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Indium is a chemical element with symbol In and atomic number 49. This rare, very soft, malleable and easily fusible post-transition metal is chemically similar to gallium and thallium, and shows intermediate properties between these two. Indium was discovered in 1863 and named for the indigo blue line in its spectrum that was the first indication of its existence in zinc ores, as a new and unknown element. The metal was first isolated in the following year. Zinc ores continue to be the primary source of indium, where it is found in compound form. Very rarely the element can be found as grains of native (free) metal, but these are not of commercial importance.
Indium's current primary application is to form transparent electrodes from indium tin oxide (ITO) in liquid crystal displays and touchscreens, and this use largely determines its global mining production. It is widely used in thin-films to form lubricated layers (during World War II it was widely used to coat bearings in high-performance aircraft). It is also used for making particularly low melting point alloys, and is a component in some lead-free solders.
Indium is not known to be used by any organism. In a similar way to aluminium salts, indium(III) ions can be toxic to the kidney when given by injection, but oral indium compounds do not have the chronic toxicity of salts of heavy metals, probably due to poor absorption in basic conditions. Radioactive indium-111 (in very small amounts on a chemical basis) is used in nuclear medicine tests, as a radiotracer to follow the movement of labeled proteins and white blood cells in the body.
Indium is a very soft, silvery-white, relatively rare poor metal with a bright luster. Like tin, when it is bent indium emits a high-pitched "cry". Like gallium, indium is able to wet glass. Like both, indium has a low melting point, 156.60 °C (313.88 °F); higher than its lighter homologue, gallium, but lower than its heavier homologue, thallium, and lower than tin. Its boiling point is, however, moderate, being 2072 °C (3762 °F), which is higher than that of thallium, but lower than that of gallium, showing opposition to melting points trend. The density of indium, 7.31 g·cm−3, is also higher than that of gallium, but lower than that of thallium.
An indium atom has 49 electrons, having an electronic configuration of [Kr]4d105s25p1. In its compounds, indium most commonly loses its three outermost electrons, becoming indium(III) ions, In3+, but in some cases the pair of 5s-electrons can stay within the atom, indium thus oxidized only to indium(I), In+. This happens due to inert pair effect, which occurs because of stabilization of 5s-orbital due to relativistic effects, which are stronger closer to the bottom of the periodic table. Thallium shows an even stronger effect, making oxidation to thallium(I) more likely than to thallium(III), making +1 the more likely oxidation state.
A number of standard electrode potentials, depending on the reaction under study, are reported for indium:
Indium is a post-transition metal and chemically, is the intermediate element between its group 13 neighbors gallium and thallium. It shows two main oxidation states, which are +1 and +3, with latter being more stable, whereas the only common oxidation state of gallium is +3 and thallium shows +1 more likely than +3, with thallium(III) being a moderately strong oxidizing agent, while indium(III) is stable and indium(I) is a powerful reducing agent.
Indium does not react with water, but it is oxidized by stronger oxidizing agents, such as halogens or oxalic acid, to give indium(III) compounds. It does not react with boron, silicon or carbon, and the corresponding boride, silicide or carbide are not known. Similarly, reaction between indium and hydrogen has not been observed, but both indium(I) and indium(III) hydrides are known.
Indium(III) oxide is formed at high temperatures during reaction between indium and oxygen, with blue flame. It is amphoteric, i. e. it can react with both acids and bases. Its reaction with water results in insoluble indium(III) hydroxide, which is also amphoteric, reacting with alkalies to give indates(III) and with acids to give indium(III) salts:
The hydrolysis of sodium indate(III) gives weak indic acid, HInO2. Out of common indium(III) salts, chloride, sulfate and nitrate are soluble. In water solutions, In3+ and [InO2]- ions are hydrolyzed to give InOH2+ and HInO2 due to generally amphoteric character of indium(III) ions. Indium(III) compounds are not well-soluble, similarly to thallium(III) compounds; however, indium(III) salts of strong acids, such as chloride, sulfate and nitrate are soluble, hydrolyzing in water solutions. The In3+ ion is colorless in solution because of the absence of unpaired electrons in the d- and f-electron shells.
Indium(I) compounds are not as common as indium(III) ones; only chloride, bromide, iodide, sulfide and cyclopentadienyl are well-characterized. Indium(I) sulfide is the product of reaction between indium and sulfur or indium and hydrogen sulfide, and can be received at 700—1000 °C. Indium(I) oxide black powder is received at 850 °C during reaction between indium and carbon dioxide or during decomposition of indium(III) oxide at 1200 °C. Cyclopentadienylindium(I), which was the first organoindium(I) compound reported, is polymer consisting of zigzag chains of alternating indium atoms and cyclopentadienyl complexes.
Less frequently, indium shows intermediate oxidation state +2, which lies between the common ones, most notably in halides, In2X4 and [In2X6]2-. Several other compounds are known to combine indium(I) and indium(III), such as InI6(InIIICl6)Cl3, InI5(InIIIBr4)2(InIIIBr6), InIInIIIBr4.
Indium occurs naturally on Earth only in two primordial nuclides, indium-113 and indium-115. Out of this two, indium-115 makes up 95.7% of all indium but it is radioactive, decaying to tin-115 via beta decay with half-life of 4.411014 years, four orders of magnitude larger than the age of the universe and nearly 50,000 times longer than that of natural thorium. This situation is uncommon among stable chemical elements; only indium, tellurium, and rhenium have been shown to have most-abundant isotopes that are radioactive. The less common natural isotope of indium, indium-113, is stable.
Indium has 39 known isotopes, ranging in mass between 97 and 135. Only one of them is stable and one has half-life exceeding 1014 years; the most stable other indium isotope is indium-111, which has half-life of approximately 2.8 days. All other isotopes have half-lives shorter than 5 hours. Indium also has 47 meta states, out of which indium-114m1 is the most stable, being more stable than ground state of any indium isotope, except for the primordial ones.
Indium is created via the long-lasting, (up to thousands of years), s-process in low-medium mass stars (which range in mass between 0.6 and 10 masses of Sun). When a silver-109 atom, which comprises approximately half of all silver in existence, catches a neutron, it undergoes a beta decay to become cadmium-110. Capturing further neutrons, it becomes cadmium-115, which decays to indium-115 via another beta decay. This explains why the radioactive isotope predominates in abundance compared to the stable one.
Indium is 61st most abundant element in the Earth's crust at approximately 49 ppb, making indium approximately as abundant as mercury. Fewer than 10 indium minerals are known, such as dzhalindite (In(OH)3) and indite (FeIn2S4), but none of these occurs in significant deposits.
Based on content of indium in zinc ore stocks, there is a worldwide reserve base of approximately 6,000 tonnes of economically viable indium. However, the Indium Corporation, the largest processor of indium, claims that, on the basis of increasing recovery yields during extraction, recovery from a wider range of base metals (including tin, copper and other polymetallic deposits) and new mining investments, the long-term supply of indium is sustainable, reliable and sufficient to meet increasing future demands.
This conclusion may be reasonable considering that silver, which is one-third as abundant as indium in the Earth's crust, is currently mined at approximately 18,300 tonnes per year, which is 40 times greater than current indium mining rates.
In 1863, the German chemists Ferdinand Reich and Hieronymous Theodor Richter were testing ores from the mines around Freiberg, Saxony. They dissolved the minerals pyrite, arsenopyrite, galena and sphalerite in hydrochloric acid and distilled raw zinc chloride. As it was known that ores from that region sometimes contain thallium they searched for the green emission lines with spectroscopy. The green lines were absent but a blue line was present in the spectrum. As no element was known with a bright blue emission they concluded that a new element was present in the minerals. They named the element with the blue spectral line indium, from the indigo color seen in its spectrum. Richter went on to isolate the metal in 1864. At the World Fair 1867 an ingot of 0.5 kg (1.1 lb) was presented.
In 1924, indium was found to have a valuable ability to stabilize non-ferrous metals, which was the first significant use for the element. It took until 1936 for the U.S. Bureau of Mines to list indium as a commodity, and even in early the 1950s only very limited applications for indium were known, the most important of which was making light-emitting diodes and coating bearings for aircraft engines during World War II. The start of production of indium-containing semiconductors started in 1952. The development and widespread use of indium-containing nuclear control rods increased demand during the 1970s, and the use of indium tin oxide in liquid crystal displays increased and became the major application by 1992.
The lack of indium mineral deposits and the fact that indium is enriched in sulfidic lead, tin, copper, iron and predominately in zinc deposits, makes zinc production the main source for indium. The indium is leached from slag and dust of zinc production. Further purification is done by electrolysis. The exact process varies with the exact composition of the slag and dust.
Indium is produced mainly from residues generated during zinc ore processing but is also found in iron, lead, and copper ores. China is a leading producer of indium (390 tonnes in 2012), followed by Canada, Japan and South Korea with 70 tonnes each. The Teck Cominco refinery in Trail, British Columbia, is a large single-source indium producer, with an output of 32.5 tonnes in 2005, 41.8 tonnes in 2004 and 36.1 tonnes in 2003. South American Silver Corporation's Malku Khota property in Bolivia is a large resource of indium with an indicated resource of 1,481 tonnes and inferred resource of 935 tonnes. Adex Mining Inc.’s Mount Pleasant Mine in New Brunswick, Canada, holds some of the world’s total known indium resources.
The amount of indium consumed is largely a function of worldwide LCD production. Worldwide production is currently 475 tonnes per year from mining and a further 650 tonnes per year from recycling. Demand has risen rapidly in recent years with the popularity of LCD computer monitors and television sets, which now account for 50% of indium consumption. Increased manufacturing efficiency and recycling (especially in Japan) maintain a balance between demand and supply. According to the UNEP, indium's end-of-life recycling rate is less than 1%. Demand increased as the metal is used in LCDs and televisions, and supply decreased when a number of Chinese mining concerns stopped extracting indium from their zinc tailings. In 2002, the price was US$94 per kilogram. The recent changes in demand and supply have resulted in high and fluctuating prices of indium, which from 2006 to 2009 ranged from US$382/kg to US$918/kg.
It has been estimated that there is less than 20 years left of indium supplies, based on current rates of extraction, demonstrating the need for additional recycling.
The first large-scale application for indium was as a coating for bearings in high-performance aircraft engines during World War II. Afterward, production gradually increased as new uses were found in fusible alloys, solders, and electronics. In the 1950s, tiny beads of it were used for the emitters and collectors of PNP alloy junction transistors. In the middle and late 1980s, the development of indium phosphide semiconductors and indium tin oxide thin films for liquid crystal displays (LCD) aroused much interest. By 1992, the thin-film application had become the largest end use.
