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Beryllium

Beryllium

Beryllium is the chemical element in the periodic table that has the symbol Be and atomic number 4. A toxic bivalent element, beryllium is a steel grey, strong, light-weight yet brittle, alkaline earth metal, that is primarily used as a hardening agent in alloys (most notably, beryllium copper).

Notable characteristics

Beryllium has one of the highest melting points of the light metals. The modulus of elasticity of beryllium is approximately 1/3 greater than that of steel. It has excellent thermal conductivity, is nonmagnetic and resists attack by concentrated nitric acid. It is highly permeable to X-rays, and neutrons are liberated when it is hit by alpha particles, as from radium or polonium (about 30 neutrons/million alpha particles). At standard temperature and pressures beryllium resists oxidation when exposed to air (although its ability to scratch glass is probably due to the formation of a thin layer of the oxide).

Applications

- Beryllium is used as an alloying agent in the production of beryllium copper. (Be has the ability to absorb large amounts of heat.) Beryllium-copper alloys are used in a wide variety of applications because of their electrical and thermal conductivity, high strength and hardness, nonmagnetic properties, along with good corrosion and fatigue resistance. These applications include the making of: spot-welding electrodes, springs, non-sparking tools and electrical contacts.
- Due to their stiffness, light weight, and dimensional stability over a wide temperature range, beryllium-copper alloys are also used in the defense and aerospace industries as light-weight structural materials in high-speed aircraft, missiles, space vehicles and communication satellites.
- Thin sheets of beryllium foil are used with X-ray detection diagnostics to filter out visible light and allow only X-rays to be detected.
- In the field of X-ray lithography beryllium is being used for the reproduction of microscopic integrated circuits.
- Because it has a low thermal neutron absorption cross section, the nuclear power industry uses this metal in nuclear reactors as a neutron reflector and moderator.
- Beryllium is sometimes used in neutron sources, in which the beryllium is mixed with an alpha emmiter such as polonium 210, radium 226, or actinium 227.
- Beryllium is also used in the making of gyroscopes, various computer equipment, watch springs and instruments where light-weight, rigidity and dimensional stability are needed.
- Beryllium oxide is useful for many applications that require an excellent heat conductor, with high strength and hardness, with a very high melting point, and that acts as an electrical insulator.
- Beryllium compounds were once used in fluorescent lighting tubes, but this use was discontinued because of berylliosis in the workers manufacturing the tubes (see below).
- The James Webb Space Telescope [http://www.jwst.nasa.gov/Telescope/mirrortale/ (Beryllium related details from NASA here)] will have 18 hexagonal beryllium sections for its mirrors. Because JWST will face a temperature of −240 degrees Celsius (30 kelvins), the mirror is made of beryllium, a material capable of handling extreme cold better than glass. Beryllium contracts and deforms less than glass — and thus remains more uniform — in such temperatures.

History

The name beryllium comes from the Greek beryllos, beryl. At one time beryllium was referred to as glucinium (from Greek glykys, sweet), due to the sweet taste of its salts. This element was discovered by Louis Vauquelin in 1798 as the oxide in beryl and in emeralds. Friedrich Wöhler and A. A. Bussy independently isolated the metal in 1828 by reacting potassium on beryllium chloride.

Occurrence

Beryllium is an essential constituent of about 100 out of about 4000 known minerals, the most important of which are bertrandite (Be₄Si₂O₇(OH)₂), beryl (Al₂Be₃Si₆O₁₈), chrysoberyl (Al₂BeO₄), and phenakite (Be₂SiO₄). Precious forms of beryl are aquamarine and emerald. Along with hydrogen, helium, and lithium, some beryllium was created in the big bang. The most important commercial sources of beryllium and its compounds are beryl and bertrandite. Beryllium metal did not become readily available until 1957. Currently, most production of this metal is accomplished by reducing beryllium fluoride with magnesium metal. The price on the US market for vacuum-cast beryllium ingots was 338 US$ per pound ($745/kg) in 2001. [http://minerals.usgs.gov/minerals/pubs/commodity/beryllium/] ;Isolation BeF₂ + Mg → MgF₂ + Be

Isotopes

US$ Of beryllium's 10 isotopes, only beryllium-9 is stable. Cosmogenic beryllium-10 is produced in the atmosphere by cosmic ray spallation of oxygen and nitrogen. Because beryllium tends to exist in solution at pH levels less than about 5.5 (and most rainwater has a pH less than 5), it will enter into solution and be transported to the Earth's surface via rainwater. As the precipitation quickly becomes more alkaline, Be drops out of solution. Cosmogenic Be-10 thereby accumulates at the soil surface, where its relatively long half-life (1.51 million years) permits a long residence time before decaying to boron-10. Be-10 and its daughter products have been used to examine soil erosion, soil formation from regolith, the development of lateritic soils, as well as variations in solar activity and the age of ice cores. The fact that Be-7 and Be-8 are unstable has profound cosmological consequences as it means that elements heavier than beryllium could not be produced by nuclear fusion in the Big Bang. Moreover, the nuclear energy levels of beryllium-8 are such that carbon can be produced within stars, thus making life possible. (See triple-alpha process and Big Bang nucleosynthesis). The shortest-lived known isotope of beryllium is Be-13 which decays through neutron emission. It has a half-life of 2.7 × 10^-21 seconds. Be-6 also is also very short-lived with a half-life of 5.0 × 10^-21 seconds.

Precautions

neutron emission Beryllium and its salts are toxic substances and potentially carcinogenic. Chronic berylliosis is a pulmonary and systemic granulomatous disease caused by exposure to beryllium. Acute beryllium disease in the form of chemical pneumonitis was first reported in Europe in 1933 and in the United States in 1943. Cases of chronic berylliosis were first described in 1946 among workers in plants manufacturing fluorescent lamps in Massachusetts. Chronic berylliosis resembles sarcoidosis in many respects, and the differential diagnosis is often difficult. Although the use of beryllium compounds in fluorescent lighting tubes was discontinued in 1949, potential for exposure to beryllium exists in the nuclear and aerospace industries and in the refining of beryllium metal and melting of beryllium-containing alloys, the manufacturing of electronic devices, and the handling of other beryllium-containing material. Early researchers tasted beryllium and its various compounds for sweetness in order to verify its presence. Modern diagnostic equipment no longer necessitates this highly risky procedure and no attempt should be made to ingest this substance. Beryllium and its compounds should be handled with great care and special precautions must be taken when carrying out any activity which could result in the release of beryllium dust (lung cancer is a possible result of prolonged exposure to beryllium laden dust). This substance can be handled safely if certain procedures are followed. No attempt should be made to work with beryllium before familiarization with correct handling procedures.

