|Boiling Point: 1746°K, 1473°C, 2683°F
Melting Point: 577°K, 304°C, 579°F
Electrons Energy Level: 2, 8, 18, 32, 18, 3
Isotopes: 30 + 2 Stable
Heat of Vaporization: 164.1 kJ/mol
Heat of Fusion: 4.142 kJ/mol
Density: 11.85 g/cm3 @ 300°K
Specific Heat: 0.13 J/g°K
Atomic Radius: 2.08Å
Ionic Radius: 1.5Å
Electronegativity: 2.04 (Pauling); 1.44 (Allrod Rochow)
Vapor Pressure: 5.33E-06 Pa @ 304°C
Crookes discovered thallium in 1861, positively identifying it by a bright green line in its spectrum (hence the name, which is from the
Greek, thallos, for "green twig") while he
was making spectroscopic determinations for Tellurium on residues from a Sulfuric Acid
plant. Rather than seeing the yellow spectral lines produced by Tellurium, he
observed a bright green line that no one had ever
seen before. In 1862 Crookes and Claude-Auguste Lamy isolated the metal independently of
each other. Although in appearance thallium resembles lead, it does not have the
corrosion resistance of Lead and so has few commercial applications.
This soft gray malleable poor metal resembles Tin but discolors when exposed to air. Thallium is highly toxic and is used in rat poisons and insecticides, Thallium Sulfate, but since it might also cause cancer (although the United States EPA does not class it as carcinogen), this use has been cut back or eliminated in many countries. It is also used in infrared detectors. It has even been used in some murders, earning the nicknames "The Poisoner's Poison" and "Inheritance powder" (alongside Arsenic).
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.
1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p1
This metal is very soft and malleable and can be cut with a knife. When it is first exposed to air, thallium has a metallic luster but quickly tarnishes with a bluish-gray 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.
Although the metal is reasonably abundant in the Earth's crust (it is 10 times more abundant than silver) 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 partially water-soluble and consequentially it can spread with groundwater when soils contain large amounts of the component. Thallium can also spread by adsorption on sludge. There are indications that thallium is fairly mobile within soils.
Thallium is found in the minerals Crookesite TlCu7Se4, Hutchinsonite TlPbAs5S9, and Lorandite TlAsS2. This metal is also contained in Pyrites and is extracted as a by-product of Sulfuric Acid production when Pyrite ore is roasted. Another way this element is obtained is 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.
There are no uses for metallic Thallium since pure Thallium quickly combines with Oxygen and water vapor from the atmosphere, forming a black, powdery substance. Thallium, used in conjunction with Sulfur or Selenium and Arsenic, forms low melting glass. Thallium Sulfate (Tl2SO4), an odorless, tasteless Thallium compound, was once used as a rat and ant poison, although it has been banned from household use in the United States since 1974. Thallium Sulfide (Tl2S), Thallium Iodide (TlI) and Thallium Bromide (TlBr) are all compounds used in devices to detect infrared radiation.
The oderless and tasteless Thallium Sulfate was widely used in the past as a rat poison and ant killer. In the United States and many other countries this use is no longer allowed due to safety concerns. Other uses:
In addition, 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.
|Crookesite, TlCu7Se4||Hutchinsonite, TlPbAs5S9|
|Lorandite, TlAsS2||Thallium Sulfate, Tl2SO4|
|Thallium Sulfide, Tl2S||Thallium Iodide, TlI|
|Thallium Bromide, TlBr|
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.
Thallium-202 (half life 12.23 days) can be made in a cyclotron while thallium-204 (half life 3.78 years) is made by the neutron activation of stable Thallium in a nuclear reactor.
|Thallium and its compounds are highly 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 oer m2 of skin in an 8-hour time-weighted average (40-hour work week). Thallium is a suspected human carcinogen.
Part of the reason for Thallium's high toxicity is that, when present in aqueous solution as the univalent Thallium (I) ion (Tl+), it exhibits some similarities with essential alkali metal cations, particularly Potassium. It can thus enter the body via potassium uptake pathways. However other aspects of Thallium's chemistry are very different from that of the alkali metals (e.g. its high affinity for Sulfur ligands), and so this substitution disrupts many cellular processes (for instance Thallium may attack Sulphur-containing proteins such as cysteine residues and ferredoxins).
Health Effects of Thallium
The element and its compounds are toxic and should be handled carefully.