Pure indium in metal form is considered nontoxic by most sources. In the welding and semiconductor industries, where indium exposure is relatively high, there have been no reports of any toxic side-effects. Indium compounds, like aluminum compounds, complex with hydroxyls to form insoluble salts in basic conditions, and are thus not well-absorbed from food, giving them fairly low oral toxicty. Soluble indium(III) is toxic when delivered parenterally, however, causing damage primarily to the kidney (both inner and outer parts), but additionally to heart and liver, and may be teratogenic. Other indium compounds are toxic when administered outside the gastrointestinal tract: for example, anhydrous indium trichloride (InCl3) and indium phosphide (InP) are quite toxic when delivered into the lungs (the latter is a suspected carcinogen). Occupational exposure to indium compounds was associated with PAP, cholesterol ester crystals and granulomas, pulmonary fibrosis, emphysema, and pneumothoraces. The available evidence suggests exposure to indium compounds causes a novel lung disease that may begin with PAP and progress to include fibrosis and emphysema, and, in some cases, premature death
Gallium is a chemical element with symbol Ga and atomic number 31. Elemental gallium does not occur in nature, but as the gallium(III) compounds in trace amounts in bauxite and zinc ores. A soft silvery metallic poor metal, elemental gallium is a brittle solid at low temperatures. Held long enough, gallium will melt in the hand as it liquefies at temperature of (slightly above room temperature). Its melting point is used as a temperature reference point. The alloy Galinstan (68.5% Ga, 21.5% In, 10% Sn) has an even lower melting point of , well below the freezing point of water. From its discovery in 1875 until the semiconductor era, gallium was used primarily as an agent to make low-melting alloys.
Today, almost all gallium is used for microelectronics. Gallium arsenide, the primary use of gallium, is used in microwave circuitry and infrared applications. Gallium nitride and indium gallium nitride, minority semiconductor uses, produce blue and violet light-emitting diodes (LEDs) and diode lasers.
Gallium has no known role in biology. Because gallium(III) and ferric salts behave similarly in biological systems, gallium ions often mimic iron ions in medical applications. Gallium-containing pharmaceuticals and radiopharmaceuticals have been developed.
Elemental gallium is not found in nature, but it is easily obtained by smelting. Very pure gallium metal has a brilliant silvery color and its solid metal fractures conchoidally like glass. Gallium metal expands by 3.1% when it solidifies, and therefore storage in either glass or metal containers is avoided, due to the possibility of container rupture with freezing. Gallium shares the higher-density liquid state with only a few materials like silicon, germanium, bismuth and water.
Gallium attacks most other metals by diffusing into their metal lattice. Gallium, for example, diffuses into the grain boundaries of Al/Zn alloys or steel, making them very brittle. Gallium easily alloys with many metals, and is used in small quantities as a plutonium-gallium alloy in the plutonium cores of nuclear bombs, to help stabilize the plutonium crystal structure.
The melting point of 302.9146 K (29.7646 °C, 85.5763 °F) is near room temperature, about the average summer time temperatures. Gallium's melting point (mp) is one of the formal temperature reference points in the International Temperature Scale of 1990 (ITS-90) established by BIPM. The triple point of gallium of 302.9166 K (29.7666 °C, 85.5799 °F), is being used by NIST in preference to gallium's melting point.
The unique melting point of gallium allows it to melt in one's hand, and then refreeze if removed. This metal has a strong tendency to supercool below its melting point/freezing point. Seeding with a crystal helps to initiate freezing. Gallium is one of the metals (with caesium, rubidium, mercury and likely francium) that are liquid at or near-normal room temperature, and can therefore be used in metal-in-glass high-temperature thermometers. It is also notable for having one of the largest liquid ranges for a metal, and (unlike mercury) for having a low vapor pressure at high temperatures. Gallium's boiling point, 2477 K, is more than eight times higher than its melting point on the absolute scale, making it the greatest ratio between melting point and boiling point of any element. Unlike mercury, liquid gallium metal wets glass and skin, making it mechanically more difficult to handle (even though it is substantially less toxic and requires far fewer precautions). For this reason as well as the metal contamination and freezing-expansion problems, samples of gallium metal are usually supplied in polyethylene packets within other containers.
Gallium does not crystallize in any of the simple crystal structures. The stable phase under normal conditions is orthorhombic with 8 atoms in the conventional unit cell. Each atom has only one nearest neighbor (at a distance of 244 pm) and six other neighbors within additional 39 pm. Many stable and metastable phases are found as function of temperature and pressure.
The bonding between the two nearest neighbors is covalent, hence Ga2 dimers are seen as the fundamental building blocks of the crystal. This explains the drop of the melting point compared to its neighbor elements aluminium and indium.
The physical properties of gallium are highly anisotropic, i.e. have different values along the three major crystallographical axes a, b and c (see table); for this reason, there is a significant difference between the linear (α) and volume thermal expansion coefficients. The properties of gallium are also strongly temperature-dependent, especially near the melting point. For example, the thermal expansion coefficient increases by several hundred percent upon melting.
Gallium is found primarily in the +3 oxidation state. The +1 oxidation is also attested in some compounds, although they tend to disproportionate into elemental gallium and gallium(III) compounds. What are sometimes referred to as gallium(II) compounds are actually mixed-oxidation state compounds containing both gallium(I) and gallium(III).
At room temperature, gallium metal is unreactive towards air and water due to the formation of a passive, protective oxide layer. At higher temperatures, however, it reacts with oxygen in the air to form gallium(III) oxide, . Reducing with elemental gallium in vacuum at 500 °C to 700 °C yields the dark brown gallium(I) oxide, .:285 is a very strong reducing agent, capable of reducing to .:207 It disproportionates at 800 °C back to gallium and .
Gallium sulfide, , has 3 possible crystal modifications.:104 It can be made by the reaction of gallium with hydrogen sulfide () at 950 °C.:162 Alternatively, can also be used at 747 °C:
Reacting a mixture of alkali metal carbonates and with leads to the formation of thiogallates containing the anion. Strong acids decompose these salts, releasing in the process.:104–105 The mercury salt, , can be used as a phosphor.
Gallium also forms sulfides in lower oxidation states, such as gallium(II) sulfide and the green gallium(I) sulfide, the latter of which is produced from the former by heating to 1000 °C under a stream of nitrogen.:94
The other binary chalcogenides, and , have zincblende structure. They are all semiconductors, but are easily hydrolysed, limiting their usefulness.:104
Strong acids dissolve gallium, forming gallium(III) salts such as and . Aqueous solutions of gallium(III) salts contain the hydrated gallium ion, .:1033 Gallium(III) hydroxide, , may be precipitated from gallium(III) solutions by adding ammonia. Dehydrating at 100 °C produces gallium oxide hydroxide, GaO(OH).:140–141
Alkaline hydroxide solutions dissolve gallium, forming gallate salts containing the anion.:1033 Gallium hydroxide, which is amphoteric, also dissolves in alkali to form gallate salts.:141 Although earlier work suggested as another possible gallate anion, this species was not found in later work.
Gallium reacts with ammonia at 1050 °C to form gallium nitride, GaN. Gallium also forms binary compounds with phosphorus, arsenic, and antimony: gallium phosphide (GaP), gallium arsenide (GaAs), and gallium antimonide (GaSb). These compounds have the same structure as ZnS, and have important semiconducting properties.:1034 GaP, GaAs, and GaSb can be synthesized by the direct reaction of gallium with elemental phosphorus, arsenic, or antimony.:99 They exhibit higher electrical conductivity than GaN.:101 GaP can also be synthesized by the reaction of with phosphorus at low temperatures.
Gallium also forms ternary nitrides; for example::99
Similar compounds with phosphorus and arsenic also exist: and . These compounds are easily hydrolyzed by dilute acids and water.:101
Gallium(III) oxide reacts with fluorinating agents such as HF or to form gallium(III) fluoride, . It is an ionic compound strongly insoluble in water. However, it does dissolve in hydrofluoric acid, in which it forms an adduct with water, . Attempting to dehydrate this adduct instead forms . The adduct reacts with ammonia to form , which can then be heated to form anhydrous .:128–129
Gallium trichloride is formed by the reaction of gallium metal with chlorine gas. Unlike the trifluoride, gallium(III) chloride exists as dimeric molecules, , with a melting point of 78 °C. This is also the case for the bromide and iodide, and .:133
Like the other group 13 trihalides, gallium(III) halides are Lewis acids, reacting as halide acceptors with alkali metal halides to form salts containing anions, where X is a halogen. They also react with alkyl halides to form carbocations and .:136–137
When heated to a high temperature, gallium(III) halides react with elemental gallium to form the respective gallium(I) halides. For example, reacts with Ga to form :
At lower temperatures, the equilibrium shifts toward the left and GaCl disproportionates back to elemental gallium and . GaCl can also be made by the reaction of Ga with HCl at 950 °C; it can then be condensed as red solid.:1036
Gallium(I) compounds can be stabilized by forming adducts with Lewis acids. For example:
The so-called "gallium(II) halides", , are actually adducts of gallium(I) halides with the respective gallium(III) halides, having the structure . For example::1036
Like aluminium, gallium also forms a hydride, , known as gallane, which may be obtained by the reaction of lithium gallanate () with gallium(III) chloride at −30 °C::1031
In the presence of dimethyl ether as solvent, polymerizes to . If no solvent is used, the dimer (digallane) is formed as a gas. Its structure is similar to diborane, having two hydrogen atoms bridging the two gallium centers,:1031 unlike α- in which aluminium has a coordination number of 6.:1008
Gallane is unstable above −10 °C, decomposing to elemental gallium and hydrogen.
In 1871, existence of gallium was first predicted by Russian chemist Dmitri Mendeleev, who named it "eka-aluminium" on the basis of its position in his periodic table. He also predicted several properties of the element, which correspond closely to real gallium properties, such as density, melting point, oxide character and bonding in chloride.
Gallium was discovered spectroscopically by French chemist Paul Emile Lecoq de Boisbaudran in 1875 by its characteristic spectrum (two violet lines) in an examination of a sphalerite sample. Later that year, Lecoq obtained the free metal by electrolysis of its hydroxide in potassium hydroxide solution. He named the element "gallia", from Latin Gallia meaning Gaul, after his native land of France. It was later claimed that, in one of those multilingual puns so beloved of men of science in the 19th century, he had also named gallium after himself, as his name, "Le coq", is the French for "the rooster", and the Latin for "rooster" is "gallus"; however, in an 1877 article Lecoq denied this supposition. (Cf. the naming of the J/ψ meson and the dwarf planet Pluto.)
From its discovery in 1875 up to the era of semiconductors, its primary uses were in high-temperature thermometric applications and in preparation of metal alloys with unusual properties of stability, or ease of melting; some being liquid at room temperature or below. The development of gallium arsenide as a direct band gap semiconductor in the 1960s ushered in the most important stage in the applications of gallium.
Gallium does not exist in free form in nature, and the few high-gallium minerals such as gallite (CuGaS2) are too rare to serve as a primary source of the element or its compounds. Its abundance in the Earth's crust is approximately 16.9 ppm. Gallium is found and extracted as a trace component in bauxite and to a small extent from sphalerite. The amount extracted from coal, diaspore and germanite in which gallium is also present is negligible. The United States Geological Survey (USGS) estimates gallium reserves to exceed 1 million tonnes, based on 50 ppm by weight concentration in known reserves of bauxite and zinc ores. Some flue dusts from burning coal have been shown to contain small quantities of gallium, typically less than 1% by weight.