Health effects

Beryllium can be harmful if inhaled. The effects depend on period of exposure. If beryllium air levels are high enough (greater than 1000 μg/m³), an acute condition can result, called acute beryllium disease, which resembles pneumonia. Occupational and community air standards are effective in preventing most acute lung damage. Some people (1-15%) become sensitive to beryllium. These individuals may develop an inflammatory reaction in the respiratory system. This condition is called chronic beryllium disease (CBD), and can occur many years after exposure to higher than normal levels of beryllium (greater than 0.2 μg/m³). This disease causes fatigue, weakness, and can cause difficulty in breathing. It can result in anorexia, weight loss, and may also lead to right-side heart enlargement and heart disease in advanced cases. Some people who are sensitized to beryllium may not have any symptoms. The general population is unlikely to develop acute or chronic beryllium disease because ambient air levels of beryllium are normally very low (0.00003-0.0002 μg/m³). Swallowing beryllium has not been reported to cause effects in humans because very little beryllium is absorbed from the stomach and intestines. Ulcers have been seen in dogs ingesting beryllium in the diet. Beryllium contact with skin that has been scraped or cut may cause rashes or ulcers. Long term exposure to beryllium can increase the risk of developing lung cancer. The United States Department of Health and Human Services (DHHS) and the International Agency for Research on Cancer (IARC) have determined that beryllium is a human carcinogen. The U.S. Environmental Protection Agency (EPA) has determined that beryllium is a probable human carcinogen. The EPA has estimated that lifetime exposure to 0.04 μg/m³ beryllium can result in a one in a thousand chance of developing cancer. There are no studies on the health effects of children exposed to beryllium. It is likely that the health effects seen in children exposed to beryllium will be similar to the effects seen in adults. It is unknown whether children differ from adults in their susceptibility to beryllium. It is unclear whether beryllium is teratogenic. Beryllium can be measured in the urine and blood. The amount of beryllium in blood or urine may not indicate time or quantity of exposure. Beryllium levels can also be measured in lung and skin samples. Another blood test, the blood beryllium lymphocyte proliferation test (BeLPT), identifies beryllium sensitization and has predictive value for CBD. Typical levels of beryllium that industries may release into the air are of the order of 0.01 μg/m³, averaged over a 30-day period, or 2 μg/m³ of workroom air for an 8-hour work shift.

References

- [http://periodic.lanl.gov/elements/4.html Los Alamos National Laboratory – Beryllium]

External links

- [http://www.webelements.com/webelements/elements/text/Be/index.html WebElements.com – Beryllium]
- [http://education.jlab.org/itselemental/ele004.html It's Elemental – Beryllium]
- [http://www-cie.iarc.fr/htdocs/monographs/vol58/mono58-1.htm IARC Monograph "Beryllium and Beryllium Compounds"]
- [http://www.compchemwiki.org/index.php?title=Beryllium Computational Chemistry Wiki]
- [http://theodoregray.com/PeriodicTable/Elements/004/ Applications and Images of Beryllium] Category:Chemical elements Category:Alkaline earth metals Category:Toxicology ko:베릴륨 ms:Berilium ja:ベリリウム simple:Beryllium th:เบริลเลียม

Chemical element

A chemical element, often called simply element, is a chemical substance that canot be divided or changed into other chemical substances by any ordinary chemical technique. The smallest unit of this kind of chemical substances is an atom. An element is a class of substances that contain the same number of protons in all its atoms.

Chemistry terminology

Earlier an element or pure element was defined as a substance which "cannot be further broken down into another compound with different chemical properties" -- which should be taken to mean it consists of atoms of one element. However, due to allotropy, the isotope effect, and the confusion with the more useful term referring to the general class of atoms (irrespective of what compound it may be in), this usage is in disfavor amongst contemporary chemists, and sees restricted, mostly historical, use. This definition was motivated by the observation that these elements could not be dissociated by chemical means into other compounds. For example, water could be converted into hydrogen and oxygen, but hydrogen and oxygen could not be further decomposed, thus "elemental". There are also many counterexamples (for example "elemental oxygen" (O₂) can be decomposed by solely chemical means into oxygen ions and atoms which have drastically different chemical properties). The remainder of this article will concern itself with the first definition.

Description

The atomic number of an element, Z, is equal to the number of protons which defines the element. For example, all carbon atoms contain 6 protons in their nucleus, so for carbon Z=6. These atoms may have different amounts of neutrons, and are known as isotopes of the element. The atomic mass of an element, A, is measured in unified atomic mass units (u) is the average mass of all the atoms of the element in an environment of interest (usually the earth's crust and atmosphere). Since electrons are light, and neutrons are barely more than the mass of the proton, this usually corresponds to the sum of the protons and neutrons in the nucleus of the most abundant isotope, though this is not always the case (notably chlorine, which is about three-quarters ³⁵Cl and a quarter ³⁷Cl). Some isotopes are radioactive and decay into other elements upon radiating an alpha or beta particle. Some elements have no nonradioactive isotopes, in particular all elements with Z >= 84. The lightest elements are hydrogen and helium. Hydrogen is thought to be the first element to appear after the Big Bang. All the heavier elements, are made naturally and artificially through various methods of nucleosynthesis. As of 2005, there are 116 known elements: 93 occur naturally on earth (including technetium and plutonium), and 94 (including promethium) have been detected so far in the universe. The 23 elements not found on earth are derived artificially; the first purportedly synthesized element was technetium, in 1937, although the trace amounts of naturally occurring technetium were not known then. All artificially derived elements are radioactive with short half-lives so that any such atoms that were present at the formation of Earth are extremely likely to have already decayed. Lists of the elements by name, by symbol, by atomic number, by density, by melting point and by boiling point are available. The most convenient presentation of the elements is in the periodic table, which groups elements with similar chemical properties together.

Nomenclature

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When it was learned, existing names (e.g., gold, mercury, iron) were kept in most countries, and national differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For example, the Germans use "Wasserstoff" for "hydrogen" and "Sauerstoff" for "oxygen," while some romance languages use "natrium" for "sodium" and "kalium" for "potassium," and the French prefer the obsolete but historic term "azote" for "nitrogen." But for international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that "aluminium" and "caesium" take the place of the US spellings "aluminum" and "cesium," while the US "sulfur" takes the place of the British "sulphur." But chemicals which are practicable to be sold in bulk within many countries, however, still have national names, and those which do not use the Latin alphabet cannot be expected to use the IUPAC name. According to IUPAC, the full name of an element is not capitalized, even if it is derived from a proper noun (unless it would be capitalized by some other rule, for instance if it begins a sentence). And in the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have too quick a decay rate to ever be sold in bulk. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question which delayed the naming of elements with atomic number of 104 and higher for a considerable time. (See element naming controversy). Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century (e.g., as "lutetium" refers to Paris, France, the Germans were reticent about relinquishing naming rights to the French, often calling it "cassiopeium"). And notably, the British discoverer of "niobium" originally named it "columbium," after the New World, though this did not catch on in Europe. The Americans had to accept the international name just when it was becoming an economically important material late in the twentieth century.

Chemical symbols

Specific chemical elements

Before chemistry became a science, alchemists had designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of one atoms combining to form molecules. With his advances in the atomic theory of matter, John Dalton devised his own simpler symbols, based on circles, which were to be used to depict molecules. These were superseded by the current typographical system in which chemical symbols are not used as mere abbreviations though each consists letters of the Latin alphabet - they are symbols intended to be used by peoples of all languages and alphabets. The first of these symbols were intended to be fully international, for they were based on the Latin abbreviations of the names of metals: Fe comes from Ferrum; Ag from Argentum. The symbols were not followed by a period (full stop) as abbreviations were. Besides a name, later chemical elements are also given a unique chemical symbol, based on the name of the element, not necessarily derived from the colloquial English name. (e.g., sodium has chemical symbol 'Na' after the Latin natrium). The same applies to "W" (wolframium) for Tungsten , "Hg" (Hydrargyrum) for mercury and "K" for potassium. Stricly taken, a symbol like Tu for tungsten or M or Me for mercury seems to be more logical. Chemical symbols are understood internationally when element names might need to be translated. There are sometimes differences; for example, the Germans have used "J" instead of "I" for iodine, so the character would not be confused with a roman numeral. The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always minuscule (small letters).

