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mercurythalliumlead
In

Tl

Uut
Appearance
silvery white
General properties
Name, symbol, number thallium, Tl, 81
Element category post-transition metal
Group, period, block 136, p
Standard atomic weight 204.3833g·mol−1
Electron configuration [Xe] 4f14 5d10 6s2 6p1
Electrons per shell 2, 8, 18, 32, 18, 3 (Image)
Physical properties
Phase solid
Density (near r.t.) 11.85 g·cm−3
Liquid density at m.p. 11.22 g·cm−3
Melting point 577 K, 304 °C, 579 °F
Boiling point 1746 K, 1473 °C, 2683 °F
Heat of fusion 4.14 kJ·mol−1
Heat of vaporization 165 kJ·mol−1
Specific heat capacity (25 °C) 26.32 J·mol−1·K−1
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 882 977 1097 1252 1461 1758
Atomic properties
Oxidation states 3, 1 (mildly basic oxide)
Electronegativity 1.62 (Pauling scale)
Ionization energies 1st: 589.4 kJ·mol−1
2nd: 1971 kJ·mol−1
3rd: 2878 kJ·mol−1
Atomic radius 170 pm
Covalent radius 170±8 pm
Van der Waals radius 196 pm
Miscellanea
Crystal structure hexagonal
Magnetic ordering diamagnetic[1]
Electrical resistivity (20 °C) 0.18 µΩ·m
Thermal conductivity (300 K) 46.1 W·m−1·K−1
Thermal expansion (25 °C) 29.9 µm·m−1·K−1
Speed of sound (thin rod) (20 °C) 818 m/s
Young's modulus 8 GPa
Shear modulus 2.8 GPa
Bulk modulus 43 GPa
Poisson ratio 0.45
Mohs hardness 1.2
Brinell hardness 26.4 MPa
CAS registry number 7440-28-0
Most stable isotopes
Main article: Isotopes of thallium
iso NA half-life DM DE (MeV) DP
203Tl 29.524% 203Tl is stable with 122 neutrons
204Tl syn 119 Ms
(3.78 y)
β 0.764 204Pb
ε 0.347 204Hg
205Tl 70.476% 205Tl is stable with 124 neutrons

Thallium (pronounced /ˈθæliəm/, THAL-ee-əm) is a chemical element with the symbol Tl and atomic number 81. This soft gray malleable poor metal resembles tin but discolors when exposed to air. Approximately 60-70% of thallium production is used in the electronics industry, and the rest is used in the pharmaceutical industry and in glass manufacturing.[2] It is also used in infrared detectors. Thallium is highly toxic and is used in rat poisons and insecticides, but its use has been cut back or eliminated in many countries. Because of its use for murder, thallium has gained the nicknames "The Poisoner's Poison" and "Inheritance Powder" (alongside arsenic).

Contents

Characteristics

Thallium is very soft and malleable and can be cut with a knife at room temperature. It has a metallic luster, but when exposed to air, it quickly tarnishes with a bluish-grey tinge that resembles lead. (It is preserved by keeping it under oil). A heavy layer of oxide builds up on thallium if left in air. In the presence of water, thallium hydroxide is formed.

History

Thallium (Greek θαλλός, thallos, meaning "a green shoot or twig")[3] was discovered by flame spectroscopy in 1862. The name comes from thallium's bright green spectral emission lines.

After the publication of the improved method of flame spectroscopy by Robert Bunsen and Gustav Kirchhoff[4] and the discovery of caesium and rubidium in the years 1859 to 1860 flame spectroscopy became an approved method to determine the composition of minerals and chemical products. William Crookes and Claude-Auguste Lamy both started to use the new method. William Crookes used it to make spectroscopic determinations for tellurium on selenium compounds deposited in the lead chamber of a sulfuric acid production plant near Tilkerode in the Harz mountains. He had obtained the samples for his research on selenium cyanide from August Hofmann years earlier.[5][6] By 1862 Crookes was able to isolate small quantities of the element and determine the properties of a few compounds.[7] Claude-Auguste Lamy used a similar spectrometer to Crookes' to determine the composition of a selenium-containing substance which was deposited during the production of sulfuric acid from pyrite. He also noticed the new green line in the spectra and concluded that a new element was present. Lamy had received this material from the sulfuric acid plant of his friend Fréd Kuhlmann and this by-product was available in large quantities. Lamy started to isolate the new element from that source.[8] The fact that Lamy was able to work ample quantities of thallium enabled him to determine the properties of several compounds and in addition he prepared a small ingot of metallic thallium which he prepared by remelting thallium he had obtained by electrolysis of thallium salts.

