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Sulfuric Acid

Sulfuric acid

Sulfuric acid (British English: sulphuric acid), H₂S O₄, 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⁺ + SO₄^2- 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 sodium sulfate. After spills are absorbed with waste material any residual acid can be neutralized with a solution of water and sodium bicarbonate. When application of fresh solution does not result in evolution of carbon dioxide the neutralization is complete. The residual materials and products can promote corrosion of ferrous metals and should be removed by further washing with water.

Comic rhyme

Sulfuric acid is one of the few compounds whose chemical formula is well known by the general public, because of a comic rhyme: Johnny was a chemist's son, but Johnny is no more. What Johnny thought was H₂O was H₂SO₄. In the U.S., a more common variant is: Little Johnny took a drink, but he shall drink no more. For what he thought was H₂O, was H₂SO₄.

References

# Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990. # N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, pp 837-845, Pergamon Press, Oxford, UK, 1984. ISBN 0080220576. # Philip J. Chenier, Survey of Industrial Chemistry, pp 45-57, John Wiley & Sons, New York, 1987. ISBN 0471010774.

External links

- [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc03/icsc0362.htm International Chemical Safety Card 0362]
- [http://www.cdc.gov/niosh/npg/npgd0577.html NIOSH Pocket Guide to Chemical Hazards]
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- [http://www.bu.edu/es/labsafety/ESMSDSs/MSSulfuricAcid.html Sulfuric Acid MSDS]
- [http://www.chemexper.com/chemicals/supplier/cas/7664-93-9.html Chemexper catalog of Chemical Suppliers]
- [http://www.sriconsulting.com/CEH/Public/Reports/781.5000/?Abstract.html Commercial overview of the sulfuric acid industry] Category:Sulfates Category:Acids ko:황산 ms:Asid Sulfurik ja:硫酸

British English

British English (BrE) is a term used to differentiate the form of the written English language in the United Kingdom from other forms of the English language. It is also used by some, particularly Americans, to describe the spoken versions of English used within England. The term is rarely heard within the United Kingdom. British people say that they speak English - but never British - and that others speak English with an accent, such as a 'South African accent'. When speaking, they will often drop the word "accent" and simply say Canadian, American, Jamaican and so on. A less ambiguous term is English English. Although British English can describe the formal written English used in the United Kingdom, the forms of spoken English used in the United Kingdom vary considerably more than in most other areas of the world where English is spoken. Dialects and accents vary not only within regions of the UK, for example in Scotland, Northern Ireland and Wales, but also within England. The written form of the language, as taught in schools, is universally Commonwealth English with a slight emphasis on a few words that might be more common in some areas than in others. For example, although the words "wee" and "small" are interchangeable, one is more likely to see "wee" written by a Scot than by a Londoner. For historical reasons dating back to the rise of London in the 9th century, the variety of language spoken in London and the East Midlands became the standard English within the Court and thus the form of language generally accepted for use in the law, government, literature and education of the British Isles. Like other forms of languages, the English used in Britain changes over time. Although British English is often used in the United States to denote the English spelling and lexicon used outside the US, the term Commonwealth English is more accurate for this purpose. The British spellings were most famously recorded in Samuel Johnson's A Dictionary of the English Language (1755). Historically, the widespread usage of English across the world is attributed to the power once held by the British Empire, and hence the most common form of English used by the British ruling class was the English used in south-east England (in the area around the capital city London, and the main English university towns of Oxford and Cambridge). This form of the language is associated with Received Pronunciation (RP), which is still regarded by many people outside the UK (especially in the United States) as "the British accent". From the second half of the 20th century to the present day, the preeminence of the English language has largely been linked to the economic, military and political dominance of the United States in world affairs, and American English is often regarded as the most prominent form of English in the world today, especially with the large amount of U.S. cultural products (such as films, books, and music) around the world, which is not matched in volume by those from other English-speaking nations. The form of English spoken and particularly written in the United Kingdom still has a major cultural influence on the English used in many Commonwealth countries, including Australia, South Africa, and India, as well as in the European Union. Although British English is taught and used in the former British colonies of Hong Kong, Singapore and Malaysia, American English is often taught in Chinese and Japanese schools, and in other schools throughout Asia.

