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Iron (II) Sulfate

Iron (II) sulfate

Iron(II) sulfate, also known as ferrous sulfate and as copperas (FeSO4) 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, Fe2- (e.g. Fe(CO)42-,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

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.

See also


- Winkler test for dissolved oxygen for instructions on how to determine the amount of oxygen dissolved in fresh water.
- Combustion
- Oxidation
- Oxygen Catastrophe in geology
- The role of oxygen as a diving breathing gas
- Oxygen depletion aquatic ecology
- Ozone layer

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

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

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.

See also


- Abu Musa Jabir Ibn Hayyan
- Sulfur
- Sulfuric acid

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 and often pleasing geometric shapes. Various shapes of such crystals are found in nature. The shape of these crystals is dependent on the types of molecular bonds between the atoms to determine the structure, as well as on the conditions under which they formed. Snowflakes, diamonds, and common salt are common examples of crystals. Some crystalline materials may exhibit special electrical properties such as the ferroelectric effect or the piezoelectric effect. The behaviour of light in crystals is described by crystal optics. In periodic dielectric structures a range of unique optical properties can be expected as described in photonic crystals. Crystallography is the scientific study of crystals and crystal formation.

See also


- Crystal habit
- Crystal structure
- Crystallite
- Crystallization processes
- Liquid crystal
- Quasicrystal
- Seed crystal
- Single crystal

External links


- [http://www.rockhounds.com/rockshop/xtal/index.html Introduction to Crystallography and Mineral Crystal Systems]
- [http://www.iucr.ac.uk/iucr-top/comm/cteach/pamphlets.html Crystallographic Teaching Pamphlets]
- [http://cst-www.nrl.navy.mil/lattice/spcgrp/ Crystal Lattice Structures] ja:結晶

Molecular weight

The molecular mass of a substance (less accurately called molecular weight and abbreviated as MW) is the mass of one molecule of that substance, relative to the unified atomic mass unit u (equal to 1/12 the mass of one atom of carbon-12). The molecular mass can be calculated as the sum of the atomic masses of all the atoms of any one molecule. The molecular mass can also be measured directly using mass spectrometry. In mass spectrometry, the molecular mass of a small molecule (< ~200 atoms of a given element) is exact, that is, the sum of the most common isotopes of each element. Of larger molecules, it is average, or calculated using the elemental mass number or the periodic table, since there is likely to be a statistical distribution of atoms representing the isotopes throughout the molecule. The molar mass of a substance is numerically equal to the molecular mass, but expressed in mass units per mole, usually as g/mol (grams per mole). : Example: : The atomic mass of hydrogen is 1.00784 u and that of oxygen is 15.9994 u; therefore, the molecular mass of water with formula H2O is (2 × 1.00784 u) + 15.9994 u = 18.01508 u. Therefore, one mole of water weighs 18.01508 grams. However, the exact mass of hydrogen-1 (the most common isotope) is 1.00783, and the exact mass of oxygen-16 (the most common isotope) is 15.9949, so the mass of the most common single molecule of water is 18.0105 u. Molecular mass or molar mass are used in stoichiometry calculations. Since molecules are created by chemical reactions, not nuclear reactions, a molecule's molecular mass exactly equals the sum of the atomic masses of its constituent atoms. The gram-molecular weight is the molecular weight of a chemical compound expressed in grams. Thus the molecular weight of calcium carbonate is 100 therefore the gram-molecular weight is 100 g. The Gram Molecular Weight is frequently used, particularly for calculating the concentration of solutions. Thus a molar solution will contain the gram molecular weight of a chemical dissolved to produce one litre of solution.

Molecular mass in polymer Chemistry

In polymer chemistry, due to the varying chain lengths between the polymer macromolecules, various types of molecular mass are used to quantify the molar mass distribution.

See also


- Avogadro's number
- Atomic mass unit
- Dalton (unit)

External links


- [http://www.vias.org/simulations/simusoft_molform.html Learning by Simulations] Calculation of Molecular Formulas from Molecular Masses Category:Mass Category:Chemical properties ja:分子量

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, Fe2- (e.g. Fe(CO)42-,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

Sulfuric acid

Sulfuric acid (British English: sulphuric acid), H2SO4, is a strong mineral acid. It is soluble in water at all concentrations. The old name for sulfuric acid was oil of vitriol. Sulfuric acid has many applications, and is produced in larger amounts than any other chemical besides water. World production in 2001 was 165 million tonnes, with an approximate value of $8 billion. Principal uses include fertilizer manufacturing, ore processing, chemical synthesis, wastewater processing, and oil refining.