General chemical symbols

There are also symbols for series of chemical elements, for comparative formulas. These are one capital letter in length, and the letters are reserved so they are not permitted to be given for the names of specific elements. For example, an "X" is used to indicate a variable group amongst a class of compounds (though usually a halogen), while "R" is used for a radical (not to be confused with radical_(chemistry), meaning a compound structure such as a hydrocarbon chain. The letter "Q" is reserved for "heat" in a chemical reaction. "Y" is also often used as a general chemical symbol, although it is also the symbol of Yttrium. "Z" is also frequently used as a general variable group. "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

Nonelement symbols

Nonelements, especially in organic and organometallic chemistry, often acquire symbols which are inspired by the elemental symbols. A few examples: Cy - cyclohexyl; Ph - phenyl; Bz - benzoyl; Bn - benzyl; Cp - Cyclopentadiene; Pr - propyl; Me - methyl; Et - ethyl; Tf - triflate; Ts - tosyl.

External links

- [http://www.vanderkrogt.net/elements/ Elementymology & Elements Multidict] word history and language dictionary

Chemical information

- [http://www.webelements.com/ WebElements]
- [http://www.vcs.ethz.ch/chemglobe/ptoe/ ChemGlobe]
- [http://pearl1.lanl.gov/periodic/default.htm Los Alamos National Laboratory]
- [http://www.chemicalelements.com/ ChemicalElements] ko:화학 원소 ms:Unsur kimia ja:元素 simple:Element th:ธาตุเคมี

Atomic number

The atomic number (Z) is a term used in chemistry and physics to represent the number of protons found in the nucleus of an atom. In an atom of neutral charge, the number of electrons also equals the atomic number. The atomic number originally meant the number of an element's place in the periodic table. When Mendeleev arranged the known chemical elements grouped by their similarities in chemistry, it was noticeable that placing them in strict order of atomic mass resulted in some mismatches. Iodine and tellurium, if listed by atomic mass, appeared to be in the wrong order, and would fit better if their places in the table were swapped. Placing them in the order which fit chemical properties most closely, their number in the table was their atomic number. This number appeared to be approximately proportional to the mass of the atom, but, as the discrepancy showed, reflected some other property than mass. The anomalies in this sequence were finally explained after research by Henry Gwyn Jeffreys Moseley in 1913. Moseley discovered a strict relationship between the x-ray diffraction spectra of elements, and their correct location in the periodic table. It was later shown that the atomic number corresponds to the electric charge of the nucleus — in other words the number of protons. It is the charge which gives elements their chemical properties, rather than the atomic mass. The atomic number is closely related to the mass number (although they should not be confused) which is the number of protons and neutrons in the nucleus of an atom. The mass number often comes after the name of the element, e.g. carbon-14 (used in carbon dating).

Bivalent

In logic, the principle of bivalence states that for any proposition P, either P is true or P is false. This is not to be confused with the law of excluded middle and the law of noncontradiction. See bivalence and related laws for a summary of the differences. In classical logic, the principle of bivalence is equivalent to the result that there are no propositions that are neither true nor false. A proposition P that is neither true nor false is undecidable. In intuitionistic logic, sometimes the truth-value of a proposition P cannot be determined (i.e. P cannot be proved nor disproved). In such a case, P simply does not have a truth-value. Other logics, e.g. multi-valued logic, may assign P an indeterminate truth-value.

Metal

:For alternative meanings see metal (disambiguation). metal (disambiguation) In chemistry, a metal (Greek: Metallon) is an element that readily forms ions (cations) and has metallic bonds, and metals are sometimes described as a lattice of positive ions (cations) in a cloud of electrons. The metals are one of the three groups of elements as distinguished by their ionisation and bonding properties, along with the metalloids and nonmetals. On the periodic table, a diagonal line drawn from boron (B) to polonium (Po) separates the metals from the nonmetals. Elements on this line are metalloids, sometimes called semi-metals; elements to the lower left are metals; elements to the upper right are nonmetals. Nonmetal elements are more abundant in nature than are metallic elements, but metals in fact constitute most of the periodic table. Some well-known metals are aluminium, copper, gold, iron, lead, silver, titanium, uranium, and zinc. The allotropes of metals tend to be lustrous, ductile, malleable, and good conductors, while nonmetals generally speaking are brittle (for solid nonmetals), lack luster, and are insulators. A more modern definition of metals is that they have overlapping conductance and valence bands in their electronic structure. This definition opens up the category for metallic polymers and other organic metals, which have been made by researchers and employed in high-tech devices. These synthetic materials often have the characteristic silvery-grey reflectiveness of elemental metals. The properties of conductivity are mainly because each atom exerts only a loose hold on its outermost electrons (valence electrons); thus, the valence electrons form a sort of sea around the close-packed metal nucleii cations. Most metals are chemically unstable, reacting with oxygen in the air to form oxides over varying timescales (iron rusts over years, potassium burns in seconds, silver tarnishes in months, although this is due to reactions with sulfur, although ozone, which is three atoms of oxygen bound together, can also play a part, as can hydrogen sulfide). The alkali metals react quickest followed by the alkaline earth metals, found in the leftmost two groups of the periodic table. The transition metals take much longer to oxidise (e.g. iron, copper, zinc, nickel), and palladium, platinum and gold do not react with atmospheric oxygen at all (which is why we make shiny jewelry from them). Some metals form a barrier layer of oxide on their surface which cannot be penetrated by further oxygen molecules and thus retain their shiny appearance and good conductivity for many decades (e.g. aluminium, some steels, titanium and more). Painting or anodising metals are good ways to prevent their oxidation.

Alloys

An alloy is a mixture with metallic properties that contains at least one metal element. Examples of alloys are steel (iron and carbon), brass (copper and zinc), bronze (copper and tin), and duralumin (aluminium and copper). Alloys specially designed for highly demanding applications, such as jet engines, may contain more than ten elements.

Physical properties

Traditionally, metals have certain characteristic physical properties: they are usually shiny (they have "lustre"), have a high density, are ductile and malleable, usually have a high melting point, are usually hard, and conduct electricity and heat well. However, this is mainly because the low density, soft, low melting point metals happen to be reactive and we rarely encounter them in their elemental, metallic form. Metals are also sonorous, which means that they conduct sound well.

Metal oxides

The oxides of metals are basic; those of nonmetals are acidic.

Astronomy usage

In the specialised usage of astronomy and astrophysics, the term "metal" is often used to refer to any element other than hydrogen or helium. See metal-rich.

