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Titan er det 22. grundstof i det periodiske system, og har det kemiske symbol N.. Under normale temperatur- og trykforhold optræder dette overgangsmetal som et sølvskinnende metal der er lige så stærkt som stål, men blot har 60 procent af stålets densitet. Titan modstår korrosion bl.a. overfor havvand, kongevand og chlor.

Titan kan bl.a. danne legeringer med jern, aluminium, vanadium, og molybdenum og danner stærke legeringer som bruges inden til rumfart (jetmotorer, missiler og rumfartøjer), militæret, industrielle processer (kemikalier og petrokemikalier, desalination plants, papirmasse og papir), biler, medicinske proteser, ortopædiske implantater, medicinske instumenter, sportsudstyr, smykker, mobiltelefoner og mange andre applikationer.[1]

Titan blev opdaget i England af William Gregor i 1791 og navngivet af Martin Heinrich Klaproth efter Titanerne fra Græsk mytologi.

Grundstoffet forekommer within a number of mineral deposits, principally rutil and ilmenit, which are widely distributed in the Jordens skorpe og lithosfæren, and it is found in almost all living things, rocks, water bodies, and soils.[1] The metal is extracted from its principal mineral ores via the Kroll process[2] or the Hunter process. Its most common compound, titanium dioxide, is a popular photocatalyst and is used in the manufacture of white pigments.[3] Other compounds include titanium tetrachloride (TiCl4), a component of smoke screens and catalysts; and titanium trichloride (TiCl3), which is used as a catalyst in the production of polypropylene).[1]

The two most useful properties of the metal form are corrosion resistance and the highest strength-to-weight ratio of any metal.[4] In its unalloyed condition, titanium is as strong as some steels, but 45% lighter.[5] There are two allotropic forms[6] and five naturally occurring isotopes of this element; 46Ti through 50Ti, with 48Ti being the most abundant (73.8%).[7] Titanium's properties are chemically and physically similar to zirconium.

CharacteristicsRediger

PhysicalRediger

A metallic element, titanium is recognized for its high strength-to-weight ratio.[6] It is a strong metal with low density that is quite ductile (especially in an oxygen-free environment),[1] lustrous, and metallic-white in color.[8] The relatively high melting point (over 1,649 °C or 3,000 °F) makes it useful as a refractory metal.

Commercial (99.2% pure) grades of titanium have ultimate tensile strength of about 63,000 psi (434 MPa), equal to that of common, low-grade steel alloys, but are 45% lighter.[5] Titanium is 60% more dense than aluminium, but more than twice as strong[5] as the most commonly used 6061-T6 aluminium alloy. Certain titanium alloys (e.g., Beta C) achieve tensile strengths of over 200.000 psi (1.400 MPa).[9] However, titanium loses strength when heated above 430 °C (806 °F).[10]

It is fairly hard although not as hard as some grades of heat-treated steel, non-magnetic and a poor conductor of heat and electricity. Machining requires precautions, as the material will soften and gall if sharp tools and proper cooling methods are not used. Like those made from steel, titanium structures have a fatigue limit which guarantees longevity in some applications.[8]

The metal is a dimorphic allotrope whose hexagonal alpha form changes into a body-centered cubic (lattice) β form at 882 °C (1.620 °F).[10] The specific heat of the alpha form increases dramatically as it is heated to this transition temperature but then falls and remains fairly constant for the β form regardless of temperature.[10] Similar to zirconium and hafnium, an additional omega phase exists, which is thermodynamically stable at high pressures, but is metastable at ambient pressures. This phase is usually hexagonal (ideal) or trigonal (distorted) and can be viewed as being due to a soft longitudinal acoustic phonon of the β phase causing collapse of (111) planes of atoms.[11]

KemiskRediger

Titans mest bemærkelsesværdige egenskab er dets modstantsdygtighed overfor korrosion - næsten ligeså resitent som platin - og modstår angreb fra syrer, fugtig chlor i vand, men er opløselig i koncentrerede syrer.[12]

 
The Pourbaix diagram for titanium in pure water, perchloric acid or sodium hydroxide[13]

Pourbaix diagramet viser at titan termodynamisk er et meget reaktivt metal, og alligevel reagerer det langsomt med både vand og luften. Dette skyldes at metallet dannee et pasivt og beskyttende oxidlag (hvilket øger korosionsevnen)[1] Oxidlaget når det dannes er først 1–2 nm tykt, men det fortsætter langsomt med at vokse, indtil det har en tykkelse på 25 nm efter fire år.[14]