As both scientists discovered thallium independently and a large part of the work, especially the isolation of the metallic thallium was done by Lamy, Crookes tried to secure his priority on the work. Lamy was awarded a medal at the International Exhibition in London 1862: For the discovery of a new and abundant source of thallium and after heavy protest Crookes also received a medal: thallium, for the discovery of the new element. The controversy between both scientists continued through 1862 and 1863. Most of the discussion ended after Crookes was elected Fellow of the Royal Society in June 1863.[9][10]

Occurrence and production

Corroded thallium rod

Although the metal is reasonably abundant in the Earth's crust at a concentration estimated to be about 0.7 mg/kg, mostly in association with potassium minerals in clays, soils, and granites, it is not generally considered to be commercially recoverable from those forms. The major source of commercial thallium is the trace amounts found in copper, lead, zinc, and other sulfide ores.

Thallium is found in the minerals crookesite TlCu7Se4, hutchinsonite TlPbAs5S9, and lorandite TlAsS2. It also occurs as trace in pyrite and extracted as a by-product of roasting this ore for sulfuric acid production.[2] The metal can be obtained from the smelting of lead and zinc rich ores. Manganese nodules found on the ocean floor also contain thallium, but nodule extraction is prohibitively expensive and potentially environmentally destructive. In addition, several other thallium minerals, containing 16% to 60% thallium, occur in nature as sulfide or selenide complexes with antimony, arsenic, copper, lead, and silver, but are rare, and have no commercial importance as sources of this element. Thallium metal can also be obtained as a by-product in the production of sulfuric acid by roasting of pyrite.[2][11]

Isotopes

Thallium has 25 isotopes which have atomic masses that range from 184 to 210. 203Tl and 205Tl are the only stable isotopes, and 204Tl is the most stable radioisotope, with a half-life of 3.78 years.

202Tl (half life 12.23 days) can be made in a cyclotron,[12] while 204Tl (half life 3.78 years) is made by the neutron activation of stable thallium in a nuclear reactor.[13]

201Tl (half-life 73 hrs), decays by electron capture, emitting Hg x-rays (~ 70-80 keV), and photons of 135 and 167 keV in 10% total abundance; therefore it has good imaging characteristics without excessive patient radiation dose. It is the most popular isotope used for thallium nuclear cardiac stress tests.

Compounds

Fluorides: Thallium(I) fluoride (TlF), Thallium(III) fluoride (TlF3)
Chlorides: Thallium(I) chloride (TlCl), Thallium(II) chloride (TlCl2), Thallium(III) chloride (TlCl3)
Bromides: Thallium(I) bromide (TlBr), Thallium(II) bromide (Tl2Br4)
Iodides: Thallium triiodide (TlI), Thallium triiodide (TlI3)
Hydrides: none listed
Oxides: Thallium(I) oxide (Tl2O), Thallium(III) oxide (Tl2O3)
Sulfides: Thallium(I) sulfide Tl2S
Selenides: Thallium(I) selenide Tl2Se
Tellurides: none listed
Nitrides: none listed

Applications

The odorless and tasteless thallium sulfate was once widely used as rat poison and ant killer. Since 1975, this use in the United States and many other countries is prohibited due to safety concerns.[2] Other uses:

The saturated solution of equal parts of thallium(I) formate (Tl(CHO2)) and thallium(I) malonate (Tl(C3H3O4)) in water is known as Clerici solution. It is a mobile odorless liquid whose color changes from yellowish to clear upon reducing the concentration of the thallium salts. With the density of 4.25 g/cm3 at 20 °C, Clerici solution is one of the heaviest aqueous solutions known. It was used in the 20th century for measuring density of minerals by the flotation method, but the use is discontinued due to the high toxicity and corrosiveness of the solution.[20][21]

Research activity with thallium is ongoing to develop high-temperature superconducting materials for such applications as magnetic resonance imaging, storage of magnetic energy, magnetic propulsion, and electric power generation and transmission. After the discovery of the first thallium barium calcium copper oxide superconductor in 1988 the research in applications started.[22]

