zirconium (Zr), chemical element, metal of Group 4 (IVb) of the periodic table, used as a structural material for nuclear reactors.

Element Properties
atomic number40
atomic weight91.22
melting point1,852 °C (3,366 °F)
boiling point3,578 °C (6,472 °F)
specific gravity6.49 at 20 °C (68 °F)
oxidation state+4
electron configuration[Kr]4d25s2

Properties, occurrence, and uses

Zirconium, obscure before the late 1940s, became a significant engineering material for nuclear energy applications because it is highly transparent to neutrons. The element was identified (1789) in zircon, ZrSiO4 (zirconium orthosilicate), from its oxide by the German chemist Martin Heinrich Klaproth, and the metal was isolated (1824) in impure form by the Swedish chemist Jöns Jacob Berzelius. The impure metal, even when 99 percent pure, is hard and brittle. The white, soft, malleable, and ductile metal of higher purity was first produced in quantity (1925) by the Dutch chemists Anton E. van Arkel and J.H. de Boer by the thermal decomposition of zirconium tetraiodide, ZrI4. In the early 1940s, William Justin Kroll of Luxembourg developed his cheaper process of making the metal based on the reduction of zirconium tetrachloride, ZrCl4, by magnesium. In the early 21st century, leading producers of zirconium included Australia, South Africa, China, and Indonesia; Mozambique, India, and Sri Lanka were additional producers.

Zirconium is relatively abundant in Earth’s crust, but not in concentrated deposits, and is characteristically observed in S-type stars. The mineral zircon, which is generally found in alluvial deposits in stream beds, ocean beaches, or old lake beds, is the only commercial source of zirconium. Baddeleyite, which is essentially pure zirconium dioxide, ZrO2, is the only other important zirconium mineral, but the commercial product is more cheaply recovered from zircon. Zirconium is produced by the same process as that used for titanium. These zirconium minerals generally have a hafnium content that varies from a few tenths of 1 percent to several percent. For some purposes separation of the two elements is not important: zirconium containing about 1 percent of hafnium is as acceptable as pure zirconium.

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The most important use of zirconium is in nuclear reactors for cladding fuel rods, for alloying with uranium, and for reactor-core structures because of its unique combination of properties. Zirconium has good strength at elevated temperatures, resists corrosion from the rapidly circulating coolants, does not form highly radioactive isotopes, and withstands mechanical damage from neutron bombardment. Hafnium, present in all zirconium ores, must be scrupulously removed from the metal intended for reactor uses because hafnium strongly absorbs thermal neutrons.

Separation of hafnium and zirconium is generally accomplished by a liquid-liquid countercurrent-extraction procedure. In the procedure, crude zirconium tetrachloride is dissolved in an aqueous solution of ammonium thiocyanate, and methyl isobutyl ketone is passed countercurrent to the aqueous mixture, with the result that the hafnium tetrachloride is preferentially extracted.

The atomic radii of zirconium and hafnium are 1.45 and 1.44 Å, respectively, while the radii of the ions are Zr4+, 0.74 Å, and Hf4+, 0.75 Å. The virtual identity of atomic and ionic sizes, resulting from the lanthanoid contraction, has the effect of making the chemical behaviour of these two elements more similar than for any other pair of elements known. Although the chemistry of hafnium has been studied less than that of zirconium, the two are so similar that only very small quantitative differences—for example, in solubilities and volatilities of compounds—would be expected in cases that have not actually been investigated.

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Zirconium absorbs oxygen, nitrogen, and hydrogen in astonishing amounts. At about 800 °C (1,500 °F) it combines chemically with oxygen to yield the oxide, ZrO2. Zirconium reduces such refractory crucible materials as the oxides of magnesium, beryllium, and thorium. This strong affinity for oxygen and other gases accounts for its use as a getter for removing residual gases in electron tubes. At normal temperatures in air, zirconium is passive because of the formation of a protective film of oxide or nitride. Even without this film, the metal is resistant to the action of weak acids and acidic salts. It is best dissolved in hydrofluoric acid, in which procedure the formation of anionic fluoro complexes is important in stabilizing the solution. At normal temperatures it is not particularly reactive but becomes quite reactive with a variety of nonmetals at elevated temperatures. Because of its high corrosion resistance, zirconium has found widespread use in the fabrication of pumps, valves, and heat exchangers. Zirconium is also used as an alloying agent in the production of some magnesium alloys and as an additive in the manufacture of certain steels.

Natural zirconium is a mixture of five stable isotopes: zirconium-90 (51.46 percent), zirconium-91 (11.23 percent), zirconium-92 (17.11 percent), zirconium-94 (17.40 percent), zirconium-96 (2.80 percent). Two allotropes exist: below 862 °C (1,584 °F) a hexagonal close-packed structure, above that temperature a body-centered cubic.

Compounds

Zirconium is predominantly in the +4 oxidation state in its compounds. Some less stable compounds, however, are known in which the oxidation state is +3. (The most important respect in which zirconium differs from titanium is that lower oxidation states are of minor importance.) The increased size of the atoms makes the oxides more basic and the aqueous chemistry somewhat more extensive and permits the attainment of coordination numbers 7 and, quite frequently, 8 in a number of zirconium compounds.

