:: wikimiki.org ::

Iron(II) Sulfate

Iron(II) sulfate

Iron(II) sulfate, also known as ferrous sulfate and as copperas (Fe S O₄) is an example of an ionic compound. It is found in various states of hydration (FeSO₄·H₂O, FeSO₄·4H₂O, FeSO₄·5H₂O, FeSO₄·7H₂O); the heptahydrate is also called green vitriol, copperas, or melanterite (a mineral that commonly occurs with pyrite). Iron(II) sulfate has a blue-green color, monoclinic crystal structure, and is water-soluble. Its molecular weight is 151.9026 g/mol. Its melting point is 64°C, and at 90°C it loses water of hydration to form the monohydrate, a white powder known as the mineral szomolnokite when it occurs naturally. Iron sulfate pentahydrate forms the mineral siderotil. Iron(II) sulfate is prepared commercially by oxidation of pyrite or by treating iron with sulfuric acid. It is used in the manufacture of inks, in wool dyeing as a mordant, and in water purification as a substitute for aluminium sulfate. It can also be used to treat iron deficiency. In the finishing of steel prior to plating or coating, the steel sheet or rod is passed through pickling lines of sulfuric acid. This produces large quantities of iron(II) sulfate as a waste product. ---- Ferrous Sulfate: An iron compound used to treat iron-deficiency anemia. ---- Category:Sulfates Category:Iron compounds

Iron

Iron is a chemical element with the symbol Fe (L.: Ferrum) and atomic number 26. Iron is a group 8 and period 4 metal. Iron is notable for being the final element produced by stellar nucleosynthesis, and thus the heaviest element which does not require a supernova or similarly cataclysmic event for its formation. It is therefore the most abundant heavy metal in the universe.

Notable characteristics

Iron is the most abundant metal on Earth, and is believed to be the tenth most abundant element in the universe. Iron is also the most abundant (by mass, 34.6%) element making up the Earth; the concentration of iron in the various layers of the Earth ranges from high at the inner core to about 5% in the outer crust; it is possible the Earth's inner core consists of a single iron crystal although it is more likely to be a mixture of iron and nickel; the large amount of iron in the Earth is thought to contribute to its magnetic field. Iron is a metal extracted from iron ore, and is hardly ever found in the free (elemental) state. In order to obtain elemental iron, the impurities must be removed by chemical reduction. Iron is used in the production of steel, which is not an element but an alloy, a solution of different metals (and some non-metals, particularly carbon). Nuclei of iron have some of the highest binding energies per nucleon, superseded only by the nickel isotope ⁶²Ni. The universally most abundant of the highly stable nucleides is, however, ⁵⁶Fe. This is formed by nuclear fusion in the stars. Although a further tiny energy gain could be extracted by synthesizing ⁶²Ni, conditions in stars are not right for this process to be favoured. When a very large star contracts at the end of its life, internal pressure and temperature rise, allowing the star to produce progressively heavier elements, despite these being less stable than the elements around mass number 60 (the "iron group"). This leads to a supernova. Some cosmological models with an open universe predict that there will be a phase where as a result of slow fusion and fission reactions, everything will become iron.

Applications

Iron is the most used of all the metals, comprising 95 percent of all the metal tonnage produced worldwide. Its combination of low cost and high strength make it indispensable, especially in applications like automobiles, the hulls of large ships, and structural components for buildings. Steel is the best known alloy of iron, and some of the forms that iron takes include:
- Pig iron has 4% – 5% carbon and contains varying amounts of contaminants such as sulfur, silicon and phosphorus. Its only significance is that of an intermediate step on the way from iron ore to cast iron and steel.
- Cast iron contains 2% – 4.0% carbon , 1% – 6% silicon , and small amounts of manganese. Contaminants present in pig iron that negatively affect the material properties, such as sulfur and phosphorus, have been reduced to an acceptable level. It has a melting point in the range of 1420–1470 K, which is lower than either of its two main components, and makes it the first product to be melted when carbon and iron are heated together. Its mechanical properties vary greatly, dependant upon the form carbon takes in the alloy. 'White' cast irons contain their carbon in the form of cementite, or iron carbide. This hard, brittle compound dominates the mechanical properties of white cast irons, rendering them hard, but unresistant to shock. The broken surface of a white cast iron is full of fine facets of the broken carbide, a very pale, silvery, shiny material, hence the appellation. In 'grey' cast iron, the carbon exists free as fine flakes of graphite , and also, renders the material brittle due to the stress-raising nature of the sharp edged flakes of graphite. A newer variant of grey iron, referred to as 'ductile iron' is specially treated with trace amounts of magnesium to alter the shape of graphite to sheroids, or nodules, vastly increasing the toughness and strength of the material.
- Carbon steel contains between 0.5% and 1.5% carbon, with small amounts of manganese, sulfur, phosphorus, and silicon.
- Wrought iron contains less than 0.2% carbon. It is a tough, malleable product, not as fusible as pig iron. It has a very small amount of carbon, a few tenths of a percent. If honed to an edge, it loses it quickly. Wrought iron is characterised, especially in old samples, by the presence of fine 'stringers' or filaments of slag entrapped in the metal.
- Alloy steels contain varying amounts of carbon as well as other metals, such as chromium, vanadium, molybdenum, nickel, tungsten, etc. They are used for structural purposes, as their alloy content raises their cost and necessitates justification of their use. Recent developments in ferrous metallurgy have produced a growing range of microalloyed steels, also termed 'HSLA' or high-strength, low alloy steels, containing tiny additions to produce high strengths and often spectacular toughness at minimal cost.
- Iron(III) oxides are used in the production of magnetic storage in computers. They are often mixed with other compounds, and retain their magnetic properties in solution.