Toxicity

Skull and crossbones.svg

Thallium and its compounds are extremely toxic, and should be handled with great care. Contact with skin is dangerous, and adequate ventilation should be provided when melting this metal. Thallium(I) compounds have a high aqueous solubility and are readily absorbed through the skin. Exposure to them should not exceed 0.1 mg per m² of skin in an 8-hour time-weighted average (40-hour work week). Thallium is a suspected human carcinogen.[23]

Treatment and internal decontamination

One of the main methods of removing thallium (both radioactive and normal) from humans is to use Prussian blue, which is a solid ion exchange material which absorbs thallium and releases potassium. Up to 20 g per day of Prussian blue is fed by mouth to the person, and it passes through their digestive system and comes out in the stool. Hemodialysis and hemoperfusion are also used to remove thallium from the blood serum. At later stage of the treatment additional potassium is used to mobilize thallium from the tissue.[24][25]

Thallium pollution

According to the United States Environmental Protection Agency (EPA), man-made sources of thallium pollution include gaseous emission of cement factories, coal burning power plants, and metal sewers. The main source of elevated thallium concentrations in water is the leaching of thallium from ore processing operations.[26]

References

  1. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.
  2. ^ a b c d e f "Chemical fact sheet — Thallium". Spectrum Laboratories. April 2001. http://www.speclab.com/elements/thallium.htm. Retrieved 2008-02-02.  
  3. ^ Liddell & Scott, A Greek-English Lexicon, sub θαλλος
  4. ^ G. Kirchhoff, R. Bunsen (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. doi:10.1002/andp.18611890702.  
  5. ^ Crookes, William (1862 - 1863). "Preliminary Researches on Thallium". Proceedings of the Royal Society of London, 12: 150–159. doi:10.1098/rspl.1862.0030. http://www.jstor.org/stable/112218.  
  6. ^ Crookes, William (1863). "On Thallium". Philosophical Transactions of the Royal Society of London, 153: 173–192. doi:10.1098/rstl.1863.0009. http://www.jstor.org/stable/108794.  
  7. ^ DeKosky, Robert K. (1973). "Spectroscopy and the Elements in the Late Nineteenth Century: The Work of Sir William Crookes". The British Journal for the History of Science 6 (4): 400–423. doi:10.1017/S0007087400012553. http://www.jstor.org/stable/4025503.  
  8. ^ Lamy, Claude-Auguste (1862). "De l'existencè d'un nouveau métal, le thallium". Comptes Rendus: 1255–. http://gallica2.bnf.fr/ark:/12148/bpt6k30115.image.r=Comptes+Rendus+Hebdomadaires.f1254.langFR.  
  9. ^ James, Frank A. J. L. (1984). "Of 'Medals and Muddles' the Context of the Discovery of Thallium: William Crookes's Early". Notes and Records of the Royal Society of London 39 (1): 65–90. doi:10.1098/rsnr.1984.0005. http://www.jstor.org/stable/531576.  
  10. ^ Emsley, John (2006). "Thallium". The Elements of Murder: A History of Poison. Oxford University Press. pp. 326–327. ISBN 9780192806000. http://books.google.de/books?id=BACSR7TXWhoC.  
  11. ^ Downs, Anthony John (1993). "Chemistry of Aluminium, Gallium, Indium, and Thallium". Springer. pp. 89 and 106. http://books.google.com/books?id=v-04Kn758yIC.  
  12. ^ Thallium Research from Department of Energy
  13. ^ Manual for reactor produced radioisotopes from the International Atomic Energy Agency
  14. ^ a b c d e f C. R. Hammond. The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304857.  
  15. ^ Nayer, P. S, Hamilton, O. (1977). "Thallium selenide infrared detector". Appl. Opt. 16: 2942. doi:10.1364/AO.16.002942. http://adsabs.harvard.edu/abs/1977ApOpt..16.2942N.  
  16. ^ Thallium Test from Walter Reed Army Medical Center
  17. ^ Thallium Stress Test from the American Heart Association
  18. ^ M. C., Lagunas-Solar; Little, F. E.; Goodart, C. D. (1982). Abstract "An integrally shielded transportable generator system for thallium-201 production". International Journal of Applied Radiation Isotopes 33 (12): 1439–1443. doi:10.1016/0020-708X(82)90183-1. http://www.medscape.com/medline/abstract/7169272 Abstract.  
  19. ^ Thallium-201 production from Harvard Medical School's Joint Program in Nuclear Medicine
  20. ^ R. H. Jahns (1939). Clerici solution for the specific gravity determination of small mineral grains. 24. p. 116. http://www.minsocam.org/ammin/AM24/AM24_116.pdf.  
  21. ^ Peter G. Read (1999). Gemmology. Butterworth-Heinemann. pp. 63-64. ISBN 0750644117. http://books.google.com/books?id=tfXa13uWiRIC&pg=PA63&lpg=PA63.  
  22. ^ Sheng, Z. Z.; Hermann A. M. (1988). "Bulk superconductivity at 120 K in the Tl–Ca/Ba–Cu–O system". Nature 332: 138–139. doi:10.1038/332138a0.  
  23. ^ "Biology of Thallium". webelemnts. http://www.webelements.com/webelements/elements/text/Tl/biol.html. Retrieved 2008-11-11.  
  24. ^ Prussian blue fact sheet from the Centers for Disease Control and Prevention
  25. ^ Malbrain, Manu L. N. G.; Lambrecht, Guy L. Y.; Zandijk, Erik; Demedts, Paul A.; Neels, Hugo M.; Lambert, Willy; De Leenheer, André P.; Lins, Robert L.; Daelemans, Ronny; (1997). "= Treatment of Severe Thallium Intoxication". Clinical Toxicology 35 (1): 97–100. doi:10.3109/15563659709001173.  
  26. ^ "Factsheet on: Thallium". http://www.epa.gov/safewater/pdfs/factsheets/ioc/thallium.pdf. Retrieved 2009-09-15.  