-ise versus -ize

Words of the sort organize/organise and their derivatives can be spelt with either s or z in British English. The -ize forms are promoted by the Oxford English Dictionary. British English with -ize is sometimes known as OED spelling, and may be marked by the registered IANA language tag 'en-GB-oed'. It is the spelling used by the Encyclopaedia Britannica, by the United Nations, and by many international organizations and academic publications. The -ize forms were used by the London Times until the mid-1980s. The -ise forms are now generally used by the British government, by the European Union and mostly taught in the British school system. They are far more prevalent in common usage. Pam Peters (2004, -ize/-ise) relates that British National Corpus data indicates the ratio of popularity for -ise forms to -ize forms in Britain is 3:2.

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

Mineral acid

A mineral acid is an acid derived from inorganic minerals by chemical reaction as opposed to organic acids. Examples are
- hydrochloric acid
- nitric acid
- phosphoric acid
- sulfuric acid Category:Acids

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

Tonne

:For the units of mass, force, and other quantities in general, see Ton. A tonne (symbol t), sometimes referred to as a metric tonne, is an measurement of weight, used in SI. Though the spelling tonne predates the introduction of the SI system in 1960 it is now used as the standard spelling for the metric weight measurement in English. The similar English units are spelt ton in English.

Definition

1 t = 1000 kg

Multiples

Origin

The spelling tonne is from French. In old English the spelling was tunne. The various spellings and meanings (tonne, ton, tun) derive from late latin tunna (cask). It may be of celtic origin. A full cask about a metre high could easily weigh 1 tonne.

Conversions

One tonne is equivalent to:
- 1 megagram (exactly). Symbol Mg. This is the SI term
- 2204.62262 pounds

Explanation

The official symbol is t, but T and mT and mt (especially in the combination mmt for "million metric tons") are also used. In France and the English-speaking countries that are predominately metric, the spelling tonne is widespread. However, in Britain, the common people consider that its measure is very close to that of the long ton and often don't bother with the distinctive spelling; for example, even the Guiness Book of World Records accepts metrification without marking this by changing the spelling. For the United States, metric ton is the name for this unit used and recommended by NIST. [http://physics.nist.gov/Pubs/SP811/sec05.html#5.1.1] In the US an unqualified mention of a "ton" almost invariably refers to a short ton. Like grams and kilograms, tonnes have also given rise to a force unit of the same name: 1 tonne-force = 9.80665 kilonewtons (kN), a unit also often called simply "tonne" or "metric ton" without identifying it as a unit of force. Note that it is only the tonne as a unit of mass which is accepted for use with SI; the tonne-force or metric ton-force is not acceptable for use with SI. The ton of TNT or tonne of trinitrotoluene is a unit of energy with the tonne as a proxy term. This unit is not acceptable for use with SI. Assuming 1000 small (thermochemical) calories per gram (4.184 kJ/g) and thus a tonne of TNT is 4.184 GJ.

References

- NIST Special Publication 811, [http://physics.nist.gov/Pubs/SP811/ Guide for the Use of the International System of Units (SI)] Category:Units of mass ko:톤 ja:トン

Fertilizer

Fertilizers or fertilisers are compounds given to plants with the intention of promoting growth; they are usually applied either via the soil, for uptake by plant roots, or by foliar spraying, for uptake through leaves. Fertilizers can be organic (composed of organic matter, i.e. carbon based), or inorganic (containing simple, inorganic chemicals). They can be naturally-occurring compounds such as peat or mineral deposits, or manufactured through natural processes (such as composting) or chemical processes (such as the Haber process). Fertilizers typically provide, in varying proportions, the three major plant nutrients (nitrogen, phosphorus, and potassium), the secondary plant nutrients (calcium, sulfur, magnesium), and sometimes trace elements (or micronutrients) with a role in plant nutrition: boron, manganese, iron, zinc, copper and molybdenum.