Thallium occurs in the environment naturally in small amounts. It is not applied very
widely by humans, merely as rat poison and as a substance in electro-technical and
chemical industries. These applications can cause human exposure to thallium substances.
The human body absorbs Thallium very effectively, especially through the skin, the breathing organs and the digestive tract.
Thallium poisoning is mainly caused by accidental uptake of rat poison, which contains large amounts of Thallium Sulphate. Consequently, stomachaches will appear and the nervous system will be damaged. In some cases the damage is so irreversible that death will soon follow. When a human survives Thallium poisoning often consequences of disturbances of the nervous system, such as trembling, paralyses and behavioural changes will remain. With unborn children Thallium poisoning can cause congenital disorders.
Due to accumulation of Thallium in the bodies of humans, chronic effects consist, such as tiredness, headaches, depressions, lack of appetite, leg pains, hair loss and disturbances of the sight.
|Amongst the distinctive effects of Thallium poisoning are loss of hair (which led it to its initial use as a depilatory before its toxicity was properly appreciated) and damage to peripheral nerves (victims may experience a sensation of walking on hot coals). Thallium was once an effective murder weapon before its effects became understood and an antidote (Prussian Blue) discovered.|
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. The Prussian Blue is fed by mouth to the person, and it passes through their digestive system and comes out in the stool.
Famous Uses as a Poison
Atomic Radius (Å): 2.08Å
Electrochemical Equivalents: 7.625 g/amp-hr
Atomic Mass Average: 204.3833
(Gr. thallos, a green shoot or twig) Thallium was discovered spectroscopically in 1861 by Crookes. the element was named after the beautiful green spectral line, which identified the element. The metal was isolated both by Crookes and Lamy in 1862 about the same time. Thallium occurs in crooksite, lorandite, and hutchinsonite. It is also present in pyrites and is recovered from the roasting of this ore in connection iwth the production of sulfuric acid. It is also obtained from the smelting of lead and zinc ores. Extraction is somewhat complex and depends on the source of the thallium. Manganese nodules, found on the ocean floor, contain thallium. When freshly exposed to air, thallium exhibits a metallic luster, but soon develops a bluish-gray tinge, resembling lead in appearance. A heavy oxide builds up on thallium if left in air, and in the presence of water the hydride is formed. The metal is very soft and malleable. It can be cut with a knife. Thirty-two isotopic forms of thallium, with atomic masses ranging from 176 to 210 are recognized. Natural thallium is a mixture of two isotopes. The element and its compounds are toxic and should be handled carefully. Contact of the metal with skin is dangerous, and when melting the metal adequate ventilation should be provided. Exposure to thallium (soluble compounds) - skin, as Tl, should not exceed 0.1 mg/m3 (8-hour time-weighted average - 40-hour work week). Thallium is suspectied of carcinogenic potential for man. Thallium sulfate has been widely employed as a rodenticide and ant killer. It is odorless and tasteless, giving no warning of its presence. Its use, however, has been prohibited in the U.S. since 1975 as a household insecticide and rodenticide. The electrical conductivity of thallium sulfide changes with exposure to infrared light, and this compound is used in photocells. Thallium bromide-iodide crystals have been used as infrared optical meterials. Thallium has been used, with sulfur or slenium and arsenic, to produce low melting glasses with become fluid between 125 and 150oC. These glasses have properties at room temperatures similar to ordinary glasses and are said to be durable and insoluble in water. Thallium oxide has been used to produce glasses with a high index of refraction. Thallium has been used in treating ringworm and other skin infections; however, its use has been limited because of the narrow margin between toxicity and therapeutic benefits. A mercury-thallium alloy, which forms a eutectic at 8.5% thallium, is reported to freeze at -60oC, some 20 degrees below the freezing point of mercury. Commercial thallium metal (99%) costs about $40/lb.
Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond
|Atomic Mass||204.383 g.mol -1|
|Electronegativity According to Pauling||1.8|
|Density||11.71 g.cm-3 at 20°C|
|Melting Point||1800 °C|
|Boiling Point||4200 °C|
|Van der Waals Radius||0.182 nm|
|Ionic Radius||0.099 nm|
|Electronic Shell||[ Xe ] 4f14 5d10 6s2 6p1|
|Energy of First Ionization||589.1 kJ.mol -1|
|Energy of Second Ionization||1970.5 kJ.mol -1|
|Energy of Third Ionization||2877.4 kJ.mol -1|
|Discovered by:||Sir William Crookes in 1861|