Alloy

An alloy is a combination, either in solution or compound, of two or more elements, which has a combination of at least one metal, and where the resultant material has metallic properties. An alloy with two components is called a binary alloy; one with three is a ternary alloy; one with four is a quaternary alloy. The result is a metallic substance with properties different from those of its components. Alloys are usually designed to have properties that are more desirable than those of their components. For instance, steel is stronger than iron, one of its main elements, and brass is more durable than copper, but more attractive than zinc. Unlike pure metals, many alloys do not have a single melting point. Instead, they have a melting range in which the material is a mixture of solid and liquid phases. The temperature at which melting begins is called the solidus, and that at which melting is complete is called the liquidus. Special alloys can be designed with a single melting point, however, and these are called eutectic mixtures. Sometimes an alloy is just named for the base metal, as 14 karat (58%) gold is an alloy of gold with other elements. The same holds for silver used in jewellery, and aluminium used structurally. Alloys include:
- aluminium bronze
- alnico
- amalgam
- brass
- bronze
- duralumin
- electrum
- galinstan
- intermetallics
- Mu-metal
- Nichrome
- pewter
- phosphor bronze
- solder
- spiegeleisen
- stainless steel
- steel
- Sterling silver
- Wood's metal
- ko:합금 ms:Aloi ja:合金 simple:Alloy

Beryllium copper

Beryllium copper is a metal alloy that contains 98% copper 2% beryllium, has significant metalworking advantages and great operating performance qualities. This is a highly ductile alloy which can be precisely stamped and shaped into a wide variety of complex patterns including very small yet critical parts. The addition of beryllium allows this alloy to be heat-treated into being a very strong and durable metal and also gives the alloy a high electrical conductivity. It is used in springs and to make other parts that need to retain their shapes over long periods of time while being subjected to repeated use. Category:Copper alloys

Metal

Alloys

Physical properties

Metal oxides

The oxides of metals are basic; those of nonmetals are acidic.

Astronomy usage

In the specialised usage of astronomy and astrophysics, the term "metal" is often used to refer to any element other than hydrogen or helium. See metal-rich.

Thermal conductivity

In physics, thermal conductivity, λ, is the intensive property of a material which relates its ability to conduct heat. Thermal conductivity is the quantity of heat, Q, transmitted through a thickness L, in a direction normal to a surface of area A, due to a temperature gradient ΔT, under steady state conditions and when the heat transfer is dependent only on the temperature gradient. :: thermal conductivity = heat flow rate × distance / (area × temperature gradient) :: λ = Q × L / (A × ΔT)

Examples

In general thermal conductivity tracks electrical conductivity; metals being good thermal conductors. There are exceptions: the most outstanding is that of diamond which has a high thermal conductivity, between 895 and 2300 W·m^-1·K^-1, while its electrical conductivity is low. The following table was compiled from data available in the [http://www.hbcpnetbase.com/ CRC Handbook of Chemistry and Physics]. All values are for materials at room temperature (298K), except where indicated. Thermal conductivity of other common materials:

Material	Thermal conductivity (298K) W·m^-1·K^-1
Diamond	895-2300
Carbon Nanotubes	1400
Silver	429
Copper	386
Gold	317
Aluminium	237
Brass	120
Platinum	71.6
Iron	80.2
Lead	35.3
Mercury	8.514
Quartz (273K)	6.8-12
Ice (273K)	2.2
Glass	1.35
Wood	0.04 (balsa) - 0.35 (fir)
Styrofoam	0.033
Wool	0.04
Silica aerogel	0.017
Air (100 kPa)	0.0262
Water	0.6062

Thermal conductivity changes with temperature. For most materials it decreases slightly as the temperature rises. A thermal conductance tester, one of the instruments of gemology, determines if gems are genuine diamonds using using diamond's uniquely high thermal conductivity, which is higher still for natural blue diamond. Diamonds of any size are notably cool to the touch because of their high thermal conductivity, perhaps the origin of the term "ice."

Related terms

The reciprocal of thermal conductivity is thermal resistivity, measured in kelvin-metres per watt (K·m·W^-1). When dealing with a known amount of material, its thermal conductance and the reciprocal property, thermal resistance, can be described. Unfortunately there are differing definitions for these terms.

First definition (general)

For general scientific use, thermal conductance is the quantity of heat that passes in unit time through a plate of particular area and thickness when its opposite faces differ in temperature by one degree. For a plate of thermal conductivity λ, area A and thickness L this is λA/L, measured in W·K^-1. This matches the relationship between electrical conductivity (A·m^-1·V^-1) and electrical conductance (A·V^-1). There is also a measure known as heat transfer coefficient: the quantity of heat that passes in unit time through unit area of a plate of particular thickness when its opposite faces differ in temperature by one degree. The reciprocal is thermal insulance. In summary:
- thermal conductance = λA/L, measured in W·K^-1
- thermal resistance = L/λA, measured in K·W^-1
- heat transfer coefficient = λ/L, measured in W·K^-1·m^-2
- thermal insulance = L/λ, measured in K·m²·W^-1. The heat transfer coefficient is also known as thermal admittance. But thermal admittance may mean other things.

Second definition (buildings)

When dealing with buildings, thermal resistance or R-value means what is described above as thermal insulance, and thermal conductance means the reciprocal. For materials in series, these thermal resistances (unlike conductances) can simply be added to give a thermal resistance for the whole. A third term, thermal transmittance, incoporates the thermal conductance of a structure along with heat transfer due to convection and radiation. It is measured in the same units as thermal conductance and is sometimes known as the composite thermal conductance. The term U value is another synonym. The term K value is a synonym for thermal conductivity. In summary, for a plate of thermal conductivity λ, area A and thickness L:
- thermal conductance = λ/L, measured in W·K^-1·m^-2
- thermal resistance (R value, thermal resistivity in scientific terms) = L/λ, measured in K·m²·W^-1.
- thermal transmittance (U value)= 1/(Σ(L/λ)) + convection + radiation, measured in W·K^-1·m^-2

Molecular origins

The thermal conductivity of a system is determined by how molecules comprising the system interact. There are no simple but correct expressions for the thermal conductivity. The simplest exact expression employs one of the Green-Kubo relations. Although this expression is exact, in order to calculate the thermal conductivity of a dense fluid or solid using this relation requires the use of molecular dynamics computer [http://rsc.anu.edu.au/~evans/evansmorrissbook.htm simulation].

External links

- http://physics.nist.gov/Pubs/SP811/appenB9.html
- http://gscassociates.com/wg8/edcs/text/unit.html
- http://www.iso.ch/iso/en/ittf/PubliclyAvailableStandards/ISO_IEC_18025_Ed1.html
- http://www.for.gov.bc.ca/hfp/pubs/silviculture_notes/sn16.pdf page 5
- http://www.npl.co.uk/thermal/faq_index.html#heat%20transfer%20property thermophysics FAQ5
- http://www.ornl.gov/roofs+walls/research/detailed_papers/rastra/dynamic.htm
- http://www.tak2000.com/data2.htm

References

Halliday, David; Resnick, Robert; & Walker, Jearl(1997). Fundamentals of Physics (5th ed.). John Wiley and Sons, INC., NY ISBN 0-471-10558-9. Category:Chemical properties Category:Thermodynamics Category:Physical quantity ja:熱伝導率