External links


1911 encyclopedia

Up to date as of January 14, 2010

From LoveToKnow 1911

Medical warning!
This article is from the 1911 Encyclopaedia Britannica. Medical science has made many leaps forward since it has been written. This is not a site for medical advice, when you need information on a medical condition, consult a professional instead.

THALLIUM [[[symbol]] Tl, atomic weight 204.0 (0 = 16)], a metallic chemical element. It was discovered in 1861 by Sir William Crookes, who, during a spectroscopic examination of the flue-dust produced in the roasting of seleniferous pyrites occurring at Tilkerode in the Harz, observed a green line foreign to all then known spectra. He concluded that the mineral contained a new element, to which he gave the name of thallium, from 9aXX6, a green twig. Crookes presumed that his thallium was something of the order of sulphur, selenium or tellurium; but Lamy, who anticipated him in isolating the new element, found it to be a metal. Our knowledge of the chemistry of thallium is based chiefly upon the labours of Crookes.

The chemical character of thallium presents striking peculiarities. Dumas once called it the "ornithorhynchus paradoxus of metals." As an elementary substance, it is very similar in its physical properties to lead; it resembles lead chemically inasmuch as it forms an almost insoluble chloride and an insoluble iodide. But the hydroxide of thallium, in most of its properties, comes very close to the alkali metals; it is strongly basic, forms an insoluble chloroplatinate, and an alum strikingly similar to the corresponding potassium compounds. Yet, unlike potassium or lead, it forms a feebly basic sesquioxide similar to manganic oxide, Mn203.

Traces of thallium exist in many kinds of pyrites, as used for vitriol-making. The only known mineral of which it forms an essential component is the rare mineral crookesite of Skrikerum, Smaland, Sweden, which, according to Nordenskidld, contains 33.3 per cent. of selenium, 45.8 per cent. of copper, 3.7 per cent. of silver, and 17.2 per cent. of thallium. The best raw materials for the preparation of thallium are the flue-dusts produced industrially in the roasting of thalliferous pyrites and the "chamber muds" accumulating in vitriol-chambers wrought with such pyrites; in both it is frequently associated with selenium. The flue-dust from the pyrites of Theux, near Spa (Belgium), according to BSttcher, contains o 5 to o 75 per cent. of thallium; that of the pyrites of Meggen, according to Carstanjen, as much as 3.5 per cent.; while that of the pyrites of Ruhrort yielded 1 per cent. of the pure chloride to Gunning.