History

The first signs of use of iron come from the Sumerians and the Egyptians, where around 4000 BC, a few items, such as the tips of spears, daggers and ornaments, were being fashioned from iron recovered from meteorites. Because meteorites fall from the sky some linguists have conjectured that the English word iron (OE īsern), which has cognates in many northern and western European languages, derives from the Etruscan aisar which means "the gods". By 3000 BC to 2000 BC, increasing numbers of smelted iron objects (distinguishable from meteoric iron by the lack of nickel in the product) appear in Mesopotamia, Anatolia, and Egypt. However, their use appears to be ceremonial, and iron was an expensive metal, more expensive than gold. In the Iliad, weaponry is mostly bronze, but iron ingots are used for trade. Some resources (see the reference What Caused the Iron Age? below) suggest that iron was being created then as a by-product of copper refining, as sponge iron, and was not reproducible by the metallurgy of the time. By 1600 BC to 1200 BC, iron was used increasingly in the Middle East, but did not supplant the dominant use of bronze. bronze In the period from the 12th to 10th century BC, there was a rapid transition in the Middle East from bronze to iron tools and weapons. The critical factor in this transition does not appear to be the sudden onset of a superior ironworking technology, but instead the disruption of the supply of tin. This period of transition, which occurred at different times in different parts of the world, is the ushering in of an age of civilization called the Iron Age. Concurrent with the transition from bronze to iron was the discovery of carburization, which was the process of adding carbon to the irons of the time. Iron was recovered as sponge iron, a mix of iron and slag with some carbon and/or carbide, which was then repeatedly hammered and folded over to free the mass of slag and oxidise out carbon content, so creating the product wrought iron. Wrought iron was very low in carbon content and was not easily hardened by quenching. The people of the Middle East found that a much harder product could be created by the long term heating of a wrought iron object in a bed of charcoal, which was then quenched in water or oil. The resulting product, which had a surface of steel, was harder and less brittle than the bronze it began to replace. In China the first irons used were also meteoric iron, with archeological evidence for items made of wrought iron appearing in the northwest, near Xinjiang, in the 8th century BC. These items were made of wrought iron, created by the same processes used in the Middle East and Europe, and were thought to be imported by non-Chinese people. In the later years of the Zhou Dynasty (ca 550 BC), a new iron manufacturing capability began because of a highly developed kiln technology. Producing blast furnaces capable of temperatures exceeding 1300 K, the Chinese developed the manufacture of cast, or pig iron. Iron was used in India as early as 250 BCE. The famous iron pillar in the Qutb complex in Delhi is made of very pure iron (98%) and has not rusted or eroded till this day. Delhi of wood annually from 1827 to 1891.]] If iron ores are heated with carbon to 1420–1470 K, a molten liquid is formed, an alloy of about 96.5% iron and 3.5% carbon. This product is strong, can be cast into intricate shapes, but is too brittle to be worked, unless the product is decarburized to remove most of the carbon. The vast majority of Chinese iron manufacture, from the Zhou dynasty onward, was of cast iron. Iron, however, remained a pedestrian product, used by farmers for hundreds of years, and did not really affect the nobility of China until the Qin dynasty (ca 221 BC). Cast iron development lagged in Europe, as the smelters could only achieve temperatures of about 1000 K. Through a good portion of the Middle Ages, in Western Europe, iron was still being made by the working of sponge iron into wrought iron. Some of the earliest casting of iron in Europe occurred in Sweden, in two sites, Lapphyttan and Vinarhyttan, between 1150 and 1350 AD. There are suggestions by scholars that the practice may have followed the Mongols across Russia to these sites, but there is no clear proof of this hypothesis. In any event, by the late fourteenth century, a market for cast iron goods began to form, as a demand developed for cast iron cannonballs. Early iron smelting (as the process is called) used charcoal as both the heat source and the reducing agent. In 18th century England, wood supplies ran down and coke, a fossil fuel, was used as an alternative. This innovation by Abraham Darby supplied the energy for the Industrial Revolution.

Occurrence

Industrial Revolution Iron is one of the more common elements on Earth, making up about 5% of the Earth's crust. Most of this iron is found in various iron oxides, such as the minerals hematite, magnetite, and taconite. The earth's core is believed to consist largely of a metallic iron-nickel alloy. About 5% of the meteorites similarly consist of iron-nickel alloy. Although rare, these are the major form of natural metallic iron on the earth's surface. Iron is also one of the least reactive metals, and therefore, it is sometimes found pure in nature.

Extraction from ore

Industrially, iron is extracted from its ores, principally hematite (nominally Fe₂O₃) and magnetite (Fe₃O₄) by a carbothermic reaction (reduction with carbon) in a blast furnace at temperatures of about 2000°C. In a blast furnace, iron ore, carbon in the form of coke, and a flux such as limestone are fed into the top of the furnace, while a blast of heated air is forced into the furnace at the bottom. In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide: :6 C + 3 O₂ → 6 CO The carbon monoxide reduces the iron ore (in the chemical equation below, hematite) to molten iron, becoming carbon dioxide in the process: :6 CO + 2 Fe₂O₃ → 4 Fe + 6 CO₂ The flux is present to melt impurities in the ore, principally silicon dioxide sand and other silicates. Common fluxes include limestone (principally calcium carbonate) and dolomite (magnesium carbonate). Other fluxes may be used depending on the impurities that need to be removed from the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (quicklime): :CaCO₃ → CaO + CO₂ Then calcium oxide combines then with silicon dioxide to form a slag. :CaO + SiO₂ → CaSiO₃ The slag melts in the heat of the furnace, which silicon dioxide would not have. In the bottom of the furnace, the molten slag floats on top of the more dense liquid iron, and spouts in the side of the furnace may be opened to drain off either the iron or the slag. The iron, once cooled, is called pig iron, while the slag can be used as a material in road construction or to improve mineral-poor soils for agriculture. Approximately 1100Mt (million tons) of iron ore was produced in the world in 2000, with a gross market value of approximately 25 billion US dollars. While ore production occurs in 48 countries, the five largest producers were China, Brazil, Australia, Russia and India, accounting for 70% of world iron ore production. The 1100Mt of iron ore was used to produce approximately 572Mt of pig iron.

Compounds

2000 production.]] Common oxidation states of iron include:
- the Iron(-II) state, Fe^2- (e.g. Fe(CO)₄^2-,Fe(CO)₂(NO)₂.
- the Iron(0) state, Fe(CO)₅, Fe(PF₃)₅.
- the Iron(I) state, [Fe(H₂O)₅NO]²⁺.
- the Iron(II) state, Fe²⁺, previously ferrous is very common.
- the Iron(III) state, Fe³⁺, previously ferric, is also very common, for example in rust.
- the Iron(IV) state, Fe⁴⁺, previously ferryl, stabilized in some enzymes (e.g. peroxidases).
- the Iron(VI) state, Fe⁶⁺ is also known, if rare, in potassium ferrate. Iron carbide Fe₃C is known as cementite.

Biological role

Iron is essential to all organisms, except for a few bacteria. It is mostly stably incorporated in the inside of metalloproteins, because in exposed or in free form it causes production of free radicals that are generally toxic to cells. To say that iron is free doesn't mean that it is free floating in the bodily fluids. Iron binds avidly to virtually all biomolecules so it will adhere nonspecifically to cell membranes, nucleic acids, proteins etc. Many animals incorporate iron into the heme complex, an essential component of cytochromes, which are proteins involved in redox reactions (including but not limited to cellular respiration), and of oxygen carrying proteins hemoglobin and myoglobin. Inorganic iron involved in redox reactions is also found in the iron-sulfur clusters of many enzymes, such as nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase. A class of non-heme iron proteins is responsible for a wide range of functions within several life forms, such as enzymes methane monooxygenase (oxidizes methane to methanol), ribonucleotide reductase (reduces ribose to deoxyribose; DNA biosynthesis), hemerythrins (oxygen transport and fixation in marine invertebrates) and purple acid phosphatase (hydrolysis of phosphate esters). When the body is fighting a bacterial infection, the body sequesters iron inside of cells (mostly stored in the storage molecule ferritin so that it cannot be used by bacteria. Iron distribution is heavily regulated in mammals, both as a defense against bacterial infection as well as the potential biological toxicity of iron. The iron absorbed from the duodenum binds to transferrin, and is carried by blood to different cells. There it gets by an as yet unknown mechanism incorporated into target proteins. [http://www.plosbiology.org/plosonline/?request=get-document&doi=10.1371%2Fjournal.pbio.0000079]. A lengthier article on the system of human iron regulation can be found in the article on human iron metabolism. Good sources of dietary iron include meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed pea, strawberries and farina. Iron provided by dietary supplements is often found as Iron (II) fumarate. The RDA for iron varies considerably based on the age, gender, and source of dietary iron (heme-based iron has higher bioavailability)[http://www.iom.edu/Object.File/Master/7/294/0.pdf]. Also note the section below on precautions.