Nitric acid

The chemical compound nitric acid (H N O₃), otherwise known as aqua fortis, is a colorless, corrosive liquid, a toxic acid which can cause severe burns. At room temperature it gives off red or yellow fumes in moist air. Commonly used as a laboratory reagent, it is used in the manufacture of explosives such as nitroglycerin and trinitrotoluene (TNT), and as well as of fertilizers such as ammonium nitrate. It has additional uses in metallurgy and refining as it reacts with most metals, and in organic syntheses. When combined with hydrochloric acid it forms aqua regia, one of the few reagents capable of dissolving gold and platinum. Nitric acid is also a component of acid rain. Nitric acid is a strong acid with a pK_a of -2: in aqueous solution, it completely dissociates into the nitrate ion NO₃⁻ and a hydrated proton, known as a hydronium ion, H₃O⁺. The salts of nitric acid (which contain the nitrate ion) are also known as nitrates. The overwhelming majority of them are very soluble in water. Nitric acid is made by mixing nitrogen dioxide (N O₂) with water. Creating a very pure nitric acid usually involves distillation with sulfuric acid, as nitric acid forms an azeotrope with water with a composition of 68% nitric acid and 32% water. Commercial grade nitric acid solutions are usually between 52% and 68% nitric acid. If the nitric acid solution contains more than 86% nitric acid, it is referred to as fuming nitric acid, and can be separated into two kinds of fuming acids, white fuming nitric acid, and red fuming nitric acid. White fuming nitric acid, also called 100% nitric acid or WFNA, is very close to the anhydrous nitric acid product. One specification for white fuming nitric acid is that it has a maximum of 2 percent water and a maximum of 0.5 percent dissolved NO₂. Red fuming nitric acid, or RFNA, contains substantial quantities of dissolved nitrogen dioxide (NO₂) leaving the solution with a reddish-brown color. One formulation of RFNA specifies a minimum of 17% NO₂, another specifies 13% NO₂. In either event, an inhibited fuming nitric acid (either IWFNA, or IRFNA) can be made by the addition of 0.6 to 0.7% hydrogen fluoride, HF. This fluoride is added for corrosion resistance in metal tanks (the fluoride creates a metal fluoride layer that protects the metal). The obvious use for such a corrosion inhibited product is as an oxidizer in liquid fuel rockets. Nitric acid is a very powerful oxidizing agent and the reactions of nitric acid with compounds such as cyanides, carbides, and metallic powders can be explosive. Reactions of nitric acid with many organic compounds, such as turpentine, are violent and hypergolic (i.e. self-igniting). Concentrated nitric acid dyes human skin yellow on contact, due to interactions with the skin protein keratin. Strangely, these yellow stains turn orange when alkalised. Commercial production of nitric acid is via the Ostwald process after Wilhelm Ostwald. Nitric acid and its salts, the nitrates, should not be confused with nitrous acid and its salts, the nitrites.

Neutron

Classification

Neutron

Properties

Mass:	1.674 927 16(13) × 10⁻²⁷ kg
	939.565 530(38) MeV/c²
Electric charge:	0 C
Spin:	½
Magnetic dipole moment:	-1.91304 μ_N
Quark composition:	2 Down, 1 Up

In physics, the neutron is a subatomic particle with no net electric charge and a mass of 939.573 MeV/c² ( kg, slightly more than a proton). Its spin is ½. Its antiparticle is called the antineutron. The neutron and proton are instances of a nucleon. The nucleus of most atoms (all except the most common isotope of hydrogen, which consists of a single proton only) consists of protons and neutrons.

Properties

Outside the nucleus, neutrons are unstable and have a mean lifetime of 886 seconds (about 15 minutes, uncertainty about 2 s [http://www.ill.fr/pages/menu_g/docs/universe2003.pdf]), decaying by emitting an electron and antineutrino to become a proton. Neutrons in this unstable form are known as free neutrons. The same decay method (beta decay) occurs in some nuclei. Particles inside the nucleus are typically resonances between neutrons and protons, which transform into one another by the emission and absorption of pions. A neutron is classified as a baryon, and consists of two down quarks and one up quark. The neutron's antimatter equivalent is the antineutron. The number of neutrons determines the isotope of an element. (For example, the carbon-12 isotope has 6 protons and 6 neutrons, while the carbon-14 isotope has 6 protons and 8 neutrons.) Isotopes are atoms of the same element that have the same atomic number but different masses due to a different number of neutrons.

Neutron Interactions

The neutron interacts through all four of the common classifications of physical interaction. These four are the electromagnetic, weak nuclear, strong nuclear and gravitational interactions. Although it is true that the neutron has zero net charge, it is nonetheless composed of electrically charged quarks, in the same way that a neutral atom is nonetheless composed of protons and electrons. As such, the neutron experiences the electromagnetic interaction. The net charge is zero, so if you are far enough away from the neutron that it appears to occupy no volume, then the total effect of the electric force will add up to zero. The movement of the charges inside the neutrons do not cancel however, and this is what gives the neutron its nonzero magnetic moment. Gravity is often not discussed when talking about neutrons. This is because neutrons are usually studied in terms of subatomic interactions. In the subatomic world, gravity is undetectable relative to the other forces which are much stronger. This having been said, a neutron accelerates at the same rate in the earth's gravitational field as a lead brick. Charged particles (such as protons, electrons, or alpha particles) and electromagnetic radiation (such as gamma rays) lose energy in passing through matter. They exert electric forces which ionize atoms of the material through which they pass. The energy taken up in ionization equals the energy lost by the charged particle, which slows down, or by the gamma ray, which is absorbed or scattered (see compton scattering). The neutron, in contrast, is seen by atoms it passes as containing no electric charge, and so does not create any ionization. As far as the nuclear forces are concerned, it is a different story. Nuclear forces play the leading role when neutrons pass through regular matter. Consequently, a free neutron goes on its way unchecked until it makes a "head-on" collision with an atomic nucleus. When this happens, the neutrons and target nuclei can be scattered (deflected or slowed down), absorbed, or transformed into something different. In the case of the reaction n + ³He → ¹H + ³H (n:neutron; ³He: nucleus consisting of two protons and one neutron; ¹H: nucleus consisting of a only proton; ³H: nucleus consisting of one proton and two neutrons) for example, the proton and the neutron appear to have exchanged places, and kinetic energy is released. In many cases, secondary particles are created and energy can be used up or released. Neutrons, like other particles, can undergo elastic collisions. A collision is elastic under the special case where kinetic energy is conserved. Billiard balls for example typically undergo elastic collisions. The law of conservation of momentum also applies as it does for any collision. If the nucleus that is struck in an elastic collision is heavy, it acquires relatively little speed, but if it is a proton, which is approximately equal in mass to the neutron, it is projected forward with a large fraction of the original speed of the neutron, which is itself correspondingly slowed.

Neutron Detection

The common means of detecting a charged particle by looking for a track of ionization does not work for neutrons directly. Neutrons that elastically scatter off of another atom can create an ionization track that is detectable, but the experiments are not as simple to carry out and other means for detecting neutrons, consisting of allowing them to interact with atomic nuclei, are more commonly used. A common method for detecting neutrons involves converting the energy released from such reactions into electrical signals. The nuclides ³He, ⁶Li, ¹⁰B, ²³³U, ²³⁵U, ²³⁷Np and ²³⁹Pu are useful for this purpose. A good discussion on neutron detection is found in chapter 14 of the book Radiation Detection and Measurement by Glenn F. Knoll (John Wiley & Sons, 1979).