For the extraction of the metal from chamber mud, the latter is boiled with water, which extracts the thallium as the sulphate. From the filtered solution the thallium is precipitated as the chloride by addition of hydrochloric acid, along, in general, with more or less of lead chloride. The mixed chlorides are boiled down to dryness with sulphuric acid to convert them into sulphates, which are then separated by boiling water, which dissolves only the thallium salt. From the filtered solution the thallium is recovered, as such, by means of pure metallic zinc, or by electrolysis. The (approximately pure) metallic sponge obtained is washed, made compact by compression, fused in a porcelain crucible in an atmosphere of hydrogen, and cast into sticks.

Metallic thallium is bluish white; it is extremely soft and almost devoid of tenacity and elasticity. Its specific gravity is 11.86. It fuses at 290° C.; at a white heat it boils and can be distilled in hydrogen gas. Its vapour density at 1728° corresponds to the molecule TI 2. Its salts colour the Bunsen flame a bright green. When heated in air it is readily oxidized, with the formation of a reddish or violet vapour. When exposed to the air it becomes quickly covered with a film of oxide; the tarnished metal when plunged into water reassumes its metallic lustre, the oxide film being quickly dissolved. When kept in contact with water and air it is gradually converted into hydroxide, T10H. It decomposes water at a red heat, liberating hydrogen and being itself converted into the hydrate. It is readily soluble in nitric and sulphuric acids, but less so in hydrochloric.

Thallium forms two series of salts: thallous, in which the metal is monovalent; and thallic, in which it is trivalent. In the thallous series many analogies with lead compounds are observed; in the thallic some resemblance to aluminium and gold.

Thallous hydroxide, T10H, is most conveniently prepared by decomposing the solution of the sulphate with baryta water. It crystallizes from its solution in long yellow needles, T10H or T10H-+H 2 0, which dissolve readily in water, forming an intensely alkaline solution, which acts as a caustic, and like it greedily absorbs carbonic acid from the atmosphere. Unlike the alkalis, it readily loses its water at too° C. and even at the ordinary temperature, to form the oxide T1 2 0, which is black or black-violet.