Isotopes

Naturally occurring iron consists of four isotopes: 5.845% of radioactive ⁵⁴Fe (half-life: >3.1E22 years), 91.754% of stable ⁵⁶Fe, 2.119% of stable ⁵⁷Fe and 0.282% of stable ⁵⁸Fe. ⁶⁰Fe is an extinct radionuclide of long half-life (1.5 million years). Much of the past work on measuring the isotopic composition of Fe has centered on determining ⁶⁰Fe variations due to processes accompanying nucleosynthesis (i.e., meteorite studies) and ore formation. The isotope ⁵⁶Fe is of particular interest to nuclear scientists. A common misconception is that this isotope represents the most stable nucleus possible, and that it thus would be impossible to perform fission or fusion on ⁵⁶Fe and still liberate energy. This is not true, as both ⁶²Ni and ⁵⁸Fe are more stable. In phases of the meteorites Semarkona and Chervony Kut a correlation between the concentration of ⁶⁰Ni, the daughter product of ⁶⁰Fe, and the abundance of the stable iron isotopes could be found which is evidence for the existence of ⁶⁰Fe at time formation of solar system. Possibly the energy released by the decay of ⁶⁰Fe contributed, together with the energy released by decay of the radionuclide ²⁶Al, to the remelting and differentiation of asteroids after their formation 4.6 billion years ago. The abundance of ⁶⁰Ni present in extraterrestrial material may also provide further insight into the origin of the solar system and its early history. Of the stable isotopes, only ⁵⁷Fe has a nuclear spin (−1/2). For this reason, ⁵⁷Fe has application as a spin isotope in chemistry and biochemistry.

Precautions

Excessive dietary iron is toxic, because excess ferrous iron reacts with peroxides in the body, producing free radicals. When iron is in normal quantity, the body's own antioxidant mechanisms can control this process. In excess, uncontrollable quantities of free radicals are produced. The lethal dose of iron in a two-year-old is about three grams of iron. One gram can induce severe poisoning. There are reported cases of children being poisoned by consuming between 10 and 50 tablets of ferrous sulfate over a period of several hours. Over-consumption of iron is the single highest cause of death in children by unintentional ingestion of pharmaceuticals. The DRI lists the Tolerable Upper Intake Level (UL) for adults as 45 mg/day. For children under fourteen years old the UL is 40 mg/day. If iron intake is excessive a number of iron overload disorders can result, such as hemochromatosis. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Blood donors are at special risk of low iron levels and are often recommended to supplement their iron intake. A specific chelating agent called Desferrioxime is used to expell excess iron from the body in case of iron toxicity.

References

- [http://periodic.lanl.gov/elements/26.html Los Alamos National Laboratory — Iron]

External links

- [http://www.webelements.com/webelements/elements/text/Fe/index.html WebElements.com – Iron]
- [http://education.jlab.org/itselemental/ele026.html It's Elemental – Iron]
- [http://hyperphysics.phy-astr.gsu.edu/hbase/nucene/nucbin2.html The Most Tightly Bound Nuclei] Category:Chemical elements Category:Transition metals ko:철 ms:Besi ja:鉄 simple:Iron th:เหล็ก

Sulfur

Sulfur (or sulphur; see spelling below) is the chemical element in the periodic table that has the symbol S and atomic number 16. It is an abundant, tasteless, odorless, multivalent non-metal. Sulfur, in its native form, is a yellow crystaline solid. In nature, it can be found as the pure element or as sulfide and sulfate minerals. It is an essential element for life and is found in several amino acids. Its commercial uses are primarily in fertilizers but it is also widely used in gunpowder, matches, insecticides and fungicides.

Notable characteristics

fungicide At room temperature, sulfur is a soft bright yellow solid. Although sulfur is infamous for its smell - frequently compared to rotten eggs - the odor is actually characteristic of hydrogen sulfide (H₂S); elemental sulfur is odorless. It burns with a blue flame that emits sulfur dioxide, notable for its peculiar suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide and other nonpolar solvents. Common oxidation states of sulfur include −2, +2, +4 and +6. Sulfur forms stable compounds with all elements except the noble gases. Sulfur in the solid state ordinarily exists as a cyclic crown-shaped S₈ molecules. Sulfur has many allotropes besides S₈. Removing one atom from the crown gives S₇, which is responsible for sulfur's distinctive yellow color. Many other rings have been prepared, including S₁₂ and S₁₈. By contrast, its lighter neighbor oxygen only exists in two states of chemical significance: O₂ and O₃. Selenium, the heavier analogue of sulfur can form rings but is more often found as a polymer chain. The crystallography of sulfur is complex. Depending on the specific conditions, the sulfur allotropes form several distinct crystal structures, with rhombic and monoclinic S₈ best known. A noteworthy property is that the viscosity of molten sulfur, unlike most other liquids, increases with temperature due to the formation of polymer chains. However, after a certain temperature is reached, the viscosity is reduced because there is enough energy to break the chains. Amorphous or "plastic" sulfur can be produced through the rapid cooling of molten sulfur. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. This form is metastable at room temperature and gradually reverts back to crystalline form. This process happens within a matter of hours to days but can be rapidly catalyzed by human saliva.

Applications

Sulfur has many industrial uses. Through its major derivative, sulfuric acid (H₂SO₄), sulfur ranks as one of the more important elements used as an industrial raw material. It is of prime importance to every sector of the world's economies. Sulfuric acid production is the major end use for sulfur, and consumption of sulfuric acid has been regarded as one of the best indices of a nation's industrial development. More sulfuric acid is produced in the United States every year than any other industrial chemical. Sulfur is also used in batteries, detergents, the vulcanization of rubber, fungicides, and in the manufacture of phosphate fertilizers. Sulfites are used to bleach paper and as a preservative in wine and dried fruit. Because of its flammable nature, sulfur also finds use in matches, gunpowder, and fireworks. Sodium or ammonium thiosulfate are used as photographic fixing agents. Magnesium sulfate, better known as Epsom salts can be used as a laxative, a bath additive, an exfoliant, or a magnesium supplement for plants. In the late 1700's, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide given off during the process of melting sulfur, the craft of sulfur inlays was soon abandoned.

Biological role

The amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids. This makes sulfur a necessary component of all living cells. Disulfide bonds between polypeptides are very important in protein assembly and structure. Homocysteine and taurine are also sulfur containing amino acids but are not coded for by DNA nor are they part of the primary structure of proteins. Some forms of bacteria use hydrogen sulfide (H₂S) in the place of water as the electron donor in a primitive photosynthesis-like process. Sulfur is absorbed by plants from soil as the sulfate ion. Inorganic sulfur forms a part of iron-sulfur clusters, and sulfur is the bridging ligand in the Cu_A site of cytochrome c oxidase. Sulfur is an important component of coenzyme A

Environmental Impact

The burning of coal and petroleum by industry and power plants liberates huge amounts of sulfur dioxide SO₂, which reacts with atmospheric water and oxygen to produce sulfuric acid. This causes acid rain which lowers the pH of soil and freshwater bodies, resulting in substantial damage to the natural environment and chemical weathering of statues and architecture. Fuel standards increasingly require sulfur to be extracted from fossil fuels to prevent the formation of acid rain. This extracted sulfur is then refined and represents a large portion of sulfur production.

History

fossil fuel Sulfur (Sanskrit, sulvere; Latin sulpur) was known in ancient times, and is referred to in the Biblical Pentateuch (Genesis). English translations of this commonly refer to sulfur as "brimstone", giving rise to the name of 'Fire and brimstone' sermons, which are sermons where hell and eternal damnation for sinners is stressed. It is from this part of the Bible that hell is thought to smell of sulfur. The word itself is almost certainly from the Arabic sufra meaning yellow, from the bright color of the naturally-occurring form. Homer mentioned "pest-averting sulfur" in the 9th century BC and in 424 BC, the tribe of Boeotia destroyed the walls of a city by burning a mixture of coal, sulfur, and tar under them. Sometime in the 12th century, the Chinese invented gun powder which is a mixture of potassium nitrate (K N O₃), carbon, and sulfur. Early alchemists gave sulfur its own alchemical symbol which was a triangle at the top of a cross. In the late 1770s, Antoine Lavoisier helped convince the scientific community that sulfur was an element and not a compound. In 1867 sulfur was discovered in underground deposits in Louisiana and Texas. The overlying layer of earth was quicksand, prohibiting ordinary mining operations. Therefore the Frasch process was utilized.