Neutron Uses

The neutron plays an important role in many nuclear reactions. For example, neutron capture often results in neutron activation, inducing radioactivity. In particular, knowledge of neutrons and their behavior has been important in the development of nuclear reactors and nuclear weapons. The development of [http://www.nature.com/nature/journal/v357/n6377/abs/357390a0.html "neutron lenses"] based on total internal reflection within hollow glass capillary tubes or by reflection from dimpled aluminum plates has driven ongoing research into [http://www.physorg.com/news599.html neutron microscopy] and [http://www.nasa.gov/vision/earth/technologies/nuggets.html neutron/gamma ray tomography]. One use of neutron emitters is the detection of light nuclei, particularly the hydrogen found in water molecules. When a fast neutron collides with a light nucleus, it loses a large fraction of its energy. By measuring the rate at which slow neutrons return to the probe after reflecting off of hydrogen nuclei, a neutron probe may determine the water content in soil.

Neutron Sources

Due to the fact that free neutrons are unstable, they (neutron radiation) can be obtained only from nuclear disintegrations, nuclear reactions, and high-energy reactions (such as in cosmic radiation showers or accelerator collisions). Free neutron beams are obtained from neutron sources by neutron transport. For access to intense neutron sources, researchers must go to specialist facilties, such as the ISIS facility in the UK, which is currently the world's most intense pulsed neutron and muon source. Neutrons' lack of total electric charge prevents engineers or experimentalists from being able to steer or accelerate them. Charged particles can be accelerated, decelerated, or deflected by electric or magnetic fields. However, these methods have almost no effect on neutrons (there is a small effect of a magnetic field on the free neutron because of its magnetic moment).

Discovery

In 1930 Walther Bothe and H. Becker in Germany found that if the very energetic alpha particles emitted from polonium fell on certain of the light elements, specifically beryllium, boron, or lithium, an unusually penetrating radiation was produced. At first this radiation was thought to be gamma radiation although it was more penetrating than any gamma rays known, and the details of experimental results were very difficult to interpret on this basis. The next important contribution was reported in 1932 by Irène Joliot-Curie and Frédéric Joliot in Paris. They showed that if this unknown radiation fell on paraffin or any other hydrogen-containing compound it ejected protons of very high energy. This was not in itself inconsistent with the assumed gamma ray nature of the new radiation, but detailed quantitative analysis of the data became increasingly difficult to reconcile with such a hypothesis. Finally (later in 1932) the physicist James Chadwick in England performed a series of experiments showing that the gamma ray hypothesis was untenable. He suggested that in fact the new radiation consisted of uncharged particles of approximately the mass of the proton, and he performed a series of experiments verifying his suggestion. Such uncharged particles were eventually called neutrons, apparently from the Latin root for neutral and the Greek ending -on (by imitation of electron and proton).

Current developments

The existence of stable clusters of four neutrons, or tetraneutrons, has been hypothesised by a team led by Francisco-Miguel Marqués at the CNRS Laboratory for Nuclear Physics based on observations of the disintegration of beryllium-14 nuclei. This is particularly interesting, because current theory suggests that these clusters should not be stable, and therefore should not exist.

Antineutron

The antineutron is the antiparticle of the neutron. It was discovered by Bruce Cork in the year 1956, a year after the antiproton was discovered. CPT-symmetry puts strong constraints on the relative properties of particles and antiparticles and, therefore, is open to stringent tests. The masses of the neutron and antineutron are equal to one part in (9±5)×10^-5.

Alpha particle

Alpha particles or alpha rays (named after the first letter in the greek alphabet, α) are a highly ionizing form of particle radiation which has low penetration. They consist of two protons and two neutrons bound together into a particle identical to a helium nucleus; hence, it can be written as He²⁺. Alpha particles are emitted by radioactive nuclei such as uranium or radium in a process known as alpha decay. This sometimes leaves the nucleus in an excited state, with the emission of a gamma ray removing the excess energy. In contrast to beta decay, alpha decay is mediated by the strong nuclear force. When an alpha particle is emitted, the atomic mass of an element goes down by roughly 4 amu, due to the loss of 4 nucleons. The atomic number of the atom goes down by 2, as the atom loses 2 protons, essentially becoming a new element. An example of this is when radium becomes radon gas due to alpha decay. Because of their charge and large mass, alpha rays are easily absorbed by materials and can travel only a few centimeters in air. They can be absorbed by tissue paper or the outer layers of human skin (about 40 micrometres, equivalent to a few cells deep) and so are not generally dangerous to life unless the source is ingested or inhaled. Because of this high mass and strong absorption, however, if alpha radiation does enter the body (most often because radioactive material has been inhaled or ingested), it is the most destructive form of ionizing radiation. It is the most strongly ionizing, and with large enough doses can cause any or all of the symptoms of radiation poisoning. It is estimated that chromosome damage from alpha particles is about 100 times greater than that caused by an equivalent amount of other radiation. The alpha emitter polonium-210 is suspected of playing a role in lung and bladder cancer related to tobacco smoking. Most smoke detectors contain a small amount of the alpha emitter americium-241. This isotope is extremely dangerous if inhaled or ingested, but the danger is minimal if the source is kept sealed. Many municipalities have established programs to collect and dispose of old smoke detectors, rather than let them go into the general waste stream. Because alpha particles occur naturally, but can have energy high enough to participate in a nuclear reaction, study of them led to much early knowledge of nuclear physics. The physicist Ernest Rutherford famously used alpha particles to infer that Lord Kelvin's "plum pudding" model of the atom was fundamentally flawed. He did this by coating a screen which flashed wherever it was struck by an alpha particle then surrounding a thin piece of gold foil with this screen. He then aimed alpha particles at the foil, hypothesizing that, assuming the "plum pudding" model of the atom was correct, the positively charged alpha particles would be only slightly deflected, if at all, by the dispersed positive charge predicted. It was found that some of the alpha particles were deflected at much larger angles than expected, with some even bouncing back. Although most of the alpha particles went straight through as expected, Rutherford commented that the few particles that were deflected was akin to shooting a cannonball at tissue paper only to have it bounce off, again assuming the "plum pudding" theory were correct. It was soon determined that the positive charge of the atom was concentrated in a small area in the center of the atom, hence making the postive charge dense enough to deflect any positively charged alpha particles that happened to come close to what was later termed the nucleus. (It was not known at the time that alpha particles were themselves nuclei nor was the existence of protons or neutrons known.) Rutherford's experiment subsequently led to the Bohr model and later the modern wave-mechanical model of the atom. In computer technology, DRAM 'soft' errors were linked to alpha particles in 1978 in Intel's DRAM chips. The discovery has led to strict control of radioactive elements in the packaging of semiconductor materials, and the problem was largely considered 'solved'.

References

- Category:Radioactivity ja:アルファ粒子

Radium

Radium is a chemical element, which has the symbol Ra and atomic number 88 (see the periodic table). Its appearance is almost pure white, but it readily oxidizes on exposure to air, turning black. Radium is an alkaline earth metal that is found in trace amounts in uranium ores. It is extremely radioactive. Its most stable isotope, Ra-226, has a half-life of 1602 years and decays into radon gas.