Gallium is a byproduct of the production of aluminium and zinc. Whereas the sphalerite for zinc production is the minor source. Most gallium is extracted from the crude aluminium hydroxide solution of the Bayer process for producing alumina and aluminium. A mercury cell electrolysis and hydrolysis of the amalgam with sodium hydroxide leads to sodium gallate. Electrolysis then gives gallium metal. For semiconductor use, further purification is carried out using zone melting, or else single crystal extraction from a melt (Czochralski process). Purities of 99.9999% are routinely achieved and commercially widely available.
In 1986, the production was estimated at 40 tons. In 2007 the production of gallium was 184 tonnes with less than 100 tonnes from mining and the rest from scrap recycling. By 2011 world production of gallium was an estimated 216 metric tons.
The semiconductor applications dominate the commercial use of gallium, accounting for 98% of applications. The next major application is for gadolinium gallium garnets.
Because of this application, extremely high-purity (99.9999+%) gallium is commercially available. Gallium arsenide (GaAs) and gallium nitride (GaN) used in electronic components represented about 98% of the gallium consumption in the United States in 2007. About 66% of semiconductor gallium is used in the U.S. in integrated circuits (mostly gallium arsenide), such as the manufacture of ultra-high speed logic chips and MESFETs for low-noise microwave preamplifiers in cell phones. About 20% is used in optoelectronics. Worldwide, gallium arsenide makes up 95% of the annual global gallium consumption.
Gallium arsenide is used in optoelectronics in a variety of infrared applications. Aluminium gallium arsenide (AlGaAs) is used in high-powered infrared laser diodes. As a component of the semiconductors indium gallium nitride and gallium nitride, gallium is used to produce blue and violet optoelectronic devices, mostly laser diodes and light-emitting diodes. For example, gallium nitride 405 nm diode lasers are used as a violet light source for higher-density compact disc data storage, in the Blu-ray Disc standard.
Multijunction photovoltaic cells, developed for satellite power applications, are made by molecular beam epitaxy or metalorganic vapour phase epitaxy of thin films of gallium arsenide, indium gallium phosphide or indium gallium arsenide.The Mars Exploration Rovers and several satellites use triple junction gallium arsenide on germanium cells. Gallium is also a component in photovoltaic compounds (such as copper indium gallium selenium sulfide or Cu(In,Ga)(Se,S)2) for use in solar panels as a cost-efficient alternative to crystalline silicon.
Gallium readily alloys with most metals, and has been used as a component in low-melting alloys. A nearly eutectic alloy of gallium, indium, and tin is a room temperature liquid that is available in medical thermometers. This alloy, with the trade-name Galinstan (with the "-stan" referring to the tin), has a low freezing point of −19 °C (−2.2 °F). It has been suggested that this family of alloys could also be used to cool computer chips in place of water. Gallium alloys have been evaluated as substitutes for mercury dental amalgams, but these materials have yet to see wide acceptance.
Because gallium wets glass or porcelain, gallium can be used to create brilliant mirrors. When the wetting action of gallium-alloys is not desired (as in Galinstan glass thermometers), the glass must be protected with a transparent layer of gallium(III) oxide.
The plutonium used in nuclear weapon pits is machined by alloying with gallium to stabilize its δ phase.
Gallium added in quantities up to 2% in common solders can aid wetting and flow characteristics.
Alloys of Al and Ga have been evaluated for hydrogen production.
It is used as alloying element in the magnetic shape-memory alloy Ni-Mn-Ga.
Although gallium has no known role in biology, it mimics iron(III), the gallium ion localizes to and interacts with many processes in the body in which iron(III) is manipulated. As these processes include inflammation, which is a marker for many disease states, several gallium salts are used, or are in development, as both pharmaceuticals and radiopharmaceuticals in medicine. When gallium ions are mistakenly taken up by bacteria such as Pseudomonas, the bacteria's ability to respire is interfered with and the bacteria die. The mechanism behind this is that iron is redox active, which allows for the transfer of electrons during respiration, but gallium is redox inactive.
Gallium nitrate (brand name Ganite) has been used as an intravenous pharmaceutical to treat hypercalcemia associated with tumor metastasis to bones. Gallium is thought to interfere with osteoclast function. It may be effective when other treatments for maligancy-associated hypercalcemia are not.
A complex amine-phenol Ga(III) compound MR045 was found to be selectively toxic to parasites that have developed resistance to chloroquine, a common drug against malaria. Both the Ga(III) complex and chloroquine act by inhibiting crystallization of hemozoin, a disposal product formed from the digestion of blood by the parasites.
Gallium-67 salts such as gallium citrate and gallium nitrate are used as radiopharmaceutical agents in a nuclear medicine imaging procedure commonly referred to as a gallium scan. The form or salt of gallium is unimportant. For these applications, the radioactive isotope 67Ga is used. The body handles Ga3+ in many ways as though it were iron, and thus it is bound (and concentrates) in areas of inflammation, such as infection, and also areas of rapid cell division. This allows such sites to be imaged by nuclear scan techniques. This use has largely been replaced by fluorodeoxyglucose (FDG) for positron emission tomography, "PET" scan and indium-111 labelled leukocyte scans. However, the localization of gallium in the body has some properties which make it unique in some circumstances from competing modalities using other radioisotopes.
Gallium-68, a positron emitter with a half life of 68 min., is now used as a diagnostic radionuclide in PET-CT when linked to pharmaceutical preparations such as DOTATOC, a somatostatin analogue used for neuroendocrine tumors investigation, and DOTA-TATE, a newer one, used for neuroendocrine metastasis and lung neuroendocrine cancer, such as certain types of microcytoma. Galium-68's preparation as a pharmaceutical is chemical and the radionuclide is extracted by elution from germanium-68, a synthetic radioisotope of germanium, in gallium-68 generators.
The Ga(III) ion of soluble gallium salts tends to form the insoluble hydroxide when injected in large amounts, and in animals precipitation of this has resulted in renal toxicity. In lower doses, soluble gallium is tolerated well, and does not accumulate as a poison.
While metallic gallium is not considered toxic, the data are inconclusive. Some sources suggest that it may cause dermatitis from prolonged exposure; other tests have not caused a positive reaction. Like most metals, finely divided gallium loses its luster and powdered gallium appears gray. Thus, when gallium is handled with bare hands, the extremely fine dispersion of liquid gallium droplets, which results from wetting skin with the metal, may appear as a gray skin stain.][
Hafnium is a chemical element with the symbol Hf and atomic number 72. A lustrous, silvery gray, tetravalent transition metal, hafnium chemically resembles zirconium and is found in zirconium minerals. Its existence was predicted by Dmitri Mendeleev in 1869. Hafnium was the penultimate stable isotope element to be discovered (rhenium was identified two years later). Hafnium is named after Hafnia, the Latin name for "Copenhagen", where it was discovered.
Hafnium is used in filaments and electrodes. Some semiconductor fabrication processes use its oxide for integrated circuits at 45 nm and smaller feature lengths. Some superalloys used for special applications contain hafnium in combination with niobium, titanium, or tungsten.
Hafnium's large neutron capture cross-section makes it a good material for neutron absorption in control rods in nuclear power plants, but at the same time requires that it be removed from the neutron-transparent corrosion-resistant zirconium alloys used in nuclear reactors.
Hafnium is a shiny, silvery, ductile metal that is corrosion-resistant and chemically similar to zirconium (due to its having the same number of valence electrons and being in the same group). The physical properties of hafnium metal samples are markedly affected by zirconium impurities, especially the nuclear properties, as these two elements are among the most difficult to separate because of their chemical similarity.
A notable physical difference between these metals is their density, with zirconium having about one-half the density of hafnium. The most notable nuclear properties of hafnium are its high thermal neutron-capture cross-section and that the nuclei of several different hafnium isotopes readily absorb two or more neutrons apiece. In contrast with this, zirconium is practically transparent to thermal neutrons, and it is commonly used for the metal components of nuclear reactors – especially the claddings of their nuclear fuel rods.
Hafnium reacts in air to form a protective film that inhibits further corrosion. The metal is not readily attacked by acids but can be oxidized with halogens or it can be burnt in air. Like its sister metal zirconium, finely divided hafnium can ignite spontaneously in air—similar to that obtained in Dragon's Breath. The metal is resistant to concentrated alkalis.
The chemistry of hafnium and zirconium is so similar that the two cannot be separated on the basis of differing chemical reactions. The melting points and boiling points of the compounds and the solubility in solvents are the major differences in the chemistry of these twin elements.
At least 34 isotopes of hafnium have been observed, ranging in mass number from 153 to 186. The five stable isotopes are in the range of 176 to 180. The radioactive isotopes' half-lives range from only 400 ms for 153Hf, to 2.0 petayears (1015 years) for the most stable one, 174Hf.
The nuclear isomer 178m2Hf was at the center of a controversy for several years regarding its potential use as a weapon.
Hafnium is estimated to make up about 5.8 ppm of the Earth's upper crust by weight. It does not exist as a free element in nature, but is found combined in solid solution for zirconium in natural zirconium compounds such as zircon, ZrSiO4, which usually has about 1 – 4% of the Zr replaced by Hf. Rarely, the Hf/Zr ratio increases during crystallization to give the isostructural mineral 'hafnon' (Hf,Zr)SiO4, with atomic Hf > Zr. An old (obsolete) name for a variety of zircon containing unusually high Hf content is alvite.
A major source of zircon (and hence hafnium) ores are heavy mineral sands ore deposits, pegmatites particularly in Brazil and Malawi, and carbonatite intrusions particularly the Crown Polymetallic Deposit at Mount Weld, Western Australia. A potential source of hafnium is trachyte tuffs containing rare zircon-hafnium silicates eudialyte or armstrongite, at Dubbo in New South Wales, Australia.
Hafnium reserves are projected to last under 10 years if the world population increases and demand grows.
The heavy mineral sands ore deposits of the titanium ores ilmenite and rutile yield most of the mined zirconium, and therefore also most of the hafnium.
Zirconium is a good nuclear fuel-rod cladding metal, with the desirable properties of a very low neutron capture cross-section and good chemical stability at high temperatures. However, because of hafnium's neutron-absorbing properties, hafnium impurities in zirconium would cause it to be far less useful for nuclear-reactor applications. Thus, a nearly complete separation of zirconium and hafnium is necessary for their use in nuclear power. The production of hafnium-free zirconium is the main source for hafnium.
The chemical properties of hafnium and zirconium are nearly identical, which makes the two difficult to separate. The methods first used — fractional crystallization of ammonium fluoride salts or the fractionated distillation of the chloride — have not proven suitable for an industrial-scale production. After zirconium was chosen as material for nuclear reactor programs in the 1940s, a separation method had to be developed. Liquid-liquid extraction processes with a wide variety of solvents were developed and are still used for the production of hafnium. About half of all hafnium metal manufactured is produced as a by-product of zirconium refinement. The end product of the separation is hafnium(IV) chloride. The purified hafnium(IV) chloride is converted to the metal by reduction with magnesium or sodium, as in the Kroll process.