Inorganic fertilizers

- Examples of naturally-occurring inorganic fertilizers include diatomaceous earth and limestone.
- Examples of manufactured or chemically-synthesized inorganic fertilisers include ammonium nitrate, potassium sulfate, and superphosphate, or triple super phosphate. Synthesized materials are also called artificial fertilizers, and may be described as straight, where the product predominantly contains the three primary ingredients of nitrogen (N), phosphorous (P) and potassium/potash (K), often described as NPK fertilizers. They are named or labelled according to the content of these three elements, thus a 5-10-5 fertilizer would have 10 percent phosphate in its ingredients. If nitrogen is the main element, they are often described as nitrogen fertilizers. Alternatively they may be described as compound where there is a mix of nutrients. Chemist Justus von Leibig (in the 19th century) contributed greatly to understanding the role of inorganic compounds in plant nutrition and devised the concept of Leibig's barrel to illustrate the significance of inadequate concentrations of essential nutrients. At the same time he deemphasized the role of humus. This theory was influential in the great expansion in use of artificial fertilizers in the 20th century. Nitrogen fertilizer is often synthesized using the Haber-Bosch process, which produces ammonia. This ammonia is applied directly to the soil or used to produce other compounds, notably ammonium nitrate, a dry, concentrated product. It can also be used in the Odda Process to produce compound fertilizers such as 15-15-15. The Haber-Bosch process uses about one percent of the Earth's total energy supply (primarily in the form of natural gas) in order to provide half of the nitrogen needed in agriculture.

Organic fertilizers

- Examples of naturally occurring organic fertilizers include manure and slurry, urine, peat, seaweed and guano. Green manure crops are also grown to add nutrients to the soil.
- Examples of manufactured organic fertilizers include compost, dried blood, bone meal and seaweed extracts. The decomposing crop residue from prior years is another source of fertility. Though not strictly considered "fertilizer", the distinction seems more a matter of words than reality. Although the density of nutrients in organic material is comparatively modest, they have some advantages. For one thing organic growers typically produce some or all of their fertilizer on-site, thus lowering operating costs considerably. Then there is the matter of how effective they are at promoting plant growth, chemical soil test results aside. The answers are encouraging. Implicit in modern theories of organic agriculture is the idea that the pendulum has swung the other way to some extent in thinking about plant nutrition. While admitting the obvious success of Leibig's theory, they stress that there are serious limitations to the current methods of implementing it via chemical fertilization. They re-emphasize the role of humus and other organic components of soil, which are believed to play several important roles:
- Mobilizing existing soil nutrients, so that good growth is achieved with lower nutrient densities while wasting less
- Releasing nutrients at a slower, more consistent rate, helping to avoid a boom-and-bust pattern
- Helping to retain soil moisture, reducing the stress due to temporary dryness
- Improving the structure of the soil Organics also have the advantage of avoiding certain long-term problems associated with the regular heavy use of artificial fertilizers;
- the possibility of "burning" plants with the concentrated chemicals
- the progressive decrease of real or perceived "soil health", apparent in loss of structure, reduced ability to absorb precipitation, lightening of soil color, etc.
- the necessity of reapplying artificial fertilizers regularly (and perhaps in increasing quantities) to maintain fertility
- the cost (substantial and rising in recent years) and resulting lack of independence In practice a compromise between the use of artificial and organic fertilizers is not uncommon, typically in the form of chemical use, supplemented with the application of such organics as may be readily available such as the return of crop residues or the application of manure. It is important to differentiate between what we mean by organic fertilizers and fertilizers approved for use in organic farming and organic gardening by organizations and authorities who provide organic certification services. Some approved fertilizers may be inorganic, naturally occurring chemical compounds, e.g. minerals.