Various zirconium compounds have important applications in industry. Among these are zirconium dioxide (also called zirconia), ZrO2, a hard, white or yellow-brown solid with a high melting point—about 2,700° C (4,892° F). It is commonly used as a gem-diamond simulant, an abrasive, a refractory material, and a component of acid- and alkali-resistant glasses and of ceramics employed in fuel cells.

Other important industrial compounds of zirconium include the tetrachloride ZrCl4 and the sulfate Zr(SO4)2∙4H2O. Prepared by the chlorination of zirconium carbide or oxide, zirconium tetrachloride is used to produce organic zirconium compounds and as a catalyst in such reactions as the cracking of petroleum and polymerization of ethylene. Zirconium sulfate, produced by the action of sulfuric acid on zirconium hydoxide, Zr(OH)4, is useful as a lubricant, a chemical reagent, and in the tanning of white leather.

The Editors of Encyclopaedia Britannica This article was most recently revised and updated by Amy Tikkanen.

titanium

chemical element
Also known as: Ti
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titanium (Ti), chemical element, a silvery gray metal of Group 4 (IVb) of the periodic table. Titanium is a lightweight, high-strength, low-corrosion structural metal and is used in alloy form for parts in high-speed aircraft. A compound of titanium and oxygen was discovered (1791) by the English chemist and mineralogist William Gregor and independently rediscovered (1795) and named by the German chemist Martin Heinrich Klaproth.

Element Properties
atomic number22
atomic weight47.867
melting point1,660 °C (3,020 °F)
boiling point3,287 °C (5,949 °F)
density4.5 g/cm3 (20 °C)
oxidation states+2, +3, +4
electron configuration[Ar]3d24s2

Occurrence, properties, and uses

Titanium is widely distributed and constitutes 0.44 percent of Earth’s crust. The metal is found combined in practically all rocks, sand, clay, and other soils. It is also present in plants and animals, natural waters and deep-sea dredgings, and meteorites and stars. The two prime commercial minerals are ilmenite and rutile. The metal was isolated in pure form (1910) by the metallurgist Matthew A. Hunter by reducing titanium tetrachloride (TiCl4) with sodium in an airtight steel cylinder.

The preparation of pure titanium is difficult because of its reactivity. Titanium cannot be obtained by the common method of reducing the oxide with carbon because a very stable carbide is readily produced, and, moreover, the metal is quite reactive toward oxygen and nitrogen at elevated temperatures. Therefore, special processes have been devised that, after 1950, changed titanium from a laboratory curiosity to an important commercially produced structural metal. In the Kroll process, one of the ores, such as ilmenite (FeTiO3) or rutile (TiO2), is treated at red heat with carbon and chlorine to yield titanium tetrachloride, TiCl4, which is fractionally distilled to eliminate impurities such as ferric chloride, FeCl3. The TiCl4 is then reduced with molten magnesium at about 800 °C (1,500 °F) in an atmosphere of argon, and metallic titanium is produced as a spongy mass from which the excess of magnesium and magnesium chloride can be removed by volatilization at about 1,000 °C (1,800 °F). The sponge may then be fused in an atmosphere of argon or helium in an electric arc and be cast into ingots. On the laboratory scale, extremely pure titanium can be made by vaporizing the tetraiodide, TiI4, in very pure form and decomposing it on a hot wire in vacuum. (For treatment of the mining, recovery, and refining of titanium, see titanium processing. For comparative statistical data on titanium production, see mining.)

Pure titanium is ductile, about half as dense as iron and less than twice as dense as aluminum; it can be polished to a high lustre. The metal has a very low electrical and thermal conductivity and is paramagnetic (weakly attracted to a magnet). Two crystal structures exist: below 883 °C (1,621 °F), hexagonal close-packed (alpha); above 883 °C, body-centred cubic (beta). Natural titanium consists of five stable isotopes: titanium-46 (8.0 percent), titanium-47 (7.3 percent), titanium-48 (73.8 percent), titanium-49 (5.5 percent), and titanium-50 (5.4 percent).

Titanium is important as an alloying agent with most metals and some nonmetals. Some of these alloys have much higher tensile strengths than does titanium itself. Titanium has excellent corrosion-resistance in many environments because of the formation of a passive oxide surface film. No noticeable corrosion of the metal occurs despite exposure to seawater for more than three years. Titanium resembles other transition metals such as iron and nickel in being hard and refractory. Its combination of high strength, low density (it is quite light in comparison to other metals of similar mechanical and thermal properties), and excellent corrosion-resistance make it useful for many parts of aircraft, spacecraft, missiles, and ships. It also is used in prosthetic devices, because it does not react with fleshy tissue and bone. Titanium has also been utilized as a deoxidizer in steel and as an alloying addition in many steels to reduce grain size, in stainless steel to reduce carbon content, in aluminum to refine grain size, and in copper to produce hardening.

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Although at room temperatures titanium is resistant to tarnishing, at elevated temperatures it reacts with oxygen in the air. This is no detriment to the properties of titanium during forging or fabrication of its alloys; the oxide scale is removed after fabrication. In the liquid state, however, titanium is very reactive and reduces all known refractories.

Titanium is not attacked by mineral acids at room temperature or by hot aqueous alkali; it dissolves in hot hydrochloric acid, giving titanium species in the +3 oxidation state, and hot nitric acid converts it into a hydrous oxide that is rather insoluble in acid or base. The best solvents for the metal are hydrofluoric acid or other acids to which fluoride ions have been added; such mediums dissolve titanium and hold it in solution because of the formation of fluoro complexes.

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