Further purification is effected by a chemical transport reaction developed by Arkel and de Boer: In a closed vessel, hafnium reacts with iodine at temperatures of 500 °C, forming hafnium(IV) iodide; at a tungsten filament of 1700 °C the reverse reaction happens, and the iodine and hafnium are set free. The hafnium forms a solid coating at the tungsten filament, and the iodine can react with additional hafnium, resulting in a steady turn over.
Hafnium and zirconium form nearly identical series of chemical compounds. Hafnium tends to form inorganic compounds in the oxidation state of +4. Halogens react with it to form hafnium tetrahalides. At higher temperatures, hafnium reacts with oxygen, nitrogen, carbon, boron, sulfur, and silicon. Due to the lanthanide contraction of the elements in the sixth period, zirconium and hafnium have nearly identical ionic radii. The ionic radius of Zr4+ is 0.79 angstrom and that of Hf4+ is 0.78 angstrom.
Hafnium(IV) chloride and hafnium(IV) iodide have some applications in the production and purification of hafnium metal. They are volatile solids with polymeric structures. These tetrachlorides are precursors to various organohafnium compounds such as hafnocene dichloride and tetrabenzylhafnium.
The white hafnium oxide (HfO2), with a melting point of 2812 °C and a boiling point of roughly 5100 °C, is very similar to zirconia, but slightly more basic. Hafnium carbide is the most refractory binary compound known, with a melting point over 3890 °C, and hafnium nitride is the most refractory of all known metal nitrides, with a melting point of 3310 °C. This has led to proposals that hafnium or its carbides might be useful as construction materials that are subjected to very high temperatures. The mixed carbide tantalum hafnium carbide () possesses the highest melting point of any currently known compound, 4215 °C.
In his report on The Periodic Law of the Chemical Elements, in 1869, Dmitri Mendeleev had implicitly predicted the existence of a heavier analog of titanium and zirconium. At the time of his formulation in 1871, Mendeleev believed that the elements were ordered by their atomic masses and placed lanthanum (element 57) in the spot below zirconium. The exact placement of the elements and the location of missing elements was done by determining the specific weight of the elements and comparing the chemical and physical properties.
The X-ray spectroscopy done by Henry Moseley in 1914 showed a direct dependency between spectral line and effective nuclear charge. This led to the nuclear charge, or atomic number of an element, being used to ascertain its place within the periodic table. With this method, Moseley determined the number of lanthanides and showed the gaps in the atomic number sequence at numbers 43, 61, 72, and 75.
The discovery of the gaps led to an extensive search for the missing elements. In 1914, several people claimed the discovery after Henry Moseley predicted the gap in the periodic table for the then-undiscovered element 72. Georges Urbain asserted that he found element 72 in the rare earth elements in 1907 and published his results on celtium in 1911. Neither the spectra nor the chemical behavior matched with the element found later, and therefore his claim was turned down after a long-standing controversy. The controversy was partly because the chemists favored the chemical techniques which led to the discovery of celtium, while the physicists relied on the use of the new X-ray spectroscopy method that proved that the substances discovered by Urbain did not contain element 72. By early 1923, several physicists and chemists such as Niels Bohr and Charles R. Bury suggested that element 72 should resemble zirconium and therefore was not part of the rare earth elements group. These suggestions were based on Bohr's theories of the atom, the X-ray spectroscopy of Mosley, and the chemical arguments of Friedrich Paneth.
Encouraged by these suggestions and by the reappearance in 1922 of Urbain's claims that element 72 was a rare earth element discovered in 1911, Dirk Coster and Georg von Hevesy were motivated to search for the new element in zirconium ores. Hafnium was discovered by the two in 1923 in Copenhagen, Denmark, validating the original 1869 prediction of Mendeleev. It was ultimately found in zircon in Norway through X-ray spectroscopy analysis. The place where the discovery took place led to the element being named for the Latin name for "Copenhagen", Hafnia, the home town of Niels Bohr. Today, the Faculty of Science of the University of Copenhagen uses in its seal a stylized image of the hafnium atom.
Hafnium was separated from zirconium through repeated recrystallization of the double ammonium or potassium fluorides by Valdemar Thal Jantzen and von Hevesey. Anton Eduard van Arkel and Jan Hendrik de Boer were the first to prepare metallic hafnium by passing hafnium tetra-iodide vapor over a heated tungsten filament in 1924. This process for differential purification of zirconium and hafnium is still in use today.
In 1923, four predicted elements were still missing from the periodic table: 43 (technetium) and 61 (promethium) are radioactive elements and are only present in trace amounts in the environment, thus making elements 75 (rhenium) and 72 (hafnium) the last two unknown non-radioactive elements. Since rhenium was discovered in 1925, hafnium was the next to last element with stable isotopes to be discovered.
Several details contribute to the fact that there are only a few technical uses for hafnium: First, the close similarity between hafnium and zirconium makes it possible to use zirconium for most of the applications; second, hafnium was first available as pure metal after the use in the nuclear industry for hafnium-free zirconium in the late 1950s. Furthermore, the low abundance and difficult separation techniques necessary make it a scarce commodity.
Most of the hafnium produced is used in the production of control rods for nuclear reactors.
The nuclei of several hafnium isotopes can each absorb multiple neutrons. This makes hafnium a good material for use in the control rods for nuclear reactors. Its neutron-capture cross-section is about 600 times that of zirconium. (Other elements that are good neutron-absorbers for control rods are cadmium and boron.) Excellent mechanical properties and exceptional corrosion-resistance properties allow its use in the harsh environment of a pressurized water reactors. The German research reactor FRM II uses hafnium as a neutron absorber.
Hafnium is used in iron, titanium, niobium, tantalum, and other metal alloys. An alloy used for liquid rocket thruster nozzles, for example the main engine of the Apollo Lunar Modules is C103, which consists of 89% niobium, 10% hafnium and 1% titanium.
Small additions of hafnium increase the adherence of protective oxide scales on nickel-based alloys. It improves thereby the corrosion resistance especially under cyclic temperature conditions that tend to break oxide scales by inducing thermal stresses between the bulk material and the oxide layer.
The electronics industry discovered that hafnium-based compound can be employed in gate insulators in the 45 nm generation of integrated circuits from Intel, IBM and others. Hafnium oxide-based compounds are practical high-k dielectrics, allowing reduction of the gate leakage current which improves performance at such scales.
Due to its heat resistance and its affinity to oxygen and nitrogen, hafnium is a good scavenger for oxygen and nitrogen in gas-filled and incandescent lamps. Hafnium is also used as the electrode in plasma cutting because of its ability to shed electrons into air.
The high energy content of 178m2Hf was the concern of a DARPA funded program in the US. This program determined the possibility of using a nuclear isomer of hafnium (the above mentioned 178m2Hf) to construct high yield weapons with X-ray triggering mechanisms—an application of induced gamma emission, was infeasible because of its expense. See Hafnium controversy.
Care needs to be taken when machining hafnium because, like its sister metal zirconium, when hafnium is divided into fine particles, it is pyrophoric and can ignite spontaneously in air—similar to that obtained in Dragon's Breath. Compounds that contain this metal are rarely encountered by most people. The pure metal is not considered toxic, but hafnium compounds should be handled as if they were toxic because the ionic forms of metals are normally at greatest risk for toxicity, and limited animal testing has been done for hafnium compounds.
Francium is a chemical element with symbol Fr and atomic number 87. It was formerly known as eka-caesium and actinium K. It is one of the two least electronegative elements, the other being caesium. Francium is a highly radioactive metal that decays into astatine, radium, and radon. As an alkali metal, it has one valence electron.
Bulk francium has never been viewed. Because of the general appearance of the other elements in its periodic table column, it is assumed that francium would appear as a highly reflective metal, if enough could be collected together to be viewed as a bulk solid or liquid. However preparing such a sample is impossible, since the extreme heat of decay (the half-life of its longest-lived isotope is only 22 minutes) would immediately vaporize any viewable quantity of the element.
Francium was discovered by Marguerite Perey in France (from which the element takes its name) in 1939. It was the last element discovered in nature, rather than by synthesis. Outside the laboratory, francium is extremely rare, with trace amounts found in uranium and thorium ores, where the isotope francium-223 continually forms and decays. As little as 20–30 g (one ounce) exists at any given time throughout the Earth's crust; the other isotopes (except for francium-221) are entirely synthetic. The largest amount produced in the laboratory was a cluster of more than 300,000 atoms.
Francium is the most unstable of the naturally occurring elements: its most stable isotope, francium-223, has a half-life of only 22 minutes. In contrast, astatine, the second-least stable naturally occurring element, has a half-life of 8.5 hours. All isotopes of francium decay into either astatine, radium, or radon. Francium is also less stable than all synthetic elements up to element 105.
Francium is an alkali metal whose chemical properties mostly resemble those of caesium. A heavy element with a single valence electron, it has the highest equivalent weight of any element. Liquid francium—if such a substance were to be created—should have a surface tension of 0.05092 N/m at its melting point. Francium's melting point was claimed to have been calculated to be around 27 °C (80 °F, 300 K). However, the melting point is uncertain because of the element's extreme rarity and radioactivity. Thus, the estimated boiling point value of 677 °C (1250 °F, 950 K) is also uncertain.
Linus Pauling estimated the electronegativity of francium at 0.7 on the Pauling scale, the same as caesium; the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium. Francium has a slightly higher ionization energy than caesium, 392.811(4) kJ/mol as opposed to 375.7041(2) kJ/mol for caesium, as would be expected from relativistic effects, and this would imply that caesium is the less electronegative of the two.
Francium coprecipitates with several caesium salts, such as caesium perchlorate, which results in small amounts of francium perchlorate. This coprecipitation can be used to isolate francium, by adapting the radiocaesium coprecipitation method of Glendenin and Nelson. It will additionally coprecipitate with many other caesium salts, including the iodate, the picrate, the tartrate (also rubidium tartrate), the chloroplatinate, and the silicotungstate. It also coprecipitates with silicotungstic acid, and with perchloric acid, without another alkali metal as a carrier, which provides other methods of separation. Nearly all francium salts are water-soluble.
Due to its instability and rarity, there are no commercial applications for francium. It has been used for research purposes in the fields of biology and of atomic structure. Its use as a potential diagnostic aid for various cancers has also been explored, but this application has been deemed impractical.
Francium's ability to be synthesized, trapped, and cooled, along with its relatively simple atomic structure have made it the subject of specialized spectroscopy experiments. These experiments have led to more specific information regarding energy levels and the coupling constants between subatomic particles. Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels which are fairly similar to those predicted by quantum theory.