Occurrence

Frasch process Frasch process, New Zealand]] Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire. These occurrences are the basis for the traditional name brimstone, since sulfur could be found near the brims of volcanic craters. Such volcanic deposits are currently exploited in Indonesia, Chile, and Japan. Significant desposits of elemental sulfur also exist in salt domes along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. The sulfur in these deposits is believed to come from the action of anaerobic bacteria on sulfate minerals, especially gypsum. Such deposits are the basis for commercial production in the United States, Poland, Russia, Turkmenistan, and Ukraine. Sulfur extracted from oil, gas and the Athabasca Oil Sands has become a glut on the market, with huge stockpiles of sulfur in existence throughout Alberta. Athabasca Oil Sands]] Common naturally-occurring sulfur compounds include the metal sulfides, such as pyrite (iron sulfide), cinnabar (mercury sulfide), Galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the metal sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). Hydrogen sulfide is the gas responsible for the odor of rotten eggs. It occurs naturally in volcanic emissions, such as from hydrothermal vents, and from bacterial action on decaying sulfur-containing organic matter. The distinctive colors of Jupiter's volcanic moon, Io, are from various forms of molten, solid and gaseous sulfur. There is also a dark area near the Lunar crater Aristarchus that may be a sulfur deposit. Sulfur is also present in many types of meteorites.

Compounds

Hydrogen sulfide has the characteristic smell of rotten eggs. Dissolved in water, hydrogen sulfide is acidic and will react with metals to form a series of metal sulfides. Natural metal sulfides are common, especially those of iron. Iron sulfide is called pyrite, the so called fool's gold. Interestingly, pyrite can show semiconductor properties.[http://home.earthlink.net/~lenyr/iposc.htm] Galena, a naturally occurring lead sulfide, was the first semiconductor discovered, and found a use as a signal rectifier in the "cat's whiskers" of early crystal radios. Many of the unpleasant odors of organic matter are based on sulfur-containing compounds such as ethyl and methyl mercaptan used to scent natural gas so that leaks are easily detectable. The odor of garlic and "skunk stink" are also caused by sulfur containing organic compounds. However, not all organic sulfur compounds smell unpleasant, margin, a sulfur containing terpene is responsible for the characteristic scent of grapefruit. Polymeric sulfur nitride has metallic properties even though it does not contain any metal atoms. This compound also has unusual electrical and optical properties. This polymer can be made from tetrasulfur tetranitride S₄N₄. Other important compounds of sulfur include: INORGANIC:
- Sulfides (S^2-) are simple compounds of sulfur with some other chemical element.
- Sulfites (SO₃^2-), the salts of sulfurous acid, H₂SO₃, created by dissolving SO₂ in water. Sulfurous acid and the corresponding sulfites are fairly strong reducing agents. Other compounds derived from SO₂ include the pyrosulfite or metabisulfite ion (S₂O₅²⁻).
- Sulfates (SO₄^2-), the salts of sulfuric acid. Related to this, sulfuric acid also reacts with SO₃ in equimolar ratios to form pyrosulfuric acid (H₂S₂O₇).
- Thiosulfates (sometimes refered as thiosulfites or "hyposulfites") (S₂O₃²⁻). Thiosulfates are used in photographic fixing (HYPO)as reducing agents and ammonium thiosulfate is being investigated as a cyanide replacement in leaching gold.[http://doccopper.tripod.com/gold/AltLixiv.html]
- Sodium dithionite, Na₂S₂O₄ from hyposulfurous/dithionous acid, a powerful reducing agent.
- Sodium dithionate (Na₂S₂O₆)
- Polythionic acids (H₂S_nO₆), where n can range from 3 to 80.
- Peroxymonosulfuric acid (H₂SO₅) and peroxydisulfuric acids (H₂S₂O₈), made from the action of SO₃ on concentrated H₂O₂, and H₂SO₄ on concentrated H₂O₂ respectively.
- Sodium polisulfides (Na₂S_x)
- Sulfur hexafluoride, SF₆, a heavy, gaseous, non-reactive and non-toxic propellant
- Tetrasulfur tetranitride S₄N₄.
- Thiocyanates are compounds containing the thiocyanate ion, SCN^-. Related to this there is thiocyanogen, (SCN)₂. ORGANIC:
- dimethylsulfoniopropionate (DMSP; (CH₃ )₂S⁺CH₂CH₂COO^-) which is the central component of the marine organic sulfur cycle.
- A thioether is a molecule with the form R-S-R′, where R and R′ are organic groups. These are the sulfur equivalents of ethers.
- A thiol (also known as a mercaptan) is a molecule with an -SH functional group. These are the sulfur equivalents of alcohols.
- A thiolate ion has an -S^- functional group attached. These are the sulfur equivalent of alkoxide ions.
- A sulfoxide is a molecule with an R-S(=O)-R′ functional group where R and R′ are organic groups. A common example of a sulfoxide is DMSO.
- A sulfone is a molecule with an R-S(=O)-R′ functional group where R and R′ are organic groups.
- Lawesson's reagent is a chemical reagent which can remove oxygen from other organic molecules and replace it with sulfur.
- Napthalen-1,8-diyl 1,3,2,4-dithiadiphosphetane 2,4-disulfide

Isotopes

Sulfur has 18 isotopes, of which four are stable: ³²S (95.02%), ³³S (0.75%), ³⁴S (4.21%), and ³⁶S (0.02%). Other than ³⁵S, the radioactive isotopes of sulfur are all short lived. Sulfur-35 is formed from cosmic ray spallation of argon-40 in the atmosphere. It has a half-life of 87 days. When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the dS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The dC-13 and dS-34 of co-existing carbonates and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation. In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites also contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can also be used in systems where there is sufficient variation in the S-34 of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different dS-34 values from lakes believed to be dominated by watershed sources of sulfate.

Precautions

Carbon disulfide, Carbon oxysulfide, hydrogen sulfide, and sulfur dioxide should all be handled with care. Although sulfur dioxide is sufficiently safe to be used as a food additive in small amounts, at high concentrations it reacts with moisture to form sulfurous acid which in sufficient quantities may harm the lungs, eyes or other tissues. In creatures without lungs such as insects or plants, it otherwise prevents respiration. Hydrogen sulfide is quite toxic (more toxic than cyanide). Although very smelly at first, it quickly deadens the sense of smell, so potential victims may be unaware of its presence until it is too late.

Spelling

The element has traditionally been spelled sulphur in the United Kingdom and India, but sulfur in the United States and Canada, while both spellings are used in Australia and New Zealand. The IUPAC adopted the spelling "sulfur" in 1990, as did the Royal Society of Chemistry Nomenclature Committee in 1992. This spelling has begun to replace its variant in educated circles, unlike aluminum, which did not stick outside the US and Canada.

References

- [http://periodic.lanl.gov/elements/16.html Los Alamos National Laboratory – Sulfur]
- R. Steudel (ed.): Elemental Sulfur and Sulfur-Rich Compounds (part I & II), Topics in Current Chemistry Vol. 230 & 231, Springer, Berlin 2003.