Environmental effects of fertilizer use

Over-application of fertilizers, or application at a time when the ground is waterlogged or the crop is not able to use the fertilizer, can lead to run-off in groundwater. This can enrich lakes and streams in a process called eutrophication and lead to algal blooms. It is possible to over-apply organic fertilizers as well, but their nutrient content, solubility, and release rate are typically lower. The problem is endemic, however, and is primarily associated with the use of artificial fertilizers, if only due to the massive quantities involved. Their high solubilities are also a factor. Storage and application of fertilizers in particular weather or soil conditions can also cause emissions of the greenhouse gas nitrous oxide (N₂O). Ammonia gas (NH₃) may be emitted following application of manure or slurry or due to inorganic fertilizers (to a lesser extent unless ammonia itself is used directly). Besides suppling nitrogen, ammonia can increase soil acidity (lower pH, or "souring"). For these reasons, it is recommended that knowledge of the nutrient content of the soil and nutrient requirements of the crop are carefully balanced with application of nutrients in organic and inorganic fertiliser. This process is called nutrient budgeting. By doing this the farmer will avoid wasting fertiliser and also avoid the cost of avoiding or cleaning up pollution.

Application

Fertilisers can be buried around a trees roots when it is planted, placed in bore holes near tree roots, spread on to soil, sprayed by hand, or one can stick a bag of fertilizer in the branches.aerial topdressing.

Wastewater

Wastewater is any water that has been adversely affected in quality by any anthropogenic influence. It therefore includes liquid waste discharged from domestic houses, industrial, agricultural or commercial processes. It does not include rain-water uncontaminated by human activities. There is a wide range of wastewaters and an equally wide range of technologies and techniques for mitigating the impacts of wastewaters on the receiving environment.

Wastewater types

Industrial

- Organic - bio-degradable - includes abattoirs, creameries, ice-cream manufacture
- Organic - non bio-degradable or difficult to treat - for example Pharmaceutical or Pesticide manufacturing
- Inorganic - for example metal working industry
- extreme pH - acid/alkali manufacturing, metal plating
- Toxic - e.g. metal plating, cyanide production, pesticide manufaturing
- Solids and Emulsions - e.g. Paper manufacturing, food stuffs, lubricating and hydraulic oil manufacture
- agricultural drainage - direct and diffuse

surface runoff

- Highway drainage
- Storm drains
- industrial site drainage
- Black water - surface water contaminated by sewage

Domestic drainage

- Sewage
- Cesspit leakage
- Septic tank discharge
- Sewage treatment plant discharge
- Grey water also known as sullage water - water from household functions such as washing dishes, laundry or bath water

Treatment

There are numerous processes that can be used to clean up waste waters depending on the type and extent of contamination. Most wastewater is treated in industrial-scale wastewater treatment plants (WWTPs) which may include physical, chemical and biological treament processes. However, the use of septic tanks is widespread in rural areas, serving up to one quarter of the homes in the U.S. The most important aerobic treatment system is the activated sludge process, based on the maintenance and recirculation of a complex biomass composed by micro-organisms able to degrade the organic matter carried in the wastewater. Anaerobic processes are widely applied in the treatment of industrial wastewaters and biological sludge. Some wastewater may be highly treated and reused as reclaimed water. For some waste waters ecological approaches using reed bed systems such as constructed wetlands may be appropriate. Modern systems include tertiary treatment by micro filtration or synthetic membranes. After membrane filtration, the treated wastewater is indistinguishable from waters of natural origin of drinking quality.

Reuse

Treated wastewater can be reused as drinking water (Singapore), in industry (cooling towers), in artificial recharge of aquifers, in agriculture (70% of Israel's irrigated agriculture is based on highly purified wastewater) and in the rehabilitation of natural ecosystems (Florida's Everglades).

Concentration

:For the psychological concept, see attention. For the game, see Concentration (game), for the game show, see Concentration (game show). In chemistry, concentration is the measure of how much of a given substance there is mixed with another substance. This can apply to any sort of chemical mixture, but most frequently is used in relation to solutions, where it refers to the amount of solute dissolved in a solvent. To concentrate a solution, one must add more solute, or reduce the amount of solvent (for instance, by selective evaporation). By contrast, to dilute a solution, one must add more solvent, or reduce the amount of solute. There exists a concentration at which no further solute will dissolve in a solution. At this point, the solution is said to be saturated. If additional solute is added to a saturated solution, it will not dissolve. Instead, phase separation will occur, leading to either coexisting phases or a suspension. The point of saturation depends on many variables such as ambient temperature and the precise chemical nature of the solvent and solute. Concentration may be expressed both qualitatively ('informally') or quantitatively ('numerically').

