolves slowly in water containing a weak acid). It resists corrosion and is relatively impenetrable to radiation. It is a poor electrical conductor. Lead forms amphoteric compounds in +2 and +4 valence states, forming plumbous and plumbic salts, as well as plumbites and plumbates. Its divalent compounds are far more numerous than its tetravalent compounds. In its very finely divided form, lead is pyrophoric. The metal is not attacked by hot water; in hard water, however, the presence of small amounts of carbonate, sulfate or silicate ions forms a protective film on the metal's surface. Lead does not evolve hydrogen readily with acids. At ordinary temperatures, it is not readily attacked by sulfuric acid. Hydrofluoric acid also has little action on the metal. Organic acids in the presence of oxygen react slowly with lead, forming their soluble salts; thus acetic acid in the presence of oxygen forms lead(II) acetate. Lead combines with fluorine, chlorine and bromine, forming bivalent lead halides.

There are numerous applications for lead in all its forms: metal, alloys and compounds. It is useful in the construction of pipelines, plumbing fixtures, wires, ammunition, containers for corrosive acids, and as a shield against short-wavelength radiation. Both the metal and its dioxide are used in storage batteries. Several lead compounds — such as lead chromate, lead sulfate, lead tetroxide and the basic carbonate — have been used in paint.

Considered an acute and chronic toxicant, lead can cause acute ataxia, headache, vomiting, stupor, hallucination, tremors and convulsions, along with chronic symptoms including weight loss, anemia, kidney damage, memory loss and brain damage.

Lead
82
207.2

La

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

138La

5,90

0,09

o

139La

99.99

99.91

o

Lanthanum was discovered in 1839 by Carl Gustaf Mosander. Its name derives from the Greek word lanthanein, meaning “to lie hidden.”

Lanthanum is a white, malleable, soft, ductile metal with a hexagonal close-pack crystal system that transforms to face-centered cubic crystals at 310 ºC, further transforming to a body-centered cubic allotropic modification at 868 ºC. It oxidizes rapidly in air and corrodes in moist air. It is soluble in acids and decomposes in water to lanthanum hydroxide and hydrogen. Combustion in air or oxygen produces lanthanum sesquioxide. Lanthanum reacts vigorously when heated with halogens above 200 ºC, forming lanthanum halides. It combines with nitrogen, carbon, sulfur and phosphorus at elevated temperatures, forming binary salts. With metalloid elements such as boron, silicon, selenium and arsenic, similar reactions occur at high temperatures, forming similar binary compounds. In nature, lanthanum never occurs in its free state; it is always found associated with other rare-earth metals.

As a pure metal, lanthanum has limited uses except for research purposes; however, in alloy form, it has several metallurgical applications. When alloyed with iron, chromium, nickel or molybdenum, it improves resistance of these metals to oxidation. It also improves the impact strength, fluidity, ductility and other mechanical properties of the alloys.

In its powdered form, lanthanum ignites spontaneously. Exposure to lanthanum may delay blood clotting and cause liver injury.

Lanthanum
57
138.91

Ce

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

136Ce

30,00-53,40

0.19

o

138Ce

13,60-41,60

0.25

o

140Ce

>99,00

88.48

o

142Ce

93,50-95,10

11.08

o

Cerium was discovered in 1803 by Wilhelm von Hisinger, Jöns Jakob Berzelius and Martin Heinrich Klaproth. It is named after the asteroid Ceres, which had been discovered just two years earlier.

Cerium is a grayish lustrous metal that is malleable and has four allotropic modifications. The common γ-form occurs at ordinary temperatures and atmospheric pressures. The β-form occurs at -16 ºC, the α-form occurs below -172 ºC, and the δ-form occurs at elevated temperatures above 725 ºC. Cerium reacts with water, while the metal is stable in dry air at ordinary temperatures. 

Compounds of cerium have many important industrial applications as catalysts, especially in the glass industry. Misch metal, an alloy of cerium with other lanthanides, is a pyrophoric substance and is used to make gas lighters and ignition devices. Some other applications of the metal or its alloys are in solid-state devices, rocket propellant compositions, as a getter in vacuum tubes, and as a diluent for plutonium in nuclear fuel. 

Cerium
58
140.115

Nd

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

142Nd

>97,50

27.13

o

143Nd

>79,00

12.18

o

144Nd

>98,50

23,80

o

145Nd

>94,00

8,30

o

146Nd

>98,80

17.19

o

148Nd

>97,40

5.76

o

150Nd

>97,60

5.64

o

Neodymium was discovered in 1885 by Carl F. Auer von Welsbach. Its name derives from the Greek phrase neos didymos, meaning “new twin.”