External links

- [http://library.tedankara.k12.tr/chemistry/vol2/allotropy/z129.htm Sulfur phase diagram.]
- [http://www.webelements.com/webelements/elements/text/S/index.html WebElements.com – Sulfur] Category:Chemical elements Category:Nonmetals Category:Chalcogens Category:Pyrotechnic chemicals ko:황 ja:硫黄 simple:Sulfur th:กำมะถัน

Oxygen

Oxygen is a chemical element in the periodic table. It has the symbol O and atomic number 8. The element is very common, found not only on Earth but throughout the universe, usually covalently bonded with other elements. Unbound oxygen (usually called molecular oxygen, O₂, a diatomic molecule) first appeared on Earth during the Paleoproterozoic era (between 2500 million years ago and 1600 million years ago) and as a product of the metabolic action of early anaerobes (archaea and bacteria). The presence of free oxygen drove most of the organisms then living to extinction. The atmospheric abundance of free oxygen in later geological epochs and up to the present has been largely driven by photosynthetic organisms, roughly three quarters by phytoplankton and algae in the oceans and one quarter from terrestrial plants.

Characteristics

At standard temperature and pressure, oxygen is mostly found as a gas consisting of a diatomic molecule with the chemical formula O₂. O₂ has two energetic forms:
- The low-energy predominant single-bonded diradical triplet oxygen. This native diradical quality of oxygen contributes to its destructive chemical nature. This form is stabilized by the degeneracy effect.
- The high-energy double-bonded molecule singlet oxygen. Oxygen is a major component of air, produced by plants during photosynthesis, and is necessary for aerobic respiration in animals. The word oxygen derives from two words in Greek, οξυς (oxys) (acid, sharp) and γεινομαι (geinomai) (engender). The name "oxygen" was chosen because, at the time it was discovered in the late 18th century, it was believed that all acids contained oxygen. The definition of acid has since been revised to not require oxygen in the molecular structure. Liquid O₂ and solid O₂ have a light blue color and both are highly paramagnetic. Liquid O₂ is usually obtained by the fractional distillation of liquid air. Liquid and solid O₃ (ozone) have a deeper color of blue. A recently discovered allotrope of oxygen, tetraoxygen (O₄), is a deep red solid that is created by pressurizing O₂ to the order of 20 GPa. Its properties are being studied for use in rocket fuels and similar applications, as it is a much more powerful oxidizer than either O₂ or O₃.

Applications

Liquid oxygen finds use as an oxidizer in rocket propulsion. Oxygen is essential to respiration, so oxygen supplementation has found use in medicine (as oxygen therapy). People who climb mountains or fly in airplanes sometimes have supplemental oxygen supplies (as air). Oxygen is used in welding (such as the oxyacetylene torch), and in the making of steel and methanol. Oxygen presents two absorption bands centered in the wavelengths 687 and 760 nanometers. Some scientists have proposed to use the measurement of the radiance coming from vegetation canopies in those oxygen bands to characterize plant health status from a satellite platform. This is because in those bands, it is possible to discriminate the vegetation's reflectance from the vegetation's fluorescence, which is much weaker. The measurement presents several technical difficulties due to the low signal to noise ratio and due to the vegetation's architecture, but it has been proposed as possibility to monitor the carbon cycle from satellite, thus in a global scale. Oxygen, as a mild euphoric, has a history of recreational use that extends into modern times. Oxygen bars can be seen at parties to this day. In the 19th century, oxygen was often mixed with nitrous oxide to promote an analgesic effect; indeed, such a mixture (Entonox) is commonly used in medicine today.

History

Oxygen was first discovered by Michał Sędziwój, Polish alchemist and philosopher in late 16th century. Sędziwój assumed the existence of oxygen by warming nitre (saltpeter). He thought of the gas given off as "the elixir of life". Oxygen was again discovered by the Swedish pharmacist Carl Wilhelm Scheele sometime before 1773, but the discovery was not published until after the independent discovery by Joseph Priestley on August 1, 1774, who called the gas dephlogisticated air (see phlogiston theory). Priestley published his discoveries in 1775 and Scheele in 1777; consequently Priestley is usually given the credit. It was named by Antoine Laurent Lavoisier after Priestley's publication in 1775.

Occurrence

Oxygen is the second most common component of the earth's atmosphere (20.947% by volume).

Compounds

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements (which is the origin of the original definition of oxidation). The only elements to escape the possibility of oxidation are a few of the noble gases. The most famous of these oxides is dihydrogen monoxide, or water (H₂O). Other well known examples include compounds of carbon and oxygen, such as carbon dioxide (CO₂), alcohols (R-OH), aldehydes, (R-CHO), and carboxylic acids (R-COOH). Oxygenated radicals such as chlorates (ClO₃⁻), perchlorates (ClO₄⁻), chromates (CrO₄²⁻), dichromates (Cr₂O₇²⁻), permanganates (MnO₄⁻), and nitrates (NO₃⁻) are strong oxidizing agents in and of themselves. Many metals such as iron bond with oxygen atoms, iron (III) oxide (Fe₂O₃). Ozone (O₃) is formed by electrostatic discharge in the presence of molecular oxygen. A double oxygen molecule (O₂)₂ is known and is found as a minor component of liquid oxygen. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Isotopes

Oxygen has fifteen known isotopes with atomic masses ranging from 12 to 26. Three of them are stable and twelve are radioactive. The radioisotopes all have half lives of less than three minutes. The stable isotopes have mass numbers of 16, 17 and 18, of which oxygen-16 is the most common (over 99%).

Precautions

Oxygen can be toxic at elevated partial pressures (i.e. high relative concentrations). This is important in some forms of scuba diving, such as with a rebreather. Certain derivatives of oxygen, such as ozone (O₃), singlet oxygen, hydrogen peroxide, hydroxyl radicals and superoxide, are also highly toxic. The body has developed mechanisms to protect against these toxic species. For instance, the naturally-occurring glutathione can act as an antioxidant, as can bilirubin which is normally a breakdown product of hemoglobin. Highly concentrated sources of oxygen promote rapid combustion and therefore are fire and explosion hazards in the presence of fuels. This is true as well of compounds of oxygen such as chlorates, perchlorates, dichromates, etc. Compounds with a high oxidative potential can often cause chemical burns. The fire that killed the Apollo 1 crew on a test launchpad spread so rapidly because the pure oxygen atmosphere was at normal atmospheric pressure instead of the one third pressure that would be used during an actual launch. (See partial pressure.) Oxygen derivatives are prone to form free radicals, especially in metabolic processes. Because they can cause severe damage to cells and their DNA, they are thought to be related to cancer and aging.