As early as 1870, chemists thought that there should be an alkali metal beyond caesium, with an atomic number of 87. It was then referred to by the provisional name eka-caesium. Research teams attempted to locate and isolate this missing element, and at least four false claims were made that the element had been found before an authentic discovery was made.
Soviet chemist D. K. Dobroserdov was the first scientist to claim to have found eka-caesium, or francium. In 1925, he observed weak radioactivity in a sample of potassium, another alkali metal, and incorrectly concluded that eka-caesium was contaminating the sample (the radioactivity from the sample was actually the naturally occurring potassium radioisotope, potassium-40). He then published a thesis on his predictions of the properties of eka-caesium, in which he named the element russium after his home country. Shortly thereafter, Dobroserdov began to focus on his teaching career at the Polytechnic Institute of Odessa, and he did not pursue the element further.
The following year, English chemists Gerald J. F. Druce and Frederick H. Loring analyzed X-ray photographs of manganese(II) sulfate. They observed spectral lines which they presumed to be of eka-caesium. They announced their discovery of element 87 and proposed the name alkalinium, as it would be the heaviest alkali metal.
In 1930, Fred Allison of the Alabama Polytechnic Institute claimed to have discovered element 87 when analyzing and lepidolite using his magneto-optical machine. Allison requested that it be named virginium after his home state of Virginia, along with the symbols Vi and Vm. In 1934, however, H.G. MacPherson of UC Berkeley disproved the effectiveness of Allison's device and the validity of this false discovery.
In 1936, Romanian physicist Horia Hulubei and his French colleague Yvette Cauchois also analyzed pollucite, this time using their high-resolution X-ray apparatus. They observed several weak emission lines, which they presumed to be those of element 87. Hulubei and Cauchois reported their discovery and proposed the name moldavium, along with the symbol Ml, after Moldavia, the Romanian province where Hulubei was born. In 1937, Hulubei's work was criticized by American physicist F. H. Hirsh Jr., who rejected Hulubei's research methods. Hirsh was certain that eka-caesium would not be found in nature, and that Hulubei had instead observed mercury or bismuth X-ray lines. Hulubei, however, insisted that his X-ray apparatus and methods were too accurate to make such a mistake. Because of this, Jean Baptiste Perrin, Nobel Prize winner and Hulubei's mentor, endorsed moldavium as the true eka-caesium over Marguerite Perey's recently discovered francium. Perey, however, continuously criticized Hulubei's work until she was credited as the sole discoverer of element 87.
Eka-caesium was discovered in 1939 by Marguerite Perey of the Curie Institute in Paris, France when she purified a sample of actinium-227 which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one which was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227. Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure which she later revised to 1%.
Perey named the new isotope actinium-K (now referred to as francium-223) and in 1946, she proposed the name catium for her newly discovered element, as she believed it to be the most electropositive cation of the elements. Irène Joliot-Curie, one of Perey's supervisors, opposed the name due to its connotation of cat rather than cation. Perey then suggested francium, after France. This name was officially adopted by the International Union of Pure and Applied Chemistry in 1949, becoming the second element after gallium to be named after France. It was assigned the symbol Fa, but this abbreviation was revised to the current Fr shortly thereafter. Francium was the last element discovered in nature, rather than synthesized, following rhenium in 1925. Further research into francium's structure was carried out by, among others, Sylvain Lieberman and his team at CERN in the 1970s and 1980s.
Francium-223 is the result of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals. In a given sample of uranium, there is estimated to be only one francium atom for every 1 × 1018 uranium atoms. It is also calculated that there is at most 30 g of francium in the Earth's crust at any time.
Francium can be synthesized in the nuclear reaction:
This process, developed by Stony Brook Physics, yields francium isotopes with masses of 209, 210, and 211, which are then isolated by the magneto-optical trap (MOT). The production rate of a particular isotope depends on the energy of the oxygen beam. An 18O beam from the Stony Brook LINAC creates 210Fr in the gold target with the nuclear reaction 197Au + 18O → 210Fr + 5n. The production required some time to develop and understand. It was critical to operate the gold target very close to its melting point and to make sure that its surface was very clean. The nuclear reaction imbeds the francium atoms deep in the gold target, and they must be removed efficiently. The atoms diffuse fast to the surface of the gold target and are released as ions; however, this does not happen every time. The francium ions are guided by electrostatic lenses until they land into a surface of hot yttrium and become neutral again. The francium is then injected into a glass bulb. A magnetic field and laser beams cool and confine the atoms. Although the atoms remain in the trap for only about 20 seconds before escaping (or decaying), a steady stream of fresh atoms replaces those lost, keeping the number of trapped atoms roughly constant for minutes or longer. Initially, about 1000 francium atoms were trapped in the experiment. This was gradually improved and the setup is capable of trapping over 300,000 neutral atoms of francium a time. Although these are neutral "metallic" atoms ("francium metal"), they are in a gaseous unconsolidated state. Enough francium is trapped that a video camera can capture the light given off by the atoms as they fluoresce. The atoms appear as a glowing sphere about 1 millimeter in diameter. This was the very first time that anyone had ever seen francium. The researchers can now make extremely sensitive measurements of the light emitted and absorbed by the trapped atoms, providing the first experimental results on various transitions between atomic energy levels in francium. Initial measurements show very good agreement between experimental values and calculations based on quantum theory. Other synthesis methods include bombarding radium with neutrons, and bombarding thorium with protons, deuterons, or helium ions. Francium has not, as of 2012[update], been synthesized in amounts large enough to weigh.
There are 34 known isotopes of francium ranging in atomic mass from 199 to 232. Francium has seven metastable nuclear isomers. Francium-223 and francium-221 are the only isotopes that occur in nature, though the former is far more common.
Francium-223 is the most stable isotope with a half-life of 21.8 minutes, and it is highly unlikely that an isotope of francium with a longer half-life will ever be discovered or synthesized. Francium-223 is the fifth product of the actinium decay series as the daughter isotope of actinium-227. Francium-223 then decays into radium-223 by beta decay (1149 keV decay energy), with a minor (0.006%) alpha decay path to astatine-219 (5.4 MeV decay energy).
Francium-221 has a half-life of 4.8 minutes. It is the ninth product of the neptunium decay series as a daughter isotope of actinium-225. Francium-221 then decays into astatine-217 by alpha decay (6.457 MeV decay energy).
The least stable ground state isotope is francium-215, with a half-life of 0.12 μs. (9.54 MeV alpha decay to astatine-211): Its metastable isomer, francium-215m, is less stable still, with a half-life of only 3.5 ns.
Atomic Number: 3
Atomic Weight: 6.941
Melting Point: 453.85 K
Boiling Point: 1615 K
Specific mass: 0.534 g/cm3
Atomic Number: 11
Atomic Weight: 22.98976928
Melting Point: 371.15 K
Boiling Point: 1156 K
Specific mass: 0.97 g/cm3
Atomic Number: 19
Atomic Weight: 39.0983
Melting Point: 336.5 K
Boiling Point: 1032 K
Specific mass: 0.86 g/cm3
Atomic Number: 37
Atomic Weight: 85.4678
Melting Point: 312.79 K
Boiling Point: 961 K
Specific mass: 1.53 g/cm3
Atomic Number: 55
Atomic Weight: 132.9054519
Melting Point: 301.7 K
Boiling Point: 944 K
Specific mass: 1.93 g/cm3
Atomic Number: 87
Atomic Weight: 
Melting Point: 300.15 K
Boiling Point: 950 K
Specific mass: 1.87 g/cm3
Mendelevium is a synthetic element with the symbol Md (formerly Mv) and the atomic number 101. A metallic radioactive transuranic element in the actinide series, mendelevium is usually synthesized by bombarding einsteinium with alpha particles. It was named after Dmitri Ivanovich Mendeleev, who created the Periodic Table. Mendeleev's periodic system is the fundamental way to classify all the chemical elements. The name "mendelevium" was accepted by the International Union of Pure and Applied Chemistry (IUPAC) in 1955 with symbol "Mv", which was changed to "Md" in the next IUPAC General Assembly (Paris, 1957).
Researchers have shown that mendelevium has a moderately stable dipositive (II) oxidation state in addition to the more characteristic (for actinide elements) tripositive (III) oxidation state, the latter being the more dominantly exhibited state in an aqueous solution (chromatography being the process used). Sometimes, mendelevium can even be shown to exhibit a monopositive (I) state. 256Md has been used to find out some of the chemical properties of this element while in an aqueous solution. There are no other known uses of mendelevium and only trace amounts of the element have ever been produced. Other isotopes of mendelevium, all radioactive, have been discovered, with 258Md being the most stable with a two-month half-life (about 55 days). Other isotopes range from 248 to 258 mass numbers and half-lives from a few seconds to about 51 days. The original 256Md had a half-life of 87 minutes.
The trivalent element is radioactive. It was expected that the reaction would be253Es (α,n) 255Md, where 255Md was α-active with a t½ of 5 minutes and the corresponding α-energy. No such α-activity was observed, but the 101 fraction showed spontaneous fission representing a t½ less than 3 hours. Because spontaneous fission was also observed in the fraction containing fermium, the α-bombardment of 253Es produced 256Md. The latter underwent electron capture to become 256Fm, which then decayed by spontaneous fission. So 256Fm was produced by the decays of cyclotron-synthesized mendelevium.
Johansson and Rosengren predicted in 1975 that Md would prefer a divalent metallic state, similar to europium (Eu) and ytterbium (Yb), rather than a trivalent one. Thermochromatographic studies conducted with trace amounts of Md concluded that Md forms a divalent metal. With the aid of empirical correlation method, a divalent metallic radius of (0.194 ± 0.010) nm has been estimated. The estimated enthalpy of sublimation is in the range of 134-142 kJ/mol.
Before the actual discovery of mendelevium, the trivalent state was the most stable one in aqueous solution. Accordingly, a similar chemical behavior to the other 3+ actinides and lanthanides was expected. The elution of Md just before Fm in the elution sequence of the trivalent actinides from the cation-exchange resin column, confirmed this prediction. Afterwards, Md in the form of insoluble hydroxides and fluorides that are quantitatively coprecipitated with trivalent lanthanides was found. The cation-exchange resin column as well as the HDEHP solvent extraction column elution date is consistent with a trivalent state for Md and an ionic radius smaller than Fm. An ionic radius of 0.0192 nm and a coordination number of 6 for Md3+ was predicted using empirical correlations. Using the known ionic radii for the trivalent rare earths and the linear correlation of log distribution coefficient with ionic radius, an average ionic radius of 0.089 nm was estimated for Md3+ and a heat of hydration of –(3654 ± 12) kJ/mol calculated using empirical models and the Born-Haber cycle. In reducing conditions, an anomalous chemical behavior of Md was found. Coprecipitation with BaSO4 and solvent extraction chromatography experiments using HDEHP were carried out in different reducing agents. These showed that Md3+ could easily be reduced to a stable Md2+ in aqueous solution. Mendelevium can also be reduced to the monovalent state in water-ethanol solutions. The cocrystallization of Md+ with salts of divalent ions is due to the formation of mixed crystals. For Md+, an ionic radius of 0.117 nm was found. The oxidation of Md3+ to Md4+ was rather unsuccessful.