Neodymium is a silvery-white, soft, malleable metal that tarnishes easily. It liberates hydrogen from water and is soluble in dilute acids. It has a high electrical resistivity and is paramagnetic. It is readily cut and machined. It exists in two allotropic forms: an alpha hexagonal form, and a beta form that has body-centered cubic crystal structure. The alpha allotrope converts to beta modification at 868 ºC. Neodymium corrodes slowly in dry atmosphere at ambient temperatures; however, in moist air, the rate of oxidation is faster, forming a hydrated oxide. Neodymium combines with many nonmetallic elements — including hydrogen, nitrogen, carbon, phosphorus and sulfur — at elevated temperatures, forming their binary compounds. The metal dissolves in dilute mineral acids, but concentrated sulfuric acid has little action on it. Neodymium is a moderately strong reducing agent. It reduces several metal oxides — such as magnesia, alumina, silica and zirconia — at elevated temperatures, converting these oxides to their metals.

The pure metal has very little commercial application; however, neodymium in the form of alloys has found some important but limited applications in metallurgy. It is added to cast iron, magnesium, aluminum, zirconium and titanium alloys, imparting high-temperature strength and creep resistance to these alloys. It acts as a “getter” for oxygen, sulfur, hydrogen, nitrogen and other elements. Small quantities of neodymium salts are used as a coloring agent for glass and porcelain, imparting a red color.

Neodymium
60
144.24

Sm

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

144Sm

88,00-93,00

3,10

o

147Sm

94,00-96,50

15.00

o

148Sm

91,00-96,50

11,30

o

149Sm

>94,00

13,80

o

150Sm

>94,00

7,40

o

152Sm

>98,40

26.70

o

154Sm

>98,50

22.70

o

Samarium was discovered in 1879 by Paul-Émile Lecoq de Boisbaudran. It is named for the mineral samarskite, from which Lecoq de Boisbaudran isolated the new element. The mineral in turn takes its name from Vasili Samarsky-Bykhovets, the one-time chief of staff of the Russian Corps of Mining Engineers.

Samarium is a hard, brittle, yellow metal, which quickly develops an oxide film in air. Its hardness is similar to that of iron. It exhibits two crystal forms: an alpha form, with a rhombohedral crystal structure at ordinary temperatures, changes to the body-centered cubic form at 917 ºC. The densities of the two forms are 7.52 g/cm3 and 7.40 g/cm3, respectively. Samarium is insoluble in water and soluble in acid. It is stable in dry air at ordinary temperatures; however, it oxidizes in moist air, forming an oxide coating. The metal ignites in air at about 150 ºC. It is an active reducing agent (it reduces several metal oxides to metals), and it liberates hydrogen from water. Among samarium's trivalent salts, the sesquioxide is commercially important, and the divalent compounds are primarily halides. The trivalent salts of these halogens are more stable than their divalent counterparts.

Samarium salts are used in optical glass, capacitors, thermoionic generating devices and sensitizers of phosphors. The metal is doped with calcium fluoride crystals for use in lasers. It is also used with other rare earths for carbon-arc lighting. Its alloys are used in permanent magnets.

Samarium
62
150.36

Eu

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

151Eu

97,50-99,24

47,80

o

153Eu

99,70

52,20

o

The discovery and isolation of europium are generally credited to Eugène-Anatole Demarçay, who successfully isolated the element in 1901. He named it after Europe.

A steel-gray metal with a body-centered cubic crystal lattice, europium is difficult to prepare. It is quite soft and malleable. It oxidizes rapidly in air and may burn spontaneously. It is the most reactive of the rare earth metals; it liberates hydrogen from water; and it reduces metallic oxides. It reacts with water and is soluble in liquid ammonia.

Europium is used for the capture of thermal neutrons for nuclear control rods in atomic power stations. While its salts are used in coatings for cathode ray tubes in color televisions, organoderivatives are used in nuclear magnetic resonance (NMR) spectroscopy.

Europium
63
151.97

Gd

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

152Gd

30,60-34,80

0,20

o

154Gd

52,10-64,20

2.18

o

155Gd

>90,00

14,80

o

156Gd

>93,30

20.47

o

157Gd

88,40

15,65

o

158Gd

>97,30

24.84

o

160Gd

Contact us

21.86

o

 Gadolinium was discovered by Jean de Marignac in 1880. It is named for gadolinite, one of the minerals in which it was found, which was in turn named for chemist Johan Gadolin.

A colorless or light yellow lustrous metal, gadolinium occurs in a hexagonal close-packed crystalline form, known as alpha-gadolinium, which transforms to a body-centered cubic allotropic form, beta-gadolinium, when it reaches 1262 ºC. It exhibits a high degree of magnetism, especially at lower temperatures. Its salts are colorless. It has a vapor pressure of 9.0 torr at 1800 ºC. It also has superconductive properties. It is combustible and burns in air to form the oxide. It reacts slowly in water, is soluble in dilute acid, and is insoluble in water. All of its compounds are trivalent. Although the metal is stable in air at ordinary temperatures, it burns in air when heated to 150-180 ºC. Gadolinium is a strong reducing agent. It reduces oxides of several metals such as iron, chromium, lead, manganese, tin and zirconium into their elements. It burns in halogen vapors above 200 ºC, forming gadolinium(III) halides.