References

- [http://periodic.lanl.gov/elements/8.html Los Alamos National Laboratory – Oxygen]
- [http://physics.nist.gov/cgi-bin/AtData/main_asd Nist atomic spectra database]
- [http://chartofthenuclides.com/default.html Nuclides and Isotopes Fourteenth Edition]: Chart of the Nuclides, General Electric Company, 1989

External links

- [http://www.priestleysociety.net Priestley Society, Dedicated to Joseph Priestley the man who discovered oxygen]
- [http://www.best-home-remedies.com/minerals/oxygen.htm Oxygen - Benefits, Deficiency Symptoms And Food Sources]
- [http://www.josephpriestley.info Joseph Priestley Information Website, about the man who discovered oxygen]
- [http://periodic.lanl.gov/elements/8.html Los Alamos National Laboratory – Oxygen]
- [http://www.webelements.com/webelements/elements/text/O/index.html WebElements.com – Oxygen]
- [http://education.jlab.org/itselemental/ele008.html It's Elemental – Oxygen]
- [http://members.tripod.com/tjaartdb0/html/oxygen_toxicity.html Oxygen Toxicity]
- [http://www.uigi.com/oxygen.html Oxygen (O2) Properties, Uses, Applications]
- [http://www.compchemwiki.org/index.php?title=Oxygen Computational Chemistry Wiki]
- [http://koti.mbnet.fi/antitz/dime/en Tests with liquid oxygen :-)] Category:Nonmetals Category:Chalcogens als:Sauerstoff ko:산소 ms:Oksigen ja:酸素 simple:Oxygen th:ออกซิเจน

Ionic compound

In chemistry, an ionic compound is a chemical compound in which ions are held together in a lattice structure by ionic bonds. To form an ionic compound, there needs to be at least one metal and one non-metal. The metal element is usually the positive charge and the non-metal element is a negative charge. Ions can be single atoms, as in common table salt sodium chloride, or more complex groups such as calcium carbonate. But to be considered ions, they must carry a positive or negative charge. Thus, in an ionic bond, one 'bonder' must have a positive charge and the other a negative one. By sticking to each other, they resolve, or partially resolve, their separate charge imbalances. Positive to positive and negative to negative ionic bonds do not occur. (For a real world analogy, experiment with a pair of bar magnets.) Ionic compounds are generally high melting and boiling. They have good electrical conductivity when molten or in solution. While ionic inorganic compounds are solids at room temperature and will usually form crystals, organic ionic liquids are increasingly of interest. When an ionic compound is named, the cation is named first and then the anion. When an elemental anion is named the suffix, -ide is added to the name of the element. There are two common types of cations: Type I and Type II. Type I cations have only one charge and their name is simply listed when the compound is named. Type II cations have more than one charge and when the ionic compound is named, a Roman numeral is used to denote the charge of the cation. In addition, there are common polyatomic anions which do not have suffixes in their name such as hypochlorite (ClO-). Category:Chemical compounds Category:Ions simple:Ionic compound

Hydrate

Hydrates are compounds formed by the union of water with some other substance, generally forming a neutral body, as certain crystallized salts. If the water is heavy water, where the hydrogen consists of the isotope deuterium, then the term deuterate may be used in place of hydrate. These substances do not contain water as such, but have their constituents (hydrogen, oxygen, hydroxyl) so arranged that water may be eliminated. Hence, hydrates are derivatives of, or compounds with, hydroxyl. Such a substance is considered to be anhydrous when it is not in such a union with water. The corresponding anhydrate also does not contain any water The notation of anhydrous compound, where n is the number of water molecules per molecule of anhydrous salt, is commonly used to show that a salt is hydrated. The n is usually a low integer, though it is possible for fractional values to exist. In a monohydrate n is one, in a hexahydrate n is 6 etc. Examples include borax, chloral hydrate, clathrate hydrates (a class of solid hydrates of gases), and chalcanthite. Gas hydrates are clathrate hydrates: water ice with gas molecules trapped within. When the gas is methane it is called a methane hydrate. Category:Chemical compounds
- Category:Oxygen compounds Category:Hydrogen compounds

Hydrogen

|- | Critical temperature || 32.19 K |- | Critical pressure || 1.315 MPa |- | Critical density || 30.12 g/L (Bohr radius) Hydrogen (Latin: hydrogenium, from Greek: hydro: water, genes: forming) is a chemical element in the periodic table that has the symbol H and atomic number 1. At standard temperature and pressure it is a colorless, odorless, nonmetallic, univalent, highly flammable diatomic gas. Hydrogen is the lightest and most abundant element in the universe. It is present in water, all organic compounds (rare exceptions exist, like buckminsterfullerene) and in all living organisms. Hydrogen is able to react chemically with most other elements. Stars in their main sequence are overwhelmingly composed of hydrogen in its plasma state. The element is used in ammonia production, as a lifting gas, as an alternative fuel, and more recently as a power source of fuel cells. Despite its ubiquity in the universe, hydrogen is surprisingly hard to produce in large quantities on the Earth. In the laboratory, the element is prepared by the reaction of acids on metals such as zinc. The electrolysis of water is a simple method of producing hydrogen, but is economically inefficient for mass production. Large-scale production is usually achieved by steam reforming natural gas. Scientists are now researching new methods for hydrogen production; if they succeed in developing a cost-efficient method of large-scale production, hydrogen may become a viable alternative to greenhouse-gas-producing fossil fuels. One of the methods under investigation involves use of green algae; another promising method involves the conversion of biomass derivatives such as glucose or sorbitol at low temperatures using a catalyst. Yet another method is the "steaming" of Carbon, whereby hydrocarbons are broken down with heat to release hydrogen.

Basic features

Hydrogen is the lightest chemical element; its most common isotope comprises just one negatively charged electron, distributed around a positively charged proton (the nucleus of the atom). The electron is bound to the proton by the Coulomb force, the electrical force that one stationary, electrically charged nanoparticle exerts on another. The hydrogen atom has special significance in quantum mechanics as a simple physical system for which there is an exact solution to the Schrödinger equation; from that equation, the experimentally observed frequencies and intensities of the hydrogen's spectral lines can be calculated. Spectral lines are dark or bright lines in an otherwise uniform and continuous spectrum, resulting from an excess or deficiency of photons in a narrow frequency range, compared with the nearby frequencies. At standard temperature and pressure, hydrogen forms a diatomic gas, H₂, with a boiling point of only 20.27 K and a melting point of 14.02 K. Under extreme pressures, such as those at the center of gas giants, the molecules lose their identity and the hydrogen becomes a liquid metal. Under the extremely low pressure in space—virtually a vacuum—the element tends to exist as individual atoms, simply because there is no way for them to combine. However, clouds of H₂ and singular hydrogen atoms are said to form in H I and H II regions and are associated with star formation, however the existence of singular hydrogen atoms is disputed.. Hydrogen plays a vital role in powering stars through the proton–proton and carbon–nitrogen cycle. These are nuclear fusion processes, which release huge amounts of energy in stars and other hot celestial bodies as hydrogen atoms combine into helium atoms. H₂ is highly soluble in water, alcohol, and ether. It has a high capacity for adsorption, in which it is attached to and held to the surface of some substances. It is an odorless, tasteless, colorless, and highly flammable gas that burns at concentrations as low as 4% H₂ in air. It reacts violently with chlorine and fluorine, forming hydrohalic acids that can damage the lungs and other tissues. When mixed with oxygen, hydrogen explodes on ignition. A unique property of hydrogen is that its flame is completely invisible in air. This makes it difficult to tell if a leak is burning, and carries the added risk that it is easy to walk into a hydrogen fire inadvertently. See also: hydrogen atom.