Mendelevium (for Dimitri Ivanovich Mendeleev, surname commonly transliterated into Latin script as Mendeleev, Mendeleyev, Mendeléef, or even Mendelejeff, and first name sometimes transliterated as Dmitry or Dmitriy) was first synthesized by Albert Ghiorso, Glenn T. Seaborg, Gregory R. Choppin, Bernard G. Harvey, and Stanley G. Thompson (team leader) in early 1955 at the University of California, Berkeley. The team produced 256Md (half-life of 87 minutes) when they bombarded an 253Es target with alpha particles (helium nuclei) in the Berkeley Radiation Laboratory's 60-inch cyclotron (256Md was the first isotope of any element to be synthesized one atom at a time). Element 101 was the ninth transuranic element synthesized. The first 17 atoms of this element were created and analyzed using the ion-exchange adsorption-elution method. During the process, mendelevium behaved very much like thulium, its naturally occurring homologue.
The discovery was based on a grand total of only 17 atoms. It is synthesized via the 253Es (α,n) 256101 reaction in the 60-Inch-Cyclotron (= 152 cm) at Berkeley (California). The target can be produced by irradiation of lighter isotopes as plutonium in the Materials Testing Reactor at the Arco Reactor Station in Idaho. Remarkable is that this target consisted of only 109 atoms of highly radioactive 253Es (with a half-life of 20.5 days). By elution through a calibrated cation exchange resin column, mendelevium was separated and chemically identified.
To predict if this method would be possible, they made use of a rough calculation. The number of atoms that would be produced, would be approximately equal to the number of atoms of target material times its cross section times the ion beam intensity times the time of bombardment related to the half-life of the product when bombarding for a time of the order of its half-life). This gave 1 atom per experiment. Thus under optimum conditions, the preparation of only one atom of element 101 per experiment could be expected. This calculation demonstrated that it was feasible to go ahead with the experiment.
The actual synthesis was done by a recoil technique, introduced by Albert Ghiorso. In this technique, the target element was placed on the opposite side of the target from the beam and caught the recoiling atoms on a catcher foil. This recoil target was made by an electroplating technique, developed by Alfred Chetham-Strode. This technique gave a very high yield, which is absolutely necessary when working with such a rare product as the einsteinium target material.
The recoil target consisted of 10−9 of 253Es which were deposited electrolytically on a thin gold foil (also Be, Al and Pt can be used). It was bombarded by 41 eV α-particles in the Berkeley cyclotron with a very high beam density of 6∙1013 particles per second over an area of 0.05 cm2. The target was cooled by water or liquid helium. The use of helium, in a gaseous atmosphere, slowed down the recoil atoms. This gas could be pumped out of the reaction chamber through a small orifice to form a ‘gas-jet’. Some fraction of the nonvolatile product atoms carried along with the gas, were deposited permanently on the foil surface. The foil could be removed periodically and a new foil could be installed. The next reaction was used for the mendelevium discovery experiment: 253Es + 4He → 256Md + 1n.
The removal of the Md atoms from the collector foil was done by acid etching or total dissolution of the thin gold foil. They can be purified and isolated from other product activities by several techniques. Separation of trivalent actinides from lanthanide fission products and La carrier can be done by a cation-exchange resin column using a 90% water/10% ethanol solution saturated with HCl as eluant. To separate Md rapidly from the catcher foil, an anion-exchange chromatography using 6M HCl as eluant can be used. The gold remained on the column while the Md and other actinides passed through. A final isolation of Md3+ from other trivalent actinides was also required. To separate fractions containing elements 99, 100 and 101, a cation-exchange resin column (Dowex-50 exchange column) treated with ammonium salts was used. A chemical identification was made on the basis of its elution position just before Fm. In series of repetitive experiments, they made use of the eluant: α-hydroxyisobutyrate solution (α-HIB). Using the ‘gas-jet’ method, the first two steps can be eliminated. There was shown that in this method it is possible to transport and collect individual product atoms in a fraction of a second some tens of meters away from the target area. Effective transport over long distances requires the presence of large clusters (KCl aerosols) in the ‘carrier’ gas. It is used frequently in the production and isolation of transeinsteinium elements.
Another possible way to separate the 3+ actinides can be achieved by solvent extraction chromatography using bis-(2-ethylhexyl) phosphoric acid (abbreviated as HDEHP) as the stationary organic phase and HNO3 as the mobile aqueous phase. The actinide elution sequence is reversed from that of the cation-exchange resin column. The Md separated by this method has the advantage to be free of organic complexing agent compared to the resin column. The disadvantage of this method is that Md elutes after Fm late in the sequence.
There was no direct detection, but by observation of spontaneous fission events arising from its electron-capture daughter 256Fm. These events were recorded during the night of February 19, 1955. The first one was identified with a "hooray" followed by a "double hooray" and a "triple hooray". The fourth one eventually officially proved the chemical identification of the 101st element, mendelevium. Additional analysis and further experimentation, showed the isotope to have mass 256 and to decay by electron capture with a half-life of 1.5 h.
Sixteen isotopes of mendelevium from mass 245 to 260 have been characterized, with the most stable being 258Md with a half-life of 51.5 days, 260Md with a half-life of 31.8 days, and 257Md with a half-life of 5.52 hours. All of the remaining radioactive isotopes have half-lives that are less than 97 minutes, and the majority of these have half-lives that are less than 5 minutes. This element also has 5 meta states, with the longest-lived being 258mMd (t½ = 58 minutes). The isotopes of mendelevium range in atomic weight from 245.091 u (245Md) to 260.104 u (260Md).
Lawrencium is a radioactive synthetic chemical element with the symbol Lr (formerly Lw) and atomic number 103. In the periodic table of the elements, it is a period 7 d-block element and the last element of the actinide series. Chemistry experiments have confirmed that lawrencium behaves as the heavier homologue to lutetium and is chemically similar to other actinides.
Lawrencium was first synthesized by the nuclear-physics team led by Albert Ghiorso on February 14, 1961, at the Lawrence Berkeley National Laboratory of the University of California. The first atoms of lawrencium were produced by bombarding a three-milligram target consisting of three isotopes of the element californium with boron-10 and boron-11 nuclei from the Heavy Ion Linear Accelerator. The team suggested the name lawrencium (after Ernest Lawrence), and the symbol "Lw", but IUPAC changed the symbol to "Lr" in 1963. It was the final and the heaviest element of the actinide series to be synthesized.
All isotopes of lawrencium are radioactive; its most stable known isotope is lawrencium-262, with a half-life of approximately 3.6 hours. All its isotopes except for lawrencium-260, -261 and -262 decay with a half-life of less than a minute.
Lawrencium was first synthesized by the nuclear-physics team of Albert Ghiorso, Torbjørn Sikkeland, Almon Larsh, Robert M. Latimer, and their co-workers on February 14, 1961, at the Lawrence Radiation Laboratory (now called the Lawrence Berkeley National Laboratory) at the University of California. The first atoms of lawrencium were produced by bombarding a three-milligram target consisting of three isotopes of the element californium with boron-10 and boron-11 nuclei from the Heavy Ion Linear Accelerator (HILAC). The Berkeley team reported that the isotope 257Lr was detected in this manner, and that it decayed by emitting an 8.6 MeV alpha particle with a half-life of about eight seconds. This identification was later corrected to be 258Lr.
In 1967, nuclear-physics researchers in Dubna, Russia, reported that they were not able to confirm assignment of an alpha emitter with a half-life of eight seconds to 257Lr. This isotope was later deduced to be 258Lr. Instead, the Dubna team reported an isotope with a half-life of about 45 seconds as 256Lr.
Further experiments have demonstrated an actinide chemistry for the new element, so by 1970 it was known that lawrencium is the last actinide. In 1971, the nuclear physics team at the University of California at Berkeley successfully performed a whole series of experiments aimed at measuring the nuclear decay properties of the lawrencium isotopes with mass numbers from 255 through 260.
In 1992, the IUPAC Trans-fermium Working Group (TWG) officially recognized the nuclear physics teams at Dubna and Berkeley as the co-discoverers of lawrencium.
The origin of the name, ratified by the American Chemical Society, is in reference to the nuclear-physicist Ernest O. Lawrence, of the University of California, who invented the cyclotron particle accelerator. The symbol Lw was used originally, but the element was assigned the Lr symbol. In August 1997, the International Union of Pure and Applied Chemistry (IUPAC) ratified the name lawrencium and the symbol Lr during a meeting in Geneva.
Lawrencium is element 103 in the periodic table. It is the first member of the 6d-block; in accordance with the Madelung rule, its electronic configuration should be [Rn]7s25f146d1. However, results from quantum mechanical research have suggested that this configuration is incorrect, and is in fact [Rn]7s25f147p1. A direct measurement of this is not possible. Though early calculations gave conflicting results, more recent studies and calculations confirm the suggestion.
A strict correlation between the periodic table blocks and the orbital-shell configurations for neutral atoms would classify lawrencium as a transition metal because it could be classed as a d-block element. However, lawrencium is classified as an actinide element according to the IUPAC recommendations.
The first gaseous-phase studies of lawrencium were reported in 1969 by a nuclear physics team at the Flerov Laboratory of Nuclear Reactions (FLNR) in the Soviet Union. They used the nuclear reaction 243Am+18O to produce lawrencium nuclei, which they then exposed to a stream of chlorine gas, and a volatile chloride product was formed. This product was deduced to be 256LrCl3, and this confirmed that lawrencium is a typical actinide element.
The first aqueous-phase studies of lawrencium were reported in 1970 by a nuclear physics team at the Lawrence Berkeley National Laboratory in California. This team used the nuclear reaction 249Cf+11B to produce lawrencium nuclei. They were able to show that lawrencium forms a trivalent ion, similar to those of the other actinide elements, but in contrast with that of nobelium. Further experiments in 1988 confirmed the formation of a trivalent lawrencium(III) ion using anion-exchange chromatography using α-hydroxyisobutyrate (α-HIB) complex. Comparison of the elution time with other actinides allowed a determination of 88.6 ± 0.3 picometers for the ionic radius for Lr3+. Attempts to reduce lawrencium in the lawrencium(III) ionization state to lawrencium(I) using the potent reducing agent hydroxylamine hydrochloride were unsuccessful.
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR. Evidence was provided for the formation of 253Lr in the 2n exit channel.
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR.
This reaction was reported in 1984 by Yuri Oganessian at the FLNR. The team was able to detect decays of 246Cf, a descendant of 254Lr.
This reaction was studied in a series of experiments in 1976 by Yuri Oganessian and his team at the FLNR. Results are not readily available.
This reaction has been used to study the spectroscopic properties of 255Lr. The team at GANIL used the reaction in 2003 and the team at the FLNR used it between 2004 and 2006 to provide further information for the decay scheme of 255Lr. The work provided evidence for an isomeric level in 255Lr.