Notable characteristics

The heaviest of the alkaline earth metals, radium is intensely radioactive and resembles Barium chemically. This metal is found (combined) in minute quantities in the uranium ore pitchblende, and various other uranium minerals. Radium preparations are remarkable for maintaining themselves at a higher temperature than their surroundings, and for their radiations, which are of three kinds: alpha rays, beta rays, and gamma rays. Radium also produces neutrons when mixed with beryllium. When freshly prepared, pure radium metal is brilliant white, but blackens when exposed to air (probably due to nitride formation). Radium is luminescent (giving a faint blue color), corrodes in water to form radium hydroxide and is a bit more volatile than barium.

Applications

Some of the practical uses of radium are derived from its radiative properties. More recently discovered radioisotopes, such as cobalt-60 and caesium-137, are replacing radium in even these limited uses because several of these are much more powerful and others are safer to handle.
- Formerly used in self-luminous paints for watches, clocks and instrument dials. More than 100 former watch dial painters who used their lips to shape the paintbrush died from the radiation. Soon afterward, the adverse effects of radioactivity were popularized. Radium was still used in dials as late as the 1950's. Objects painted with this paint may still be dangerous, and must be handled properly. Currently, tritium is used instead of radium. Although tritium still carries some risks, it is considered by many to be safer than radium.
- When mixed with Beryllium it is a Neutron source for physics experiments.
- Radium (usually in the form of radium chloride) is used in medicine to produce radon gas which in turn is used as a cancer treatment.
- One unit for radioactivity, the non-SI curie, is based on the radioactivity of radium-226 (see Radioactivity).

History

Radium (Latin radius, ray) was discovered by Marie Curie and her husband Pierre in 1898 in pitchblende/uraninite from North Bohemia. While studying pitchblende the Curies removed uranium from it and found that the remaining material was still radioactive. They then separated out a radioactive mixture mostly consisting of barium which gave a brilliant red flame color and spectral lines which had never been documented before. In 1902 radium was isolated into its pure metal by Curie and Andre Debierne through the electrolysis of a pure radium chloride solution by using a mercury cathode and distilling in an atmosphere of hydrogen gas. Historically the decay products of radium were known as Radium A, B, C, etc. These are now known to be isotopes of other elements as follows: :Radium emanation - radon-222 :Radium A - polonium-218 :Radium B - lead-214 :Radium C - bismuth-214 :Radium C₁ - polonium-214 :Radium C₂ - thallium-210 :Radium D - lead-210 :Radium E - bismuth-210 :Radium F - polonium 210 On February 4, 1936 radium E became the first radioactive element to be made synthetically. During the 1930s it was found that worker exposure to radium by handling luminescent paints caused serious health effects which included sores, anemia and bone cancer. This use of radium was stopped soon afterward. This is because radium is treated as calcium by the body, and deposited in the bones, where radioactivity degrades marrow, and can mutate bone cells. Handling of radium has since been blamed for Marie Curie's premature death.

Occurrence

Radium is a decay product of uranium and is therefore found in all uranium-bearing ores. Radium was originally acquired from pitchblende ore from Joachimsthal, Bohemia (7 metric tons of pitchblende yields 1 gram of radium). Carnotite sands in Colorado provide some of the element, but richer ores are found in the Democratic Republic of the Congo, the Great Lakes area of Canada and can also be extracted from uranium processing waste. Large uranium deposits are located in Ontario, New Mexico, Utah, Australia, and in other places.

Compounds

Its compounds color flames crimson carmine (rich red or crimson color with a shade of purple) and give a characteristic spectrum. Due to its geologically short half life and intense radioactivity, radium compounds are quite rare, occurring almost exclusively in uranium ores.
- radium fluoride (RaF₂)
- radium chloride (RaCl₂)
- radium bromide (RaBr₂)
- radium iodide (RaI₂)
- radium oxide (RaO)

Isotopes

Radium has 25 different isotopes, four of which are found in nature, with radium-226 being the most common. Ra-223, Ra-224, Ra-226 and Ra-228 are all generated in the decay of either U or Th. Ra-226 is a product of U-238 decay, and is the longest-lived isotope of radium with a half-life of 1602 years; next longest is Ra-228, a product of Th-232 breakdown, with a half-life of 6.7 years.

Radioactivity

Radium is over one million times more radioactive than the same mass of uranium. Its decay occurs in at least seven stages; the successive main products have been studied and were called radium emanation or exradio (this is radon), radium A (polonium), radium B (lead), radium C (bismuth), etc. (The radon is a heavy gas, the later products are solids.) These products are themselves radioactive elements, each with an atomic weight a little lower than its predecessor. Radium loses about 1% of its activity in 25 years, being transformed into elements of lower atomic weight with lead being a final product of disintegration. The SI unit of radioactivity is the becquerel (Bq), equal to one disintegration per second. The curie is a non-SI unit defined as that amount of radioactivity which has the same disintegration rate as 1 gram of Ra-226 (3.7 x 10¹⁰ disintegrations per second, or 37 GBq).

Precautions

Radium is highly radioactive and its decay product, radon gas is also radioactive. Since radium is chemically similar to calcium, it has the potential to cause great harm by replacing it in bone. Inhalation, injection, ingestion or body exposure to radium can cause cancer and other body disorders. Stored radium should be ventilated to prevent accumulation of radon. Emitted energy from the decay of radium ionizes gases, affects photographic plates, causes sores on the skin, and produces many other detrimental effects.

References

- Guide to the Elements - Revised Edition, Albert Stwertka, (Oxford University Press; 1998) ISBN 0-19-508083-1
- [http://periodic.lanl.gov/elements/88.html Los Alamos National Laboratory - Radium]
- [http://www.nytimes.com/library/national/science/100698sci-radium.html A Glow in the Dark, and a Lesson in Scientific Peril]

External links

- [http://www.webelements.com/webelements/elements/text/Ra/index.html WebElements.com - Radium] (also used as a reference)
- [http://www.lateralscience.co.uk/radium/RaDisc.html Lateral Science - Radium Discovery]
- [http://www.markwshead.com/stuffHappens/radium.html Photos of Radium Water Bath in Oklahoma] Category:Chemical elements Category:Alkaline earth metals ko:라듐 ja:ラジウム th:เรเดียม

Polonium

Polonium is a chemical element in the periodic table that has the symbol Po and atomic number 84. A rare radioactive metalloid, polonium is chemically similar to tellurium and bismuth and occurs in uranium ores. Polonium had been studied for possible use in heating spacecraft.

Notable characteristics

This radioactive substance dissolves readily in dilute acids, but is only slightly soluble in alkalis. It is closely related chemically to bismuth and tellurium. Polonium is a volatile metal with 50% being vaporized in air after 45 hours at 328 K. Polonium has no stable isotopes and has over 50 potential isotopes. Polonium is extremely toxic and highly radioactive. Polonium has been found in tobacco smoke as a contaminant and in uranium ores.

Applications

When it is mixed or alloyed with beryllium, polonium can be a neutron source. Other uses;
- This element has also been used in devices that eliminate static charges in textile mills and other places. However, beta sources are more commonly used and are less dangerous.
- Polonium is used on brushes that remove accumulated dust from photographic films. The polonium in these brushes is sealed and controlled thus minimizing radiation hazards.