Titan brænder i luft når det opvarmes til 1200 °C og i ren oxygen når det opvarmes til 610 °C eller højere, hvorved dannes titandioxid, TiO2[6] As a result, the metal cannot be melted in open air as it burns before the melting point is reached, so melting is only possible in an inert atmosphere or in a vacuum. It is also one of the few elements that burns in pure nitrogen gas (it burns at 800 °C or 1,472 °F and forms titanium nitride, which causes embrittlement).[15] Titanium is resistant to dilute sulfuric acid and hydrochloric acid, along with chlorine gas, chloride solutions, and most organic acids.[2] It is paramagnetic (weakly attracted to magnets) and has fairly low electrical and thermal conductivity.[1]

Experiments have shown that natural titanium becomes radioactive after it is bombarded with deuterons, emitting mainly positrons and hard gamma rays.[2] When it is red hot the metal combines with oxygen, and when it reaches 550 °C (1.022 °F) it combines with chlorine.[2] It also reacts with the other halogens and absorbs hydrogen.[3]

ForbindelserRediger

 
TiN coated drill bit

Oxidationstrin +4 dominerer titans kemi,[16] men forbindelser i oxidationstrin +3 er også almindelige.[17] Pga. det høje oxidationstrin har bindingerne i titans forbindelser en høj grad af kovalent karakter.

Rubiner har deres asterism fra urenheder af titandioxid, TiO2.[14] Titanater er forbindelser med titandioxid. Bariumtitanat har piezoelektriske egenskaber, som gør det muligt at anvende det til en transducer in the  interconversion of lyd og elektricitet.[6] Estere af titan dannes ved reaktion mellem alkohol og titantetrachlorid og bruges til at waterproof fabrics.[6]

Titannitrid (TiN) har en hårdhed svarende til safir og carborundum (9,0 på Mohs Scale),[18] og benyttes derfor ofte til skæreværktøjer, som fx such as bor.[19] It also finds use as a gold-colored decorative finish, and as a barrier metal in semiconductor fabrication.[20]

Titantetrachlorid (titan(IV)chlorid, TiCl4 er en farveløs væske som er et mellemprodukt i syntesen af titandioxid, TiO2, til maling.[21] Derudover bruges TiCl4 i organisk kemi som Lewis syre, for eksempel til Mukaiyama aldolkondensation.[22] Titan danner også et lavere chlorid, titan(III)chlorid (TiCl3), som er et reduceringsmiddel.[23]

Titanocendichlorid er en vigtig katalysator til dannelse af carbon-carbon bindinger. Titanisopropoxid benyttes i Sharpless epoxidering. Andre forbindelser inkluderer titaniumbromid (used in metallurgy, superalloys, and high-temperature electrical wiring and coatings) and titanium carbide (found in high-temperature cutting tools and coatings).[3]

OccurrenceRediger

Produktion af titandioxid i 2003, i tusind ton.[24]
Producent Produktion % af ialt
  Australien 1291,0 30,6
  Sydafrika 850,0 20,1
  Canada 767,0 18,2
  Norge 382,9 9,1
  Ukraine 357,0 8,5
Øvrige lande 573,1 13,6
Totalt 4221,0 100,0
Pga. afrunding summer værdierne ikke til 100%.

Titani er altid bundet til andre grundstoffer i naturen. Det er det niende mest udbredte grundstof i Jordens skorpe (0,63% i masseprocent)[25] og det syvende mest udbredte metal. It is present in most igneous rocks and in sediments derived from them (as well as in living things and natural bodies of water).[1][2] Of the 801 types of igneous rocks analyzed by the United States Geological Survey, 784 contained titanium.[25] Its proportion in soils is approximately 0.5 to 1.5%.[25]

It is widely distributed and occurs primarily in the minerals anatase, brookite, ilmenite, perovskite, rutile, titanite (sphene), as well in many iron ores.[14] Of these minerals, only rutile and ilmenite have any economic importance, yet even they are difficult to find in high concentrations. Significant titanium-bearing ilmenite deposits exist in western Australia, Canada, China, India, New Zealand, Norway, and Ukraine.[14] Large quantities of rutile are also mined in North America and South Africa and help contribute to the annual production of 90,000 tonnes of the metal and 4.3 million tonnes of titanium dioxide.[14] Total reserves of titanium are estimated to exceed 600 million tonnes.[14]

Titanium is contained in meteorites and has been detected in the sun and in M-type stars;[2] the coolest type of star with a surface temperature of 3.200 °C (5.790 °F).[26] Rocks brought back from the moon during the Apollo 17 mission are composed of 12.1% TiO2.[2] It is also found in coal ash, plants, and even the human body.