This reaction was first studied in 1965 by the team at the FLNR. They were able to detect activity with a characteristic decay of 45 seconds, which was assigned to 256Lr or 257Lr. Later work suggests an assignment to 256Lr. Further studies in 1968 produced an 8.35–8.60 MeV alpha activity with a half-life of 35 seconds. This activity was also initially assigned to 256Lr or 257Lr and later to solely 256Lr.
This reaction was studied in 1970 by the team at the FLNR. They were able to detect an 8.37 MeV alpha activity with a half-life of 22s. This was assigned to 255Lr.
This reaction was studied in 1971 by the team at the LBNL in their large study of lawrencium isotopes. They were able to assign alpha activities to 260Lr,259Lr and 258Lr from the 3-5n exit channels.
This reaction was studied in 1988 at the LBNL in order to assess the possibility of producing 262Lr and 261Lr without using the exotic 254Es target. It was also used to attempt to measure an electron capture (EC) branch in 261mRf from the 5n exit channel. After extraction of the Lr(III) component, they were able to measure the spontaneous fission of 261Lr with an improved half-life of 44 minutes. The production cross-section was 700 pb. On this basis, a 14% electron capture branch was calculated if this isotope was produced via the 5n channel rather than the p4n channel. A lower bombarding energy (93 MeV c.f. 97 MeV) was then used to measure the production of 262Lr in the p3n channel. The isotope was successfully detected and a yield of 240 pb was measured. The yield was lower than expected compared to the p4n channel. However, the results were judged to indicate that the 261Lr was most likely produced by a p3n channel and an upper limit of 14% for the electron capture branch of 261mRf was therefore suggested.
This reaction was studied briefly in 1958 at the LBNL using an enriched 244Cm target (5% 246Cm). They observed a ~9 MeV alpha activity with a half-life of ~0.25 seconds. Later results suggest a tentative assignment to 257Lr from the 3n channel
This reaction was studied briefly in 1958 at the LBNL using an enriched 244Cm target (5% 246Cm). They observed a ~9 MeV alpha activity with a half-life of ~0.25s. Later results suggest a tentative assignment to 257Lr from the 3n channel with the 246Cm component. No activities assigned to reaction with the 244Cm component have been reported.
This reaction was studied in 1971 by the team at the LBNL in their large study of lawrencium isotopes. They were able to detect an activity assigned to 260Lr. The reaction was repeated in 1988 to study the aqueous chemistry of lawrencium. A total of 23 alpha decays were measured for 260Lr, with a mean energy of 8.03 MeV and an improved half-life of 2.7 minutes. The calculated cross-section was 8.7 nb.
This reaction was first studied in 1961 at the University of California by Albert Ghiorso by using a californium target (52% 252Cf). They observed three alpha activities of 8.6, 8.4 and 8.2 MeV, with half-lives of about 8 and 15 seconds, respectively. The 8.6 MeV activity was tentatively assigned to 257Lr. Later results suggest a reassignment to 258Lr, resulting from the 5n exit channel. The 8.4 MeV activity was also assigned to 257Lr. Later results suggest a reassignment to 256Lr. This is most likely from the 33% 250Cf component in the target rather than from the 7n channel. The 8.2 MeV was subsequently associated with nobelium.
This reaction was first studied in 1961 at the University of California by Albert Ghiorso by using a californium target (52% 252Cf). They observed three alpha activities of 8.6, 8.4 and 8.2 MeV, with half-lives of about 8 and 15 seconds, respectively. The 8.6 MeV activity was tentatively assigned to 257Lr. Later results suggest a reassignment to 258Lr. The 8.4 MeV activity was also assigned to 257Lr. Later results suggest a reassignment to 256Lr. The 8.2 MeV was subsequently associated with nobelium.
This reaction was studied in 1971 at the LBNL. They were able to identify a 0.7 s alpha activity with two alpha lines at 8.87 and 8.82 MeV. This was assigned to 257Lr.
This reaction was first studied in 1970 at the LBNL in an attempt to study the aqueous chemistry of lawrencium. They were able to measure a Lr3+ activity. The reaction was repeated in 1976 at Oak Ridge and the 26s lifetime of 256Lr was confirmed by measurement of coincident X-rays.
This reaction was studied in 1971 by the team at the LBNL. They were able to detect an activity assigned to 258Lr from the p2n channel.
This reaction was studied in 1971 by the team at the LBNL. They were able to detect an activities assigned to 258Lr and 257Lr from the α2n and α3n and channels. The reaction was repeated in 1976 at Oak Ridge and the synthesis of 258Lr was confirmed.
This reaction was studied in 1987 at the LLNL. They were able to detect new spontaneous fission (SF) activities assigned to 261Lr and 262Lr, resulting from transfer from the 22Ne nuclei to the 254Es target. In addition, a 5 ms SF activity was detected in delayed coincidence with nobelium K-shell X-rays and was assigned to 262No, resulting from the electron capture of 262Lr.
Isotopes of lawrencium have also been identified in the decay of heavier elements. Observations to date are summarised in the table below:
Eleven isotopes of lawrencium plus one isomer have been synthesized with 262Lr being the longest-lived and the heaviest, with a half-life of 216 minutes. 252Lr is the lightest isotope of lawrencium to be produced to date.
A study of the decay properties of 257Db (see dubnium) in 2001 by Hessberger et al. at the GSI provided some data for the decay of 253Lr. Analysis of the data indicated the population of two isomeric levels in 253Lr from the decay of the corresponding isomers in 257Db. The ground state was assigned spin and parity of 7/2-, decaying by emission of an 8794 KeV alpha particle with a half-life of 0.57s. The isomeric level was assigned spin and parity of 1/2-, decaying by emission of an 8722 KeV alpha particle with a half-life of 1.49 s.
Recent work on the spectroscopy of 255Lr formed in the reaction 209Bi(48Ca,2n)255Lr has provided evidence for an isomeric level.
Promethium, originally prometheum, is a chemical element with the symbol Pm and atomic number 61. All of its isotopes are radioactive; it is one of only two such elements that are followed in the periodic table by elements with stable forms, a distinction shared with technetium. Chemically, promethium is a lanthanide, which forms salts when combined with other elements. Promethium shows only one stable oxidation state of +3; however, a few +2 compounds may exist.
In 1902, Bohuslav Brauner suggested there was an element with properties intermediate between those of the known elements neodymium (60) and samarium (62); this was confirmed in 1914 by Henry Moseley who, having measured the atomic numbers of all the elements then known, found there was no element with atomic number 61. In 1926, an Italian and an American group claimed to have isolated a sample of element 61; both "discoveries" were soon proven to be false. In 1938, during a nuclear experiment conducted at Ohio State University, a few radioactive nuclides were produced that certainly were not radioisotopes of neodymium or samarium, but there was a lack of chemical proof that element 61 was produced, and the discovery was not generally recognized. Promethium was first produced and characterized at Oak Ridge National Laboratory in 1945 by the separation and analysis of the fission products of uranium fuel irradiated in a graphite reactor. The discoverers proposed the name "prometheum" (the spelling was subsequently changed), derived from Prometheus, the Titan in Greek mythology who stole fire from Mount Olympus and brought it down to humans, to symbolize "both the daring and the possible misuse of mankind's intellect." However, a sample of the metal was made only in 1963.
There are two possible sources for natural promethium: rare decays of natural europium-151 (producing promethium-147), and uranium (various isotopes). Practical applications exist only for chemical compounds of promethium-147, which are used in luminous paint, atomic batteries and thickness measurement devices, even though promethium-145 is the most stable promethium isotope. Since natural promethium is exceedingly scarce, the element is typically synthesized by bombarding uranium-235 (enriched uranium) with thermal neutrons to produce promethium-147.
A promethium atom has 61 electrons, arranged in the configuration [Xe]4f56s2. In forming compounds, the atom loses its two outermost electrons and one of the 4f-electrons, which belongs to an open subshell. The element's atomic radius is the third largest among all the lanthanides but is only slightly greater than those of the neighboring elements. It is the only exception to the general trend of the contraction of the atoms with increase of atomic radius (caused by the lanthanide contraction) that is not caused by the filled (or half-filled) 4f-subshell.
Many properties of promethium rely on its position among lanthanides and are intermediate between those of neodymium and samarium. For example, the melting point, the first three ionization energies, and the hydration energy are greater than those of neodymium and lower than those of samarium; similarly, the estimate for the boiling point, ionic (Pm3+) radius, and standard heat of formation of monatomic gas are greater than those of samarium and less those of neodymium.
Promethium has a double hexagonal close packed (dhcp) structure and a hardness of 63 kg/mm2. This low-temperature alpha form converts into a beta, body-centered cubic (bcc) phase upon heating to 890 °C.
Promethium belongs to the cerium group of lanthanides and is chemically very similar to the neighboring elements. Because of its instability, chemical studies of promethium are incomplete. Even though a few compounds have been synthesized, they are not fully studied; in general, they tend to be pink or red in color. Treatment of acidic solutions containing Pm3+ ions with ammonia results in a gelatinous light-brown sediment of hydroxide, Pm(OH)3, which is insoluble in water. When dissolved in hydrochloric acid, a water-soluble yellow salt, PmCl3, is produced; similarly, when dissolved in nitric acid, a nitrate results, Pm(NO3)3. The latter is also well-soluble; when dried, it forms pink crystals, similar to Nd(NO3)3. The electron configuration for Pm3+ is [Xe] 4f4, and the color of the ion is pink. The ground state term symbol is 5I4. The sulfate is slightly soluble, like the other cerium group sulfates. Cell parameters have been calculated for its octahydrate; they lead to conclusion that the density of Pm2(SO4)3·8 H2O is 2.86 g/cm3. The oxalate, Pm2(C2O4)3·10 H2O, has the lowest solubility of all lanthanide oxalates.
Unlike the nitrate, the oxide is similar to the corresponding samarium salt and not the neodymium salt. As-synthesized, e.g. by heating the oxalate, it is a white or lavender-colored powder with disordered structure. This powder crystallizes in a cubic lattice upon heating to 600 °C. Further annealing at 800 °C and then at 1750 °C irreversibly transforms it to a monoclinic and hexagonal phases, respectively, and the last two phases can be interconverted by adjusting the annealing time and temperature.
Promethium forms only one stable oxidation state, +3, in the form of ions; this is in line with other lanthanides. According to its position in the periodic table, the element cannot be expected to form stable +4 or +2 oxidation states; treating chemical compounds containing Pm3+ ions with strong oxidizing or reducing agents showed that the ion is not easily oxidized or reduced.
Promethium is the only lanthanide and one of only two elements among the first 82 that has no stable (or even long-lived) isotopes; this is a result of a rarely occurring effect of the liquid drop model and stabilities of neighbor element isotopes; it is also the least stable element of the first 84. The primary decay products are neodymium and samarium isotopes (promethium-146 decays to both, the lighter isotopes generally to neodymium via positron decay and electron capture, and the heavier isotopes to samarium via beta decay). Promethium nuclear isomers may decay to other promethium isotopes and one isotope (145Pm) has a very rare alpha decay mode to praseodymium.