Thallic oxide, T10 or T1202, was obtained by 0. Rabe (Abst. J.C.S., 1907, ii. 769) by acting with hydrogen peroxide on an alkaline solution of thallous sulphate at low temperatures, an initial red precipitate rapidly changing into a bluish-black compound. It melts at 720° and decomposes rapidly above 800°, giving oxygen and thallous oxide. Thallous chloride, T1C1, is readily obtained from the solution of any thallous salt, by the addition of hydrochloric acid, as a white precipitate similar in appearance to silver chloride, like which it turns violet in the light and fuses below redness into a (yellow) liquid which freezes into a horn-like flexible mass. It is also formed when the metal is burnt in chlorine. The specific gravity of this "horn" thallium is 7.02. One part of the precipitated chloride dissolves at o° C. in 500 parts of water, and in 70 parts at loo° C. It is less soluble in dilute hydrochloric acid. Carbonate of soda solution dissolves it pretty freely. Thallous iodide, T11, is obtained as a yellow precipitate, which requires 16,000 parts of cold water for its solution, by the addition of potassium iodide to a solution of a thallous salt, or by the direct union of its components. The yellow crystals melt at 190°, and when cooled down assume a red colour, which changes to the original yellow on standing. Thallous bromide, TIBr, is a light yellow crystalline powder; it is formed analogously to the chloride. Thallous fluoride, T1F, forms white glistening octahedra; it is obtained by crystallizing a solution of the carbonate in hydrofluoric acid. It resembles potassium fluoride in forming an acid salt, T1HF 2. Thallous chloroplatinate, T1 2 PtC1 6, readily obtainable from thallous salt solutions by addition of platinum chloride, is a yellow precipitate soluble in no less than 15,600 parts of cold water. Thallous Perchlorate, T1C10 4, and periodate, Tl10 4, are interesting inasmuch as they are isomorphous with the corresponding potassium salts. Other instances of the isomorphism of thallous with potassium salts are the nitrates, phosphates, hydrazoates, sulphates, chromates, selenates, and the analogously constituted double salts, and also the oxalates, racemates and picrates. Thallous carbonate, T1 2 CO 31 more nearly resembles the lithium compound than any other ordinary carbonate. It is produced by the exposure of thallous hydrate to carbon dioxide, and therefore is obtained when the moist metal is exposed to the air. It forms resplendent monoclinic prisms, soluble in water. Thallous sulphate, T1 2 SO 4, forms rhombic prisms, soluble in water, which melt at a red heat with decomposition, sulphur dioxide being evolved. It unites with sulphuric acid giving an acid salt, T1HSO 4.3H 2 O, and with aluminium, chromium and iron sulphates to form an "alum." It also forms double salts of the type T1 2 SO 4 (Mg,Fe,ZnSO 4) 6H 2 O. Thallous sulphide, T1 2 S, is obtained as a black precipitate by passing sulphuretted hydrogen into a thallous solution. It is insoluble in water and in the alkalis, but readily dissolves in the mineral acids. On thallium sulphides see H. Pelabon, Abst. J.C.S., 1907, ii. 770. Thallous nitrate, T1NO 31 is obtained as white, rhombic prisms by crystallizing a solution of the metal, oxide, carbonate, &c., in nitric acid. Various thallous phosphates are known. The normal salt, T1 3 PO 4, is soluble in 200 parts of water, and may be obtained by precipitation. On thallous salts see W. Stortenbeker, Abst. J.C.S., 1907, ii. 770. Thallic oxide, T1203, is obtained as a dark reddish powder, insoluble in water and alkalis, by plunging molten thallium into oxygen, or by electrolysing water, using a thallium anode. Thallic hydroxide, TI(OH) 31 is obtained as a brown precipitate by adding a hot solution of thallous chloride in sodium carbonate to a solution of sodium hypochlorite. On drying it has the composition TIO(OH). Hydrochloric acid gives thallous chloride and chlorine; sulphuric acid gives off oxygen; and on heating it first gives the trioxide and afterwards the monoxide. The hydroxide is obtained as brown hexagonal plates by fusing thallic oxide with potash to which a little water has been added. Thallic chloride, T1C1 3, is obtained by treating the monochloride with chlorine under water; evaporation in a vacuum gives colourless deliquescent crystals of T1C1,.H20. By heating the metal or thallous chloride in chlorine, T1C1 T1C1 3 is obtained, which on further heating gives3TlCI.T1C13. as a yellowish brown mass. The chloride when anhydrous is a crystalline mass which melts at 24°. It forms several double salts, e.g. with hydrochloric acid and the alkaline chlorides, and also with nitrosyl chloride. The chlorine is not completely precipitated by silver nitrate in nitric acid solution, the ionization apparently not proceeding to all the chlorine atoms. Thallic iodide, T11 3, is interesting on account of its isomorphism with rubidium and caesium tri-iodides, a resemblance which suggests the formula T11 (12) for the salt, in which the metal is obviously monovalent. On the halogen compounds see V. Thomas, Abst. J.C.S. (1907), ii. 547. Thallic sulphate, T1 2 (SO 4) 3.7H 2 O, and thallic nitrate, Tl(NO 3) 3.8H 2 0, are obtained as colourless crystals on the evaporation of a solution of the oxide in the corresponding acid. The sulphate decomposes into sulphuric acid and the trioxide on warming with water, and differs from aluminium sulphate in not forming alums.

Analysis

All thallium compounds volatile or liable to dissociation at the temperature of the flame of a Bunsen lamp impart to such flame an intense green colour. The spectrum contains a bright green of wave-length 5351. From solutions containing it as thallous salt the metal is easily precipitated as chloride, iodide, or chloroplatinate by the corresponding reagents. Sulphuretted hydrogen, in the presence of free mineral acid, gives no precipitate; sulphide of ammonium, from neutral solutions, precipitates T12S as a dark brown or black precipitate, insoluble in excess of reagent. Thallic salts are easily reduced to thallous by means of solution of sulphurous acid, and thus rendered amenable to the above reactions. The chloroplatinate serves for the quantitative estimation. L. F. Hawley employs sodium thiostannate which precipitates thallium as T1 2 SnS 4, insoluble in water, and which may be dried on a Gooch filter at 105°. It may be noted that all thallium compounds are poisonous.

The atomic weight of thallium was determined very carefully by Crookes, who found T1=204.2 (0= 16); this figure was confirmed by Lepierre in 1893.