Qualitative notation

Qualitatively, solutions of relatively low concentration are described using adjectives such as "dilute," or "weak," while solutions of relatively high concentration are described as "concentrated," or "strong." As a rule, the more concentrated a chromatic solution is, the more intensely coloured it is. chromatic

Quantitative notation

Quantitative notation of concentration is far more informative and useful from a scientific point of view. There are a number of different ways to quantitatively express concentration; the most common are listed below. Note: Many units of concentration require measurement of a substance's volume, which is variable depending on ambient temperature and pressure. Unless otherwise stated, all the following measurements are assumed to be at standard state temperature and pressure (that is, 25 degrees Celsius at 1 atmosphere or 101.325 kPa).

Mass percentage

Mass percentage denotes the mass of a substance in a mixture as a percentage of the mass of the entire mixture. For instance: if a bottle contains 40 grams of ethanol and 60 grams of water, then it contains 40% ethanol by mass. Commercial concentrated aqueous reagents such as acid and bases are often labeled in concentrations of weight percentage with the specific gravity also listed. In older texts and references this is sometimes referred to as weight-weight percentage (abbreviated as w/w).

Mass-volume percentage

Mass-volume percentage, (sometimes referred to as weight-volume percentage and often abbreviated as % m/v or % w/v) denotes the mass of a substance in a mixture as a percentage of the volume of the entire mixture. Mass-volume percentage is often used for solutions made from solid reagents. It is the mass of the solute in grams multiplied by one hundred divided by the volume of solution in milliliters.

Volume-volume percentage

Volume-volume percentage or % (v/v) describes the volume of the solute in mL per 100 mL of the resulting solution. This is most useful when a liquid - liquid solution is being prepared. For example, beer is about 5% ethanol by volume. This means every 100 mL beer contains 5 mL ethanol (ethyl alcohol).

Molarity

Molarity (M) denotes the number of moles of a given substance per litre of solution. For instance: 4.0 litres of liquid, containing 2.0 moles of dissolved particles, constitutes a solution of 0.5 M. Such a solution may be described as "0.5 molar." (Working with moles can be highly advantageous, as they enable measurement of the absolute number of particles in a solution, irrespective of their weight and volume. This is often more useful when performing stoichiometric calculations.). See molar solution for further information.

Molality

Molality (m) denotes the number of moles of a given substance per kilogram of solvent. For instance: 2.0 kilograms of solvent, containing 1.0 moles of dissolved particles, constitutes a molality of 0.5 mol/kg. Such a solution may be described as "0.5 molal." The advantage of molality is, it does not change with the temperature, as it deals with the mass of solvent rather than the volume of solution. Volume typically increases with increase in temperature resulting in decrease in molarity. Molality of a solution is always constant irrespective of the physical conditions like temperature and pressure.

Molinity

Molinity is a rarely-used term that denotes the number of moles of a given substance per kilogram of solution. For instance: imagine 2.0 kg of solvent, plus 1.0 mol of dissolved particles, weighs a total of 2.5 kg. The molinity of the solution is therefore 1 mol / 2.5 kg = 0.4 mol/kg. :Note: molarity and molinity are calculated using the volume of the entire solution, but molality is calculated using the mass of solvent only. :Warning: There may be confusion between above terms, which look and sound very similar; also, the abbreviations 'M' (denoting molarity) and 'm' (denoting molality) can be ambiguous. Special care should be exercised; if there is any risk of confusion, one should fully describe the measure being used.

Normality

Normality is a concept related to molarity, usually applied to acid-base solutions and reactions. For acid-base reactions, the equivalent is the mass of acid or base that can accept or donate exactly one mole of protons (H⁺ ions). Normality is also used for redox reactions. In this case the equivalent is the quantity of oxidizing or reducing agent that can accept or furnish one mole of electrons. Whereas molarity measures the number of particles per litre of solution, normality measures the number of equivalents per litre of solution. In practice, this simply means one multiplies the molarity of a solution by the valence of the ionic solute. A bit more complex for redox reactions. Note: The normality is always equal to, or greater than the molarity for acid-base reactions. However, for redox reactions the normality is typically equal to or less than the molarity.