Applications

Large quantities of hydrogen are needed in the chemical and petrolium industries, notably in the Haber process for the production of ammonia, which by mass ranks as the world's fifth most highly produced industrial compound. Hydrogen is used in the hydrogenation of fats and oils (into items such as margarine), and in the production of methanol. Hydrogen is used in hydrodealkylation, hydrodesulfurization, and hydrocracking. The element has several other important uses.
- The element is used in the manufacture of hydrochloric acid, in welding processes, and in the reduction of metallic ores.
- It is an ingredient in rocket fuels.
- It is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas.
- Liquid hydrogen is used in cryogenic research, including superconductivity studies.
- Since hydrogen is 14.5 times lighter than air, it was once widely used as a lifting agent in balloons and airships. However, this use was curtailed when the Hindenburg disaster convinced the public that the gas was too dangerous for this purpose.
- Deuterium, an isotope of hydrogen (hydrogen-2), is used in nuclear fission applications as a moderator to slow neutrons, and in nuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects.
- Tritium (hydrogen-3), produced in nuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints. There are no "hydrogen wells" or "hydrogen mines" on Earth, so hydrogen cannot be considered a primary energy source like fossil fuels or uranium. Hydrogen can however be burned in internal combustion engines, an approach advocated by BMW's experimental hydrogen car. However, it is currently difficult and dangerous to store and handle in sufficient quantity for motor fuel use. Hydrogen fuel cells are being investigated as mobile power sources with lower emissions than hydrogen-burning internal combustion engines. The low emissions of hydrogen in internal combustion engines and fuel cells are currently offset by the pollution created by hydrogen production. This may change if the substantial amounts of electricity required for water electrolysis can be generated primarily from low pollution sources such as nuclear energy or wind. Research is being conducted on hydrogen as a replacement for fossil fuels. It could become the link between a range of energy sources, carriers and storage. Hydrogen can be converted to and from electricity (solving the electricity storage and transport issues), from bio-fuels, and from and into natural gas and diesel fuel. All of this can theoretically be achieved with zero emissions of CO₂ and toxic pollutants.

History

Hydrogen was first produced by Theophratus Bombastus von Hohenheim (1493–1541)—also known as Paracelsus—by mixing metals with acids. He was unaware that the explosive gas produced by this chemical reaction was hydrogen. In 1671, Robert Boyle described the reaction between two iron fillings and dilute acids, which results in the production of gaseous hydrogen. In 1766, Henry Cavendish was the first to recognize hydrogen as a discrete substance, by identifying the gas from this reaction as "inflammable" and finding that the gas produces water when burned in air. Cavendish stumbled on hydrogen when experimenting with acids and mercury. Although he wrongly assumed that hydrogen was a compound of mercury—and not of the acid—he was still able to accurately describe several key properties of hydrogen. Antoine Lavoisier gave the element its name and proved that water is composed of hydrogen and oxygen. One of the first uses of the element was for balloons. The hydrogen was obtained by mixing sulfuric acid and iron. Harold C. Urey discovered Deuterium, an isotope of hydrogen, by repeated distilling the same sample of water. For this discovery, Urey received the Nobel prize for in 1934. In the same year, the third isotope, tritium, was discovered. Because of its relatively simple structure, hydrogen has often been used in models of how an atom works.

Electron energy levels

The ground state energy level of the electron in a Hydrogen atom is 13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm. With the Bohr Model the energy levels of Hydrogen can be calculated fairly accurately. This is done by modeling the electron as revolving around the proton, much like the earth revolving around the sun. Except the sun holds earth in orbit with the force of gravity, but the proton holds the electron in orbit with the force of electromagnetism. Another difference between the Earth-Sun system and the Electron-Proton system is that, in this model, due to quantum mechanics the electron is allowed to only be at very specific distances from the proton. Modeling the hydrogen atom in this fashion yields the correct energy levels and spectrum.

Occurrence

quantum mechanics.]] Hydrogen is the most abundant element in the universe, making up 75% of normal matter by mass and over 90% by number of atoms. This element is found in great abundance in stars and gas giant planets. It is very rare in the Earth's atmosphere (1 ppm by volume), because being the lightest gas causes it to escape Earth's gravity, though when compounds are considered, it is the tenth most abundant element on Earth. The most common source for this element on Earth is water, which is composed two parts hydrogen to one part oxygen (H₂O). Other sources include most forms of organic matter (currently all known life forms) including coal, natural gas, and other fossil fuels. Methane (CH₄) is an increasingly important source of hydrogen. Throughout the Universe, hydrogen is mostly found in the plasma state whose properties are quite different to molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity, even when the gas is only partially ionised. The charged particles are highly influenced by magnetic and electric fields, for example, in the Solar Wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen can be prepared in several different ways: steam on heated carbon, hydrocarbon decomposition with heat, reaction of a strong base in an aqueous solution with aluminium, water electrolysis, or displacement from acids with certain metals. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. At high temperatures (700–1100 °C), steam reacts with methane to yield carbon monoxide and hydrogen. :CH₄ + H₂O → CO + 3 H₂ Additional hydrogen can be recovered from the carbon monoxide through the water-gas shift reaction: :CO + H₂O → CO₂ + H₂

Compounds

The lightest of all gases, hydrogen combines with most other elements to form compounds. Hydrogen has an electronegativity of 2.2, so it forms compounds where it is the more nonmetallic and where it is the more metallic element. The former are called hydrides, where hydrogen either exists as H^- ions or just as a solute within the other element (as in palladium hydride). The latter tend to be covalent, since the H⁺ ion would be a bare nucleus and so has a strong tendency to pull electrons to itself. These both form acids. Thus even in an acidic solution one sees ions like hydronium (H₃O⁺) as the protons latch on to something. Although exotic on earth, one of the most common ions in the universe is the H₃⁺ ion. Hydrogen combines with oxygen to form water, H₂O, and releases a lot of energy in doing so, burning explosively in air. Deuterium oxide, or D₂O, is commonly referred to as heavy water. Hydrogen also forms a vast array of compounds with carbon. Because of their association with living things, these compounds are called organic compounds, and the study of the properties of these compounds is called organic chemistry. organic chemistry

Forms

Under normal conditions, hydrogen gas is a mix of two different kinds of molecules which differ from one another by the relative spin of the nuclei. These two forms are known as ortho- and para-hydrogen (this is different from isotopes, see below). In ortho-hydrogen the nuclear spins are parallel (form a triplet), while in para they are antiparallel (form a singlet). At standard conditions hydrogen is composed of about 25% of the para form and 75% of the ortho form (the so-called "normal" form). The equilibrium ratio of these two forms depends on temperature, but since the ortho form has higher energy (is an excited state), it cannot be stable in its pure form. In low temperatures (around boiling point), the equilibrium state is comprised almost entirely of the para form. The conversion process between the forms is slow, and if hydrogen is cooled down and condensed rapidly, it contains large quantities of the ortho form. It is important in preparation and storage of liquid hydrogen, since the ortho-para conversion produces more heat than the heat of its evaporation, and a lot of hydrogen can be lost by evaporation in this way during several days after liquefying. Therefore, some catalysts of the ortho-para conversion process are used during hydrogen cooling. The two forms have also slightly different physical properties. For example, the melting and boiling points of parahydrogen are about 0.1 K lower than of the "normal" form.

Isotopes

Hydrogen is the only element that has different names for its isotopes. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used, although one element, radon, has a name that originally applied to only one of its isotopes.) The symbols D and T (instead of ²H and ³H) are sometimes used for deuterium and tritium, although this is not officially sanctioned. (The symbol P is already in use for phosphorus and is not available for protium.) ;¹H The most common isotope of hydrogen, this stable isotope has a nucleus consisting of a single proton; hence the descriptive, although rarely used, name protium. The spin of a protium atom is 1/2+. ;²H The other stable isotope is deuterium, with an extra neutron in the nucleus. Deuterium comprises 0.0184%–0.0082% of all hydrogen (IUPAC); ratios of deuterium to protium are reported relative to the VSMOW standard reference water. The spin of a deuterium atom is 1+. ;³H The third naturally occurring hydrogen isotope is the radioactive tritium. The tritium nucleus contains two neutrons in addition to the proton. It decays through beta decay and has a half-life of 12.32 years. Tritium occurs naturally due to cosmic rays interacting with atmospheric gases. Like ordinary hydrogen, tritium reacts with the oxygen in the atmosphere to form T₂O. This radioactive "water" molecule constantly enters the Earth's seas and lakes in the form of slightly radioactive rain, but its half-life is short enough to prevent a buildup of hazardous radioactivity. The spin of a tritium atom is 1/2+. ;⁴H Hydrogen-4 was synthesized by bombarding tritium with fast-moving deuterium nuclei. It decays through neutron emission and has a half-life of 9.93696x10^-23 seconds. The spin of a hydrogen-4 atom is 2-. ;⁵H In 2001 scientists detected hydrogen-5 by bombarding a hydrogen target with heavy ions. It decays through neutron emission and has a half-life of 8.01930x10^-23 seconds. ;⁶H Hydrogen-6 decays through triple neutron emission and has a half-life of 3.26500^-22 seconds. ;⁷H In 2003 hydrogen-7 was created ([http://physicsweb.org/articles/news/7/3/3 article]) at the RIKEN laboratory in Japan by colliding a high-energy beam of helium-8 atoms with a cryogenic hydrogen target and detecting tritons—the nuclei of tritium atoms—and neutrons from the breakup of hydrogen-7, the same method used to produce and detect hydrogen-5.