The most stable isotope of the element is promethium-145, which has a specific activity of 940 Ci (35 TBq)/g and a half-life of 17.7 years via electron capture. Because it has 84 neutrons (two more than 82, which is a magic number which corresponds to a stable neutron configuration), it may emit an alpha particle (which has 2 neutrons) to form praseodymium-141 with 82 neutrons. Thus it is the only promethium isotope with an experimentally observed alpha decay. Its partial half-life for alpha decay is about 6.3109 years, and the relative probability for a 145Pm nucleus to decay in this way is 2.810−7%. Several other Pm isotopes (144Pm, 146Pm, 147Pm etc.) also have a positive energy release for alpha decay; their alpha decays are predicted to occur but have not been observed.
The element also has 18 nuclear isomers, with mass numbers of 133 to 142, 144, 148, 149, 152, and 154 (some mass numbers have more than one isomer). The most stable of them is promethium-148m, with a half-life of 43.1 days; this is longer than the half-lives of the ground states of all promethium isotopes, except only for promethium-143 to 147 (note that promethium-148m has a longer half-life than the ground state, promethium-148).
In 1934, Willard Libby found weak beta activity in pure neodymium, which was attributed to a half-life over 1012 years. Almost 20 years later, it was claimed that the element occurs in natural neodymium in equilibrium in quantities below 10−20 grams of promethium per one gram of neodymium. However, these observations were disproved by newer investigations, because for all seven naturally occurring neodymium isotopes, any single beta decays (which can produce promethium nuclides) are forbidden by energy conservation. In particular, careful measurements of atomic masses show that the mass difference 150Nd-150Pm is negative (−87 keV), which absolutely prevents the single beta decay of 150Nd to 150Pm.
Both isotopes of natural europium have larger mass excesses than sums of those of their potential alpha daughters plus that of an alpha particle; therefore, they (stable in practice) may alpha decay. Research at Laboratori Nazionali del Gran Sasso showed that europium-151 experimentally decays to promethium-147 with the half-life of 51018 years. It has been shown that europium is "responsible" for about 12 grams of promethium in the Earth's crust. Alpha decays for europium-153 have not been found yet, and its theoretically calculated half-life is so high (due to low energy of decay) that this process will probably never be observed.
Finally, promethium can be formed in nature as a product of spontaneous fission of uranium-238. Only trace amounts can be found in naturally occurring ores: a sample of pitchblende has been found to contain promethium at a concentration of four parts per quintillion (1018) by mass. Uranium is thus "responsible" for 560 g promethium in Earth's crust.
Promethium has also been identified in the spectrum of the star HR 465 in Andromeda; it also has been found in HD 101065 (Przybylski's star) and HD 965. Because of the short half-life of promethium isotopes, they should be formed near the surface of those stars.
In 1902, Czech chemist Bohuslav Brauner found out that the difference between neodymium and samarium is the largest of all neighboring lanthanides pairs; as a conclusion, he suggested there was an element with intermediate properties between them. This prediction was supported in 1914 by Henry Moseley who, having discovered that atomic number was an experimentally measurable property of elements, found a few atomic numbers had no element to correspond: the gaps were 43, 61, 72, 75, 85, and 87. With the knowledge of a gap in the periodic table several groups started to search for the predicted element among other rare earths in the natural environment.
The first claim of a discovery was published by Luigi Rolla and Lorenzo Fernandes of Florence, Italy. After separating a mixture of a few rare earth elements nitrate concentrate from the Brazilian mineral monazite by fractionated crystallization, they yielded a solution containing mostly samarium. This solution gave x-ray spectra attributed to samarium and element 61. In honor of their city, they named element 61 "florentium." The results were published in 1926, but the scientists claimed that the experiments were done in 1924. Also in 1926, a group of scientists from the University of Illinois at Urbana-Champaign, Smith Hopkins and Len Yntema published the discovery of element 61. They named it "illinium," after the university. Both of these reported discoveries were shown to be erroneous because the spectrum line that "corresponded" to element 61 was identical to that of didymium; the lines thought to belong to element 61 turned out to belong to a few impurities (barium, chromium, and platinum).
In 1934, Josef Mattauch finally formulated the isobar rule. One of the indirect consequences of was this rule was that element 61 was unable to form stable isotopes. In 1938, a nuclear experiment was conducted by H. B. Law et al. at Ohio State University. The nuclides produced certainly were not radioisotopes of neodymium or samarium, and the name "cyclonium" was proposed, but there was a lack of chemical proof that element 61 was produced and the discovery not largely recognized.
Promethium was first produced and characterized at Oak Ridge National Laboratory (Clinton Laboratories at that time) in 1945 by Jacob A. Marinsky, Lawrence E. Glendenin and Charles D. Coryell by separation and analysis of the fission products of uranium fuel irradiated in the graphite reactor; however, being too busy with military-related research during World War II, they did not announce their discovery until 1947. The original proposed name was "clintonium", after the laboratory where the work was conducted; however, the name "prometheum" was suggested by Grace Mary Coryell, the wife of one of the discoverers. It is derived from Prometheus, the Titan in Greek mythology who stole fire from Mount Olympus and brought it down to humans and symbolizes "both the daring and the possible misuse of the mankind intellect." The spelling was then changed to "promethium," as this was in closer in accordance with other metals.
Jacob A. Marinsky
Lawrence E. Glendenin
Charles D. Coryell
In 1963, promethium(III) fluoride was used to make promethium metal. Provisionally purified from impurities of samarium, neodymium, and americium, it was put into a tantalum crucible which was located in another tantalum crucible; the outer crucible contained lithium metal (10 times excess compared to promethium). After creating a vacuum, the chemicals were mixed to produce promethium metal:
The promethium sample produced was used to measure a few of the metal's properties, such as its melting point.
In 1963, ion-exchange methods were used at ORNL to prepare about ten grams of promethium from nuclear reactor fuel processing wastes.
Today, promethium is still recovered from the byproducts of uranium fission; it can also be produced by bombarding 146Nd with neutrons, turning it into 147Nd which decays into Pm147 through beta decay with a half-life of 11 days.
The production methods for different isotopes vary, and only that for promethium-147 is given because it is the only isotope with industrial applications. Promethium-147 is produced in large quantities (compared to other isotopes) by bombarding uranium-235 with thermal neutrons. The output is relatively high, at 2.6% of the total product. Another way to produce promethium-147 is via neodymium-147, which decays to promethium-147 with a short half-life. Neodymium-147 can be obtained either by bombarding enriched neodymium-146 with thermal neutrons or by bombarding a uranium carbide target with energetic protons in a particle accelerator. Another method is to bombard uranium-238 with fast neutrons to cause fast fission, which, among multiple reaction products, creates promethium-147.
As early as the 1960s, Oak Ridge National Laboratory could produce 650 grams of promethium per year and was the world's only large-volume synthesis facility. Gram-scale production of promethium has been discontinued in the U.S. in the early 1980s, but will possibly be resumed after 2010 at the High Flux Isotope Reactor. Currently, Russia is the only country producing promethium-147 on a relatively large scale.
Most promethium is used only for research purposes, except for promethium-147, which can be found outside laboratories. It is obtained as the oxide or chloride, in milligram quantities. This isotope does not emit gamma rays, and its radiation has a relatively small penetration depth in matter and a relatively long half-life.
Some signal lights use a luminous paint, containing a phosphor that absorbs the beta radiation emitted by promethium-147 and emits light. This isotope does not cause aging of the phosphor, as alpha emitters do, and therefore the light emission is stable for a few years. Originally, radium-226 was used for the purpose, but it was later replaced by promethium-147 and tritium (hydrogen-3). Promethium may be favored over tritium for safety reasons.
In atomic batteries, the beta particles emitted by promethium-147 are converted into electric current by sandwiching a small Pm source between two semiconductor plates. These batteries have a useful lifetime of about five years. The first promethium-based battery was assembled in 1964 and generated "a few milliwatts of power from a volume of about 2 cubic inches, including shielding".
Promethium is also used to measure the thickness of materials by evaluating the amount of radiation from a promethium source that passes through the sample. It has possible future uses in portable X-ray sources, and as auxiliary heat or power sources for space probes and satellites (although the alpha emitter plutonium-238 has become standard for most space-exploration-related uses).
The element, like other lanthanides, has no biological role. Promethium-147 can emit X-rays during its beta decay, which are dangerous for all lifeforms. Interactions with tiny quantities of promethium-147 are not hazardous if certain precautions are observed. In general, gloves, footwear covers, safety glasses, and an outer layer of easily removed protective clothing should be used.
It is not known what human organs are affected by interaction with promethium; a possible candidate is the bone tissues. Sealed promethium-147 is not dangerous. However, if the packaging is damaged, then promethium becomes dangerous to the environment and humans. If radioactive contamination is found, the contaminated area should be washed with water and soap, but, even though promethium mainly affects the skin, the skin should not be abraded. If a promethium leak is found, the area should be identified as hazardous and evacuated, and emergency services must be contacted. No dangers from promethium aside from the radioactivity are known.
A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus. Elements are divided into metals, metalloids, and non-metals. Familiar examples of elements include carbon, oxygen (non-metals), silicon, arsenic (metalloids), aluminium, iron, copper, gold, mercury, and lead (metals).
The lightest chemical elements, including hydrogen, helium (and smaller amounts of lithium, beryllium and boron), are thought to have been produced by various cosmic processes during the Big Bang and cosmic-ray spallation. Production of heavier elements, from carbon to the very heaviest elements, proceeded by stellar nucleosynthesis, and these were made available for later solar system and planetary formation by planetary nebulae and supernovae, which blast these elements into space. The high abundance of oxygen, silicon, and iron on Earth reflects their common production in such stars, after the lighter gaseous elements and their compounds have been subtracted. While most elements are generally viewed as stable, a small amount of natural transformation of one element to another also occurs at the present time through decay of radioactive elements as well as other natural nuclear processes.
The periodic table is a tabular arrangement of the chemical elements, organized on the basis of their atomic numbers, electron configurations, and recurring chemical properties. Elements are presented in order of increasing atomic number (the number of protons in the nucleus). The standard form of the table consists of a grid of elements laid out in 18 columns and 7 rows, with a double row of elements below that. The table can also be deconstructed into four rectangular blocks: the s-block to the left, the p-block to the right, the d-block in the middle, and the f-block below that.
The rows of the table are called periods; the columns are called groups, with some of these having names such as halogens or noble gases. Since, by definition, a periodic table incorporates recurring trends, any such table can be used to derive relationships between the properties of the elements and predict the properties of new, yet to be discovered or synthesized, elements. As a result, a periodic table—whether in the standard form or some other variant—provides a useful framework for analyzing chemical behavior, and such tables are widely used in chemistry and other sciences.
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