Mole fraction

The mole fraction χ, chi (also called molar fraction) denotes the number of moles of solute as a proportion of the total number of moles in a solution. For instance: 1 mole of solute dissolved in 9 moles of solvent would have a mole fraction of 1/10 or 0.1.

Formal

The formal (F) is yet another measure of concentration similar to molarity. It is used rarely. It is calculated based on the formula weights of chemicals per litre of solution. The difference between formal and molar concentrations is that the formal concentration indicates moles of the original chemical formula in solution, without regard for the species that actually exist in solution. Molar concentration, on the other hand, is the concentration of species in solution. For example: if one dissolves sodium carbonate (Na₂CO₃) in a litre of water, the compound dissociates into the Na⁺ and CO₃^2- ions. Some of the CO₃^2- reacts with the water to form HCO₃^- and H₂CO₃. If the pH of the solution is low, there is practically no Na₂CO₃ left in the solution. So, although we have added 1 mol of Na₂CO₃ to the solution, it does not contain 1 M of that substance. (Rather, it contains a molarity based on the other constituents of the solution.) However, one can still say that the solution contains 1 F of Na₂CO₃.

"Parts-per" notation

The parts-per notation is used for extremely low concentrations. This is often used to denote the relative abundance of trace elements in the Earth's crust, trace elements in forensics or other analyses, or levels of pollutants in the environment.
- Parts per hundred (denoted by '%' and very rarely 'pph') - denotes one particle of a given substance for every 99 other particles. This is the common percent. 1 part in 10².
- Parts per thousand (denoted by '‰' [the per mil symbol], and occasionally 'ppt') denotes one particle of a given substance for every 999 other particles. This is roughly equivalent to one drop of ink in a cup of water, or one second per 17 minutes. 'Parts per thousand' is often used to record the salinity of seawater. 1 part in 10³.
- Parts per million ('ppm') denotes one particle of a given substance for every 999,999 other particles. This is roughly equivalent to one drop of ink in a 40 gallon drum of water, or one second per 280 hours. 1 part in 10⁶.
- Parts per billion ('ppb') denotes one particle of a given substance for every 999,999,999 other particles. This is roughly equivalent to one drop of ink in a canal lock full of water, or one second per 32 years. 1 part in 10⁹.
- Parts per trillion ('ppt') denotes one particle of a given substance for every 999,999,999,999 other particles. This is roughly equivalent to one drop of ink in an Olympic-sized swimming pool, or one second every 320 centuries. 1 part in 10¹².
- Parts per quadrillion ('ppq'?) denotes one particle of a given substance for every 999,999,999,999,999 other particles. This is roughly equivalent to a drop of ink in a medium-sized lake, or one second every 32,000 millennia. There are no known analytical techniques that can measure with this degree of accuracy; nevertheless, it is still used in some mathematical models of toxicology and epidemiology. 1 part in 10¹⁵. Warning: although 'ppt' is usually used to denote 'parts per trillion', it is also on occasion used to denote 'parts per thousand'. If there is any chance of ambiguity, one should describe the abbreviation in full. According to the U.S. National Institute of Standards and Technology (NIST) Guide for the Use of the International System of Units (SI), "the language-dependent terms part per million, part per billion, and part per trillion ... are not acceptable for use with the SI to express the values of quantities." [http://physics.nist.gov/Pubs/SP811/sec07.html#7.10.3] which lists examples of alternative expressions. Notes for clarity: :The indication given above is that parts per notation refers to numbers of particles (equivalent to moles), whereas in the last column of the chart below it is given by mass (grams per kilogram). Those using the notation need to state their usage to avoid confusion. :In atmospheric chemistry the parts per notation is commonly expressed with a v following, such as ppmv (or ppvm is some usages), to indicate parts per million by volume. In gases ppmv is equivalent to ppm by particles (Avogadro's law). This works fine for gases, but may have problems with cloud droplets and smoke or other atmospheric particulate matter.