References

# # # # # #
- [http://www.riken.go.jp/engn/r-world/research/lab/wako/ribeam/ RIKEN Beam Science Laboratory, Japan - Heavy hydrogen research]
- [http://chartofthenuclides.com/default.html Nuclides and Isotopes] Fourteenth Edition: Chart of the Nuclides, General Electric Company, 1989 ;Book references:
-
-
-
-

External links

- [http://www.hydropole.ch/Hydropole/Intro/Phasediag.gif Hydrogen phase diagram.]
- [http://www.compchemwiki.org/index.php?title=Hydrogen Computational Chemistry Wiki] Category:Nonmetals Category:Fuels Category:Chemical elements ko:수소 ms:Hidrogen ja:水素 simple:Hydrogen th:ไฮโดรเจน

Vitriol

Vitriol is the name that alchemists gave to sulfuric acid. The name was also used for various sulfate salts, such as copper(II) sulfate (blue vitriol, or rarely Roman vitriol), zinc sulfate (white vitriol), iron(II) sulfate (green vitriol), iron(III) sulfate (vitriol of Mars), or cobalt(II) sulfate (red vitriol). Oil of vitriol is concentrated sulfuric acid so named due to its oily appearance. Vitriol is also a quality of abusive or malicious forms of speech or feelings.

Extraction

In antiquity, the vitriol salts were extracted from the runoff that collected inside mines of sulfide ores; the sulfates were formed naturally by the action of air on the wet sulfide minerals, and washed down by percolating water.

Uses

Vitriol was the most important alchemical substance, intended to be used as a philosopher's stone. Highly purified vitriol was used as a medium to react substances in. This was largely because the acid does not react with gold, often the final aim of alchemical processes. The word Vitriol is formed from the initial letters of the alchemical motto VISITA� INTERIORA� TERRA� RECTIFICANDO� INVENIES� OCCULTUM� LAPIDEM (Visit the interior of the earth and rectifying (i.e. purifying) you will find the hidden/secret stone -- the reference is evidently to the legendary Philosopher's Stone).

Manufacture of sulfuric acid

The famous Persian alchemist Al-Razi (864-930) discovered sulfuric acid by the dry distillation of vitriol salts, thus setting in motion a chain of discoveries that would form the foundation of modern chemistry and chemical engineering. (Nowadays the reverse process is generally used, namely the metal sulfates are made by reacting oxide or other metal compound with the acid, which is obtained by other means).

Agriculture

Blue (copper) and, to a lesser extent, white (zinc) vitriol are still occasionally used as chemical defensives in agriculture. In typical applications, a solution of the vitriol is mixed with lime (calcium hydroxide) to produce a fine copper hydroxide suspension, which is sprayed on the plant.

Iron-gall ink

Green (iron(II)) vitriol was much used in the middle ages to make writing iron-gall nut ink.

External links

- [http://www.triad-publishing.com/stone20e.html Triad Publishing's Article] Category:Alchemy

Monoclinic

In crystallography, the monoclinic crystal system is one of the 7 lattice point groups. A crystal system is described by three vectors. In the monoclinic system, the crystal is described by vectors of unequal length, as in the orthorhombic system. They form a rectangular prism with a parallelogram as base. Hence two pairs of vectors are perpendicular, while the third pair make an angle other than 90°. There exist two monoclinic Bravais lattices: the simple monoclinic and the centered monoclinic lattices, with layers with a rectangular and rhombic lattice, respectively. The crystal classes that fall under this crystal system are listed below, followed by their representations in international notation and Schoenflies notation, and mineral examples. The number of space groups for each crystal class is 6, 3, and 4, respectively. The three monoclinic hemimorphic space groups are as follows:
- a prism with as cross-section wallpaper group p2
- ditto with screw axes instead of axes
- ditto with screw axes as well as axes, parallel, in between; in this case an additional translation vector is one half of a translation vector in the base plane plus one half of a perpendicular vector between the base planes The four monoclinic hemihedral space groups include:
- those with pure reflection at the base of the prism and halfway
- those with glide planes instead of pure reflection planes; the glide is one half of a translation vector in the base plane
- those with both in between each other; in this case an additional translation vector is this glide plus one half of a perpendicular vector between the base planes Category:Crystallography

Crystal

:This article is about the form of solid matter. For other uses of this word, see Crystal (disambiguation). Crystal (disambiguation) A crystal is a solid in which the constituent atoms, molecules, or ions are packed in a regularly ordered, repeating pattern extending in all three spatial dimensions. Generally, fluid substances form crystals when they undergo a process of solidification. Under ideal conditions, the result may be a single crystal, where all of the atoms in the solid fit into the same lattice or crystal structure but, generally, many crystals form simultaneously during solidification, leading to a polycrystalline solid. For example, most metals encountered in everyday life are polycrystals. Crystals are often symmetrically intergrown to form crystal twins. Which crystal structure the fluid will form depends on the chemistry of the fluid, the conditions under which it is being solidified, and also on the ambient pressure. The process of forming a crystalline structure is often referred to as crystallization. pressure While the process of cooling usually results in the generation of a crystalline material, under certain conditions the fluid may be frozen in a noncrystalline state. In most cases, this involves cooling the fluid so rapidly that atoms cannot travel to their lattice sites before they lose mobility. A noncrystalline material, which has no long-range order, is called an amorphous, vitreous, or glassy material. It is also often referred to as an amorphous solid, although there are distinct differences between solids and glasses: most notably, the process of forming a glass does not release the latent heat of fusion. For this reason, many scientists consider glassy materials to be viscous liquids rather than solids, although this is a controversial topic; see the entry on glass for more details. glass Crystalline structures occur in all classes of materials, with all types of chemical bonds. Almost all metal exists in a polycrystalline state; amorphous or single-crystal metals must be produced synthetically, often with great difficulty. Ionically bonded crystals can form upon solidification of salts, either from a molten fluid or when it condenses from a solution. Covalently bonded crystals are also very common, notable examples being diamond, silica, and graphite. Polymer materials generally will form crystalline regions, but the lengths of the molecules usually prevents complete crystallization. Weak Van der Waals forces can also play a role in a crystal structure; for example, this type of bonding loosely holds together the hexagonal-patterned sheets in graphite. Most crystalline materials have a variety of crystallographic defects. The types and structures of these defects can have a profound effect on the properties of the materials. crystallographic defect crystallographic defect.]] While the term "crystal" has a precise meaning within materials science and solid-state physics, colloquially "crystal" refers to solid objects that exhibit well-defined