The most important application of this metal is as control rod material for shielding in nuclear power reactors. Other uses are in thermoelectric generating devices, as a thermoionic emitter, in yttrium-iron garnets in microwave filters to detect low-intensity signals, as an activator in many phosphors, for deoxidation of molten titanium, and as a catalyst.

Gadolinium
64
157.25

Yb

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

168Yb

35,00-87,00

0.13

o

170Yb

>70,00

3,05

o

171Yb

95,00-97,50

14,30

o

172Yb

>98,00

21,90

o

173Yb

>92,00

16.12

o

174Yb

>99,00

31,80

o

176Yb

>99,60

12,70

o

Ytterbium was discovered in 1878 by Jean de Marignac. It is named after the village of Ytterby, near Vaxholm, Sweden.

Ytterbium is a silvery, lustrous metal that is soft, malleable and ductile. The metal exists in two allotropic forms: an alpha form, which has a face-centered cubic structure and a density of 6.98 g/cm3 and is stable at room temperature; and a beta form, with a body-centered cubic modification and a density of 6.54 g/cm3.The beta form appears when the alpha form is heated to 798 ºC. Ytterbium reacts slowly with water and is soluble in dilute acids and liquid ammonia. It reacts with oxygen above 200 ºC. It forms two oxides: a monoxide and a more stable sesquioxide. The metal dissolves in dilute and concentrated mineral acids. Similar to other rare earth metals, ytterbium is corroded slowly at ordinary temperatures by caustic alkalis, ammonium hydroxide and sodium nitrate solutions. The metal also dissolves in liquid ammonia, forming a deep blue solution. Reactions with halogens are slow at room temperature but progress rapidly above 200 ºC, forming ytterbium trihalides. At elevated temperatures, ytterbium forms many binary, metalloid and intermetallic compounds with a number of elements.

Ytterbium metal has very little commercial use. In elemental form it is a laser source, a portable x-ray source, and a dopant in garnets. When added to stainless steel, it improves grain refinement, strength and other properties. Some other applications include carbon rods for industrial lighting, titanate-insulated capacitors, and additives to glass. The radioactive isotope Ytterbium-169 is used in portable devices to examine defects in thin steel and aluminum. The metal and its compounds are used in fundamental research.

Ytterbium
70
173.04

Lu

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

175Lu

99.80

97,41

o

176Lu

60,00-84,60

2,59

o

Lutetium was discovered in 1907 by Georges Urbain and Carl Auer von Welsbach. Its name originates with the Latin Lutetia, the name of a Roman town on the site of modern-day Paris.

Lutetium is a silvery-white, lustrous metal with a hexagonal close-packed structure. It is soft, ductile and slightly paramagnetic. It reacts slowly with water and is soluble in dilute acids. In aqueous media, lutetium occurs as tripositive Lu3+ ion. Aqueous solutions of all its salts are colorless, while in dry form they are white crystalline solids. Lutetium's soluble salts — such as chloride, bromide, iodide, nitrate, sulfate and acetate — form hydrates upon crystallization. The oxide, hydroxide, fluoride, carbonate, phosphate and oxalate of the metal are insoluble in water. The metal dissolves in acids, forming the corresponding salts upon evaporation of the solution and crystallization.

There is very limited commercial application for lutetium. The metal emits beta particles after thermal neutron activation, and it is used to catalyze organic reactions as well as for dating meteorites. Stable lutetium can be used in petroleum cracking in refineries, as well as for alkylation, hydrogenation and polymerization applications. Its synthetic isotope Lutetium-177, when bound to octreotate, is used experimentally in targeted radionuclide therapy for neuroendocrine tumors.

Lutetium
71
174.97

Dy

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

156Dy

18,00-20,70

0.06

o

158Dy

14,10-23,70

0,10

o

160Dy

67,70-70,10

2.34

o

161Dy

92,20-93,50

18.90

o

162Dy

>94,70

25,50

o

163Dy

>94,50

24.90

o

164Dy

95,90-98,45

28,20

o

 

Dysprosium
-
-

Er

Isotope

Maximum Available
Enrichment
(ATOMIC %)

Natural
Abundance
(ATOMIC %)

Chemical
Form(s)

162Er

28,20-39,20

0.14

o

164Er

51,30-75,30

1.61

o

166Er

94,70-98,10

33.60

o

167Er

95,30-96,30

22.95

o

168Er

>98,00

26.80

o

170Er

97,70-98,20

14.90

o

 

Erbium
-
-

Izotope Order

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