Physical properties

Forms of sulfuric acid

Although 100% sulfuric acid can be made, this loses SO3 at the boiling point to produce 98.3% acid. The 98% grade is also more stable for storage, making it the usual form for "concentrated" sulfuric acid. Other concentrations of sulfuric acid are used for different purposes. Some common concentrations are:
- 33.5%, battery acid (used in lead-acid batteries)
- 62.18%, chamber or fertilizer acid
- 77.67%, tower or Glover acid
- 98%, concentrated Different purities are also available. Technical grade H2SO4 is impure and often colored, but it is suitable for making fertiliser. Pure grades such as US Pharmacopoeia (USP) grade are used for making pharmaceuticals and dyestuffs. When high concentrations of SO3(g) are added to sulfuric acid, H2S2O7 forms. This is called fuming sulfuric acid or oleum or, less commonly, Nordhausen acid. Concentrations of oleum are either expressed in terms of % SO3 (called % oleum) or as "% H2SO4 (the amount made if H2O were added); common concentrations are 40% oleum (109% H2SO4) and 65% oleum (114.6% H2SO4). Pure H2S2O7 is in fact a solid, melting point 36 °C.

Polarity and conductivity

Anhydrous H2SO4 is a very polar liquid, with a dielectric constant of around 100. This is due to the fact that it can dissociate by protonating itself, a process known as autoprotolysis, which occurs to a high degree, more than 10 billion times the level seen in water: : 2 H2SO4 H3SO4+ + HSO4− This allows protons to be highly mobile in H2SO4. It also makes sulfuric acid an excellent solvent for many reactions. In fact, the equilibrium is more complex than shown above. 100% H2SO4 contains the following species at equilibrium (figures shown as mmol per kg solvent): HSO4− (15.0), H3SO4+ (11.3), H3O+ (8.0), HS2O7− (4.4), H2S2O7 (3.6), H2O (0.1).

Chemical properties

Reaction with water

The hydration reaction of sulfuric acid is highly exothermic. If water is added to concentrated sulfuric acid, it can boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. This can be remembered through mnemonics such as "Do as you oughta: add acid to water", "A.A.: Add Acid", or "Drop acid, not water." Note that part of this problem is due to the relative densities of the two liquids. Water is less dense than sulfuric acid and will tend to float above the acid. The reaction is best thought of as forming hydronium ions, as such: H2SO4 + H2O → H3O+ + HSO4- And then: HSO4- + H2O → H3O+ + SO42- Because the hydration of sulfuric acid is thermodynamically favorable (ΔH = 880 kJ/mol), sulfuric acid is an excellent dehydrating agent, and is used to prepare many dried fruits. The affinity of sulfuric acid for water is sufficiently strong that it will take hydrogen and oxygen atoms out of other compounds; for example, mixing starch (C6H12O6)n and concentrated sulfuric acid will give elemental carbon and water which is absorbed by the sulfuric acid (which becomes slightly diluted): (C6H12O6)n → 6C + 6H2O. The effect of this can bee seen when concentrated sulphuric acid spilled on paper; the the starch reacts to give a burned appearance, the carbon appears as soot would in a fire.

Other reactions of sulfuric acid

As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, copper(II) sulfate, the familiar blue salt of copper used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid: : CuO + H2SO4 → CuSO4 + H2O Sulfuric acid can be used to displace weaker acids from their salts, for example sodium acetate gives acetic acid: H2SO4 + CH3COONa → NaHSO4 + CH3COOH Likewise the reaction of sulfuric acid with potassium nitrate can be used to produce nitric acid, along with a precipitate of potassium bisulfate. With nitric acid itself, sulfuric acid acts as both an acid and a dehydrating agent, forming the nitronium ion NO2+, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction where protonation occurs on an oxygen atom, is important in many reactions in organic chemistry, such as Fischer esterification and dehydration of alcohols. Sulfuric acid reacts with most metals in a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese and nickel, but tin and copper require hot concentrated acid. Lead and tungsten are, however, resistant to sulfuric acid. The reaction with iron (shown) is typical for most of these metals, but the reaction with tin is unusual in that it produces sulfur dioxide rather than hydrogen. :Fe(s) + H2SO4(aq) → H2(g) + FeSO4(aq) :Sn(s) + 2 H2SO4(l) → SnSO4 + 2 H2O + SO2

Environmental aspects

Sulfuric acid is a constituent of acid rain, being formed by atmospheric oxidation of water and sulfur dioxide. Sulfur dioxide is the main product when sulfur-containing fuels such as coal or oil are burned. Sulfuric acid is a major component in the hot atmosphere of the planet Venus, and as a result exploration of Venus with spacecraft is difficult.