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Wiktionary

Up to date as of January 15, 2010

Definition from Wiktionary, a free dictionary

See also thallium

German

Chemical Element: Tl (atomical number 81)

Noun

Thallium n

  1. thallium

Simple English

[[File:|thumb|Thallium]] Thallium is a chemical element. It has the chemical symbol Tl. It has the atomic number 81. Its atomic mass is 204.4. It is found in Group 13 of the periodic table. Thallium is a gray metal that is very toxic.

Contents

Properties

Physical properties

Thallium is a soft, malleable, grayish post-transition metal. It can be cut with a knife at room temperature. It melts at a low temperature, 304°C. This is typical of a post-transition metal. Thallium has 25 isotopes and two stable (nonradioactive) ones. It is extremely toxic.

Chemical properties

File:Thallium rod
A corroded thallium rod

Thallium is a moderately reactive metal. It corrodes easily in air with a blue-gray color that is similar to lead. If it is kept in air for a long time, a large amount of thallium(I) oxide will build up. It corrodes in the presence of water to make the hydroxide. It burns with a greenish flame. It reacts with most acids.

Chemical compounds

Thallium makes chemical compounds in two oxidation states: +1 and +3. The +1 state is more common and less reactive. Its chemical compounds are very similar to potassium or silver compounds. It makes a hydroxide that in a strong base when dissolved in water. Most other transition metal and post-trnansition metal hydroxides do not dissolve in water. This reacts with carbon dioxide to make thallium(I) carbonate, which is also water-soluble and very heavy. It is the only heavy metal carbonate that can dissolve in water. Other compounds are similar to silver compounds. Thallium(I) bromide turns yellow when exposed to light, similar to silver(I) bromide. Thallium(I) sulfide is black, similar to silver(I) sulfide. The +3 state compounds are oxidizing agents. The black oxide, thallium(III) oxide and the hydroxide, thallium(III) hydroxide, are the only stable +3 compounds. They break down to oxygen and thallium(I) oxide when heated. Thallium and its compounds are rare because they are toxic and polluting.

File:Thallium(I)
Thallium(I) chloride

History

Thallium was found by spectroscopy in 1861 by a bright green line in its spectrum. The main use for thallium, rat poison, was banned in many countries in the 1970's. Thallium was used to poison people, similar to the more popular arsenic.[1]

Occurrence

[[File:|thumb|A mineral that has thallium in it]] Thallium is found most in certain clays and granites. It cannot be gotten easily from these, though. Thallium is normally gotten from the waste after other ores like galena are processed.[2] Hutchinsonite is another mineral that has thallium in it.

Preparation

When lead and zinc are taken from their ores, many impurities are left behind. Sulfuric acid is used to dissolve the thallium from it as thallium(I) sulfate. Then the thallium(I) sulfate is electrolyzed to make thallium metal.

Uses

It is used in rat poisons and insecticides. The use of thallium as a poison has been reduced or banned in many countries because these countries think that thallium might cause cancers. It is also used in infrared detectors. It has been used in some murders. Like arsenic, the use of thallium in murders has given it the name "inheritance powder". Thallium compounds are used in glass for infrared light. Thallium was also used to kill skin infections, but it is too toxic to be used for that now. A superconductor that can work at higher temperatures than normal ones do uses thallium. A radioactive thallium isotope was used for nuclear scans. An alloy of thallium and mercury has a low freezing temperature and is a liquid. A very dense solution of a thallium compound was used to test minerals for specific gravity, but it is too toxic for use.

Safety

Thallium is very toxic. It can be absorbed through skin. Many of its salts easily dissolve. Some are colorless, tasteless, and odorless, but are very toxic. Some think that it is a carcinogen. Thallium can be a pollutant if the thallium waste from metal processing is washed away.

References

  1. Hasan, Heather (2009). The Boron Elements: Boron, Aluminum, Gallium, Indium, Thallium. Rosen Publishing Group. p. 14. ISBN 9781435853331. 
  2. Peter, A; Viraraghavan, T (2005). [Expression error: Unexpected < operator "Thallium: a review of public health and environmental concerns"]. Environment International 31 (4): 493–501. doi:10.1016/j.envint.2004.09.003. PMID 15788190. 








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