IsotoperRediger

Titan har fem naturligt forekommende og stabile isotoper: 46Ti, 47Ti, 48Ti, 49Ti, og 50Ti, hvor 48Ti er den mest udbredte (73,8% natural abundance). Derudover er elleve radioaktive isotoper blevet karakteriseret, og blandt disse er 44Ti med en halveringstid på 63 år, 45Ti med en halveringstid på 184,8 minutter, 51Ti med en halveringstid 5,76 minutter, samt 52Ti med en halveringstid på 1,7 minutter. De resterende radioaktive isotoper har halveringstid som er mindre end 33 sekunder, og flertallet af disse har en halveringstid på mindre end ét sekund.[7]

Atomvægten af titans isotoper går fra 39,99 u 40Ti til 57,966 u (58Ti). The primary decay mode before the most abundant stable isotope, 48Ti, is electron capture and the primary mode after is beta emission. The primary decay products before 48Ti are element 21 (scandium) isotopes and the primary products after are element 23 (vanadium) isotopes.[7]


ReferencerRediger

  1. ^ a b c d e f g "Titanium". Encyclopædia Britannica. (2006). Hentet 2006-12-29. 
  2. ^ a b c d e f g Lide, D. R., (red.) (2005). CRC Handbook of Chemistry and Physics (86th udgave). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5. 
  3. ^ a b c Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements: A Reference Guide (2nd edition). Westport, CT: Greenwood Press. ISBN 0313334382. 
  4. ^ Matthew J. Donachie, Jr. (1988). TITANIUM: A Technical Guide. Metals Park, OH: ASM International. s. 11. ISBN 0871703092. 
  5. ^ a b c Barksdale 1968, s. 738
  6. ^ a b c d e "Titanium". Columbia Encyclopedia (6th edition). (2000–2006). New York: Columbia University Press. ISBN 0-7876-5015-3. 
  7. ^ a b c Barbalace, Kenneth L. (2006). "Periodic Table of Elements: Ti - Titanium". Hentet 2006-12-26. 
  8. ^ a b Fodnotefejl: Ugyldigt <ref>-tag; ingen tekst er angivet for referencer med navnet Stwertka1998
  9. ^ Matthew J. Donachie, Jr. (1988). Titanium: A Technical Guide. Metals Park, OH: ASM International. Appendix J, Table J.2. ISBN 0871703092. 
  10. ^ a b c Barksdale 1968, s. 734
  11. ^ Sikka, S. K.; Vohra, Y. K., Chidambaram, R. (1982). "Omega phase in materials". Progress in Materials Science. 27: 245–310. doi:10.1016/0079-6425(82)90002-0. 
  12. ^ Casillas, N.; Charlebois, S.; Smyrl, W. H.; White, H. S. (1994). "Pitting Corrosion of Titanium". J. Electrochem. Soc. 141 (3): 636–642. doi:10.1149/1.2054783. 
  13. ^ Ignasi Puigdomenech, Hydra/Medusa Chemical Equilibrium Database and Plotting Software (2004) KTH Royal Institute of Technology, freely downloadable software at [1]
  14. ^ a b c d e f Emsley 2001, s. 453
  15. ^ "Titanium". Microsoft Encarta. (2005). Hentet 2006-12-29. 
  16. ^ Greenwood 1997, s. 958
  17. ^ Greenwood 1997, s. 970
  18. ^ Schubert, E.F. "The hardness scale introduced by Friederich Mohs" (PDF). 
  19. ^ Truini, Joseph. "Drill Bits". Popular Mechanics. Hearst Magazines. 165 (5): 91. ISSN 0032-4558. 
  20. ^ Baliga, B. Jayant (2005). Silicon carbide power devices. World Scientific. s. 91. ISBN 9812566058. 
  21. ^ Johnson, Richard W. (1998). The Handbook of Fluid Dynamics. Springer. s. 38–21. ISBN 3540646124. 
  22. ^ Coates, Robert M.; Paquette, Leo A. (2000). Handbook of Reagents for Organic Synthesis. John Wiley and Sons. s. 93. ISBN 0470856254. 
  23. ^ Grimmett, M. Ross (1997). Imidazole and benzimidazole synthesis. Academic Press. s. 155. ISBN 0123031907. 
  24. ^ Cordellier, Serge; Didiot, Béatrice (2004). L'état du monde 2005: annuaire économique géopolitique mondial. Paris: La Découverte. 
  25. ^ a b c Fodnotefejl: Ugyldigt <ref>-tag; ingen tekst er angivet for referencer med navnet Barksdale1968p732
  26. ^ Emsley 2001, s. 451