History of sulfuric acid

The discovery of sulfuric acid is credited to the 9th century Persian physician and alchemist Ibn Zakariya al-Razi (Rhases), who obtained the subtance by dry distillation of minerals including iron (II) sulfate heptahydrate, FeSO4 • 7H2O, called green vitriol, and copper(II) sulfate pentahydrate, CuSO4 • 5H2O, called blue vitriol. When heated, these compounds decompose to iron(II) oxide and copper(II) oxide, respectively, giving off water and sulfur trioxide, which combine to produce a dilute solution of sulfuric acid. This method was popularized in Europe through translations of Islamic treatises and books by European alchemists, such as the 13th-century German Albertus Magnus. For this reason, sulfuric acid was known to medieval European alchemists as oil of vitriol and spirit of vitriol, among other names. In the 17th century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO3), in the presence of steam. As the saltpeter decomposes, it oxidizes the sulfur to SO3, which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid. In 1746 in Birmingham, John Roebuck began producing sulfuric acid this way in lead-lined chambers, which were stronger, less expensive, and could be made larger than the glass containers which had been used previously. This lead chamber process allowed the effective industrialization of sulfuric acid production, and with several refinements remained the standard method of production for almost two centuries. John Roebuck's sulfuric acid was only about 35–40% sulfuric acid. Later refinements in the lead-chamber process by the French chemist Joseph-Louis Gay-Lussac and the British chemist John Glover improved this to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product, and throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS2) was heated in air to yield iron (II) sulfate, FeSO4, which was oxidized by further heating in air to form iron(III) sulfate, Fe2(SO4)3, which when heated to 480 °C decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. The expense of this process prevented the large-scale use of concentrated sulfuric acid. In 1831, the British vinegar merchant Peregrine Phillips patented a far more economical process for producing sulfur trioxide and concentrated sulfuric acid, now known as the contact process. Essentially all of the world's supply of sulfuric acid is now produced by this method.

Manufacture

Sulfuric acid is produced from sulfur, oxygen and water via the contact process. In the first step sulfur is burned to produce sulfur dioxide. This is oxidised to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. Finally the sulfur trioxide is treated with water (in the form of 97-98% H2SO4) to produce 98-99% sulfuric acid. Alternatively, the SO3 is absorbed into H2SO4 to produce oleum (H2S2O7), which is then diluted to form sulfuric acid. :(1) (Solid|s) + O2(Gas|g) → SO2(g) (2) 2 SO2 + O2(g) → 2 SO3(g)     (in presence of V2O5) (3) SO3(g) + H2O(l) → H2SO4(l) In 1993, US production of sulfuric acid amounted to 36.4 million tonnes. World production in 2001 was 165 million tonnes.

Uses

Sulfuric acid is a very important commodity chemical, and indeed a nation's sulfuric acid production is a good indicator of its industrial strength. The major use (60% of total worldwide) for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilisers as well as sodium triphosphate for detergents. In this method phosphate rock is used, and more than 100 million tonnes is processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as fluorosilicic acid. The overall process can be represented as :Ca5F(PO4)3 + 5 H2SO4 + 10 H2O → 5 CaSO4·2 H2O + HF + 3 H3PO4 Sulfate fertilisers such as ammonium sulfate are manufactured using sulfuric acid, although in smaller quantities than phosphates. Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as papermaker's alum. This can react with small amounts of soap on paper pulp fibres to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibres into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid: :Al2O3 + 3 H2SO4 → Al2(SO4)3 + 3 H2O Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanoneoxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H2SO4 is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also important in the manufacture of dyestuffs. A mixture of sulfuric acid and water is used as the electrolyte in various types of lead-acid battery where it undergoes a reversible reaction where lead and lead dioxide are converted to lead(II) sulfate. Sulfuric acid is also the principal ingredient in some drain cleaners, used to clear blockages consisting of paper, rags, and other materials not easily dissolved by caustic solutions.

Emergencies involving sulfuric acid

The boiling of sulfuric acid when water is added often causes clouds of sulfuric acid vapor to be released, this vapor being extremely hot as well as very acidic. Fires near or involving sulfuric acid are usually fought using foam or dry earth agents, to avoid the possibilty of the acid boiling. Where water must be used, the aim is to pour water on as much and as fast as possible to cool the resulting reaction. Fire fighters wear splash suits when dealing with sulfuric acid, to protect themselves against both the vapor and any splashes or spills.

Precautions

When mixing with water, sulfuric acid should always be added to water, never the other way round. See above for more information. As a strong acid and an oxidiser, sulfuric acid should be stored away from bases and reducing agents. It is highly corrosive even when dilute, attacking many metals such as iron and aluminium. Gloves and protective goggles should be worn when handling dilute H2SO4, and the concentrated acid also requires use of a face shield and PVC apron. When working with acids, water should be available to flush the skin and eyes, and a neutralizing agent such as sodium bicarbonate should be handy to handle spills on other surfaces. (Sodium bicarbonate solution must not be applied to acid burns on the skin, as the heat liberated by the reaction can cause further damage to tissue.) Sodium bicarbonate will react with the acid to form carbon dioxide and