Polonium-210

This isotope of polonium is an alpha emitter that has a half-life of 138.39 days. A milligram of polonium-210 emits as many alpha particles as 5 grams of radium. A great deal of energy is released by its decay with a half a gram quickly reaching a temperature above 750 K. A few curies (gigabecquerels) of polonium-210 emit a blue glow which is caused by excitation of surrounding air. A single gram of polonium-210 generates 140 watts of heat energy. Since nearly all alpha radiation can be easily stopped by ordinary containers and upon hitting its surface releases its energy, Polonium-210 has been used as a lightweight heat source to power thermoelectric cells in artificial satellites. Because of its short halflife though polonium-210 cannot provide power for long-term space missions and has been phased out of use in this application.

History

Also called Radium F, polonium was discovered by Marie Curie and her husband Pierre Curie in 1898 and was later named after Marie's home land of Poland. Poland at the time was under Russian, Prussian and Austrian domination, and not recognized as an independent country. It was Marie's hope that naming the element after her home land would add notoriety to its plight. Polonium may be the first element named to highlight a political controversy. This element was the first one discovered by the Curies while they were investigating the cause of pitchblende radioactivity. The pitchblende, after removal of uranium and radium, was more radioactive than both radium and uranium put together. This spurred them on to find the element. The electroscope showed it separating with bismuth.

Occurrence

A very rare element in nature, polonium is found in uranium ores at about 100 micrograms per metric ton (1:10¹⁰). Its natural abundance is approximately 0.2% of radium's. In 1934 an experiment showed that when natural bismuth (Bi-209) is bombarded with neutrons, Bi-210, which is the parent of polonium, was created. Polonium may now be made in milligram amounts in this procedure which uses high neutron fluxes found in nuclear reactors.

Isotopes

Polonium has many isotopes all of which are radioactive. There are 25 known isotopes of polonium with atomic masses that range from 194 u to 218 u. Polonium-210 is the most widely available. Po-209 (half-life 103 years) and Po-208 (half-life 2.9 years) can be made through the alpha, proton, or deuteron bombardment of lead or bismuth in a cyclotron. However these isotopes are expensive to produce.

Precautions

Polonium is a highly radioactive and toxic element and is dangerous to handle. Even milligram or microgram amounts, handling polonium-210 is very dangerous and requires special equipment used with strict procedures. Direct damage occurs from energy absorption into tissues from alpha particles. The maximum allowable body burden for ingested polonium is only 1100 becquerels (0.03 microcurie), which is equivalent to a particle weighing only 6.8 x 10^-12 gram. Weight for weight polonium is approximately 2.5 x 10¹¹ times as toxic as hydrocyanic acid. The maximum permissible concentration for airborne soluble polonium compounds is about 7,500 Bq/m³ (2 x 10^-11 µCi/cm³).

References

- [http://periodic.lanl.gov/elements/84.html Los Alamos National Laboratory – Polonium]

External links

- [http://www.webelements.com/webelements/elements/text/Po/index.html WebElements.com – Polonium]
- [http://www.globalsecurity.org/wmd/intro/polonium.htm History of Polonium] Category:Metalloids Category:Chalcogens ko:폴로늄 ja:ポロニウム th:พอโลเนียม

Oxidation

Redox reactions include all chemical processes in which atoms have their oxidation number (oxidation state) changed. This can be a simple redox process, such as the combustion of carbon to yield carbon dioxide, it could be the reduction of carbon by hydrogen to yield methane, or it could be the oxidation of sugar in the human body, through a series of very complex electron transfer processes. The term redox comes from the two concepts of reduction and oxidation. :Oxidation describes the loss of an electron by a molecule, atom, or ion; loss of hydrogen, or gain of oxygen. It also means an increase in oxidation number. :Reduction describes the uptake of an electron by a molecule, atom, or ion; loss of oxygen and gain of hydrogen. It also means a decrease in oxidation number. These two terms go together, because in a chemical reaction, one cannot occur without the other; electrons lost by one compound must be gained by another. Reduction can also be considered to be the reducing of an atom's positive charge, and oxidation its opposite (gaining positive charge). oxygen

Oxidizing and reducing agents

Substances that have the ability to oxidize (Commonwealth English oxidise) other substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers. Put in another way, the oxidant removes electrons from the other substance, and is thus reduced itself. Oxidants are usually chemical substances with elements in high oxidation numbers (e.g. H₂O₂, MnO₄^-, CrO₃, Cr₂O₇^2-, OsO₄) or highly electronegative substances that can gain one or two extra electrons by oxidizing a substance (O₂, O₃, F₂, Cl₂, Br₂). Substances that have the ability to reduce other substances are said to be reductive and are known as reductive agents, reductants, or reducers. Put in another way, the reductant transfers electrons to the substance. Reductants in chemistry are very diverse. Metal reduction - electropositive elemental metals can be used (Li, Na, Mg, Fe, Zn, Al). These metals donate or give away electrons readily. Other kinds of reductants are hydride transfer reagents (NaBH₄, LiAlH₄), these reagents are widely used in organic chemistry, primarily in the reduction of carbonyl compounds to alcohols. Another useful method is reductions involving hydrogen gas (H₂) with a palladium, platinum, or nickel catalyst. These catalytic reductions are primarily used in the reduction of carbon-carbon double or triple bonds. The chemical way to look at redox processes is that the reductant transfers electrons to the oxidant. Thus, at the end of the reaction, the reductant will have been oxidized and the oxidant will have been reduced.

Former meaning (oxygen/hydrogen)

Formerly, oxidation simply meant the addition of oxygen or the removing of hydrogen (hence the name oxidation), and reduction was removal of oxygen or the addition of hydrogen. Currently, however, the terms are normally used in a more general sense, describing electron movement.

Examples of redox reactions

A good example is the reaction between hydrogen and fluorine: :H₂ + F₂ → 2HF We can write this overall reaction as two half-reactions: an oxidation reaction: :H₂ → 2H⁺ + 2e^- and a reduction reaction: :F₂ + 2e^- → 2F^- Elements always have an oxidation number of zero. In the first half reaction hydrogen is oxidized from an oxidation num

Beryllium

Notable characteristics

Applications

History

Occurrence

Isotopes

Precautions

Health effects

References

External links

Chemical element

Chemistry terminology

Description

Nomenclature

Chemical symbols

Specific chemical elements

General chemical symbols

Nonelement symbols

See also

External links

Chemical information

Atomic number

See also

Bivalent

See also

Metal

Alloys

Physical properties

Metal oxides

Astronomy usage

See also

Alloy

Beryllium copper

Metal

Alloys

Physical properties

Metal oxides

Astronomy usage

See also

Thermal conductivity

Examples

Related terms

First definition (general)

Second definition (buildings)

Molecular origins

See also

External links

References

Nitric acid

Neutron

Properties

Neutron Interactions

Neutron Detection

Neutron Uses

Neutron Sources

Discovery

Current developments

Antineutron

See also

Fields concerning neutrons

Types of neutrons

Objects containing neutrons

Neutron sources

Processes involving neutrons

Alpha particle

See also

References

Radium

Notable characteristics

Applications

History

Occurrence

Compounds

Isotopes

Radioactivity

Precautions

Further reading

References

External links

Polonium