Techniques used to determine concentration

- Spectrophotometry
- Chromatography
- Various titration methods

Table of concentration measures

Frequently used standards of concentration
Measurement	Notation	Generic formula	Typical units
Mass percentage	-	$\left ( \frac \right )$	%
Mass-volume percentage	-	$\left ( \frac \right )$	% though strictly %kg/L
Volume-volume percentage	-	$\left ( \frac \right )$	%
Molarity	M	$\left ( \frac \right )$	mol/L (or M)
Molinity	-	$\left ( \frac \right )$	mol/kg
Molality	m	$\left ( \frac \right )$	mol/kg (or m)
Molar fraction	χ (chi)	$\left ( \frac \right )$	(fraction)
Formal	F	$\left ( \frac \right )$	mol/L (or F)
Normality	N	$\left ( \frac \times valence\ of\ solute \right )$	N
Parts per hundred	% (or pph)	$\left ( \frac \right )$	da.g/kg
Parts per thousand	‰ (or ppt - )	$\left ( \frac \right )$	g/kg
Parts per million	ppm	$\left ( \frac \right )$	mg/kg
Parts per billion	ppb	$\left ( \frac \right )$	μg/kg
Parts per trillion	ppt -	$\left ( \frac \right )$	ng/kg
Parts per quadrillion	ppq	$\left ( \frac \right )$	pg/kg

- Although 'ppt' is usually used to denote 'parts per trillion', it is on occasion used for 'parts per thousand'. Sometimes 'ppt' is also used as an abbreviation for precipitate. Note (1) : The table above is described in terms of solvents and solutes; however the units given often also apply to other types of mixture. Note (2) : The use of billion, trillion, quadrillion above follows the short scale usage of these words. Category:Analytical chemistry ja:モル濃度

Lead-acid battery

Lead-acid batteries, invented in 1859 by French physicist Gaston Planté, are a type of galvanic cell and are the most commonly used rechargeable batteries today. They also represent the oldest design with one of the worst energy-to-weight ratios, although the power-to-weight ratio can be quite good. Also, the energy-to-volume ratio is good compared to other types of batteries. They are cheap and can supply high surge currents needed in starter motors. Every reasonably modern car uses a lead-acid battery for this purpose. They are also used in vehicles such as forklifts, in which the low energy-to-weight ratio may in fact be considered a benefit since the battery can be used as a counterweight. Lead-acid car batteries consist of six cells of 2 V nominal voltage. Each cell contains (in the charged state) electrodes of lead metal (Pb) and lead (IV) oxide (PbO₂) in an electrolyte of about 37 % w/w sulfuric acid (H₂SO₄). Modern designs have gelified electrolytes. In the discharged state both electrodes turn into lead(II) sulfate and the electrolyte turns into water. (This is why discharged lead-acid batteries can freeze.) Lead acid batteries for automotive use are not designed for deep discharge and should always be kept at maximum charge, using constant voltage at 13.8 V (for six element car batteries). Their capacity will severely suffer from deep cycling. Specially designed deep-cycle cells are much less susceptible to this problem, and are required for applications where the batteries are regularly discharged.
- Quiescent(open-circuit) voltage at full battery: 12.6 V
- Unloading-end voltage: 11.8 V
- Charge with 13.2-14.4 V
- Gassing voltage: 14.4 V
- Continuous-preservation charge with max. 13.2 V
- After full charge the terminal voltage will drop quickly to 13.2 V and then slowly to 12.6 V. The energy to weight ratio, or specific energy, is in the range of 108 kJ/kg (30 Wh/kg). The chemical reactions are (charged to discharged): Anode (oxidation):

\mbox (s) +\mbox_^ (aq) \leftrightarrow \mbox_ (s) +2e^- \quad\epsilon^o = 0.356 V

Cathode (reduction):

\mbox_ (s) +\mbox_^ (aq) +4\mbox^++2e^- \leftrightarrow \mbox_ (s) +2\mbox_2\mbox (l) \quad\epsilon^o = 1.685 V

Because of the open cells with liquid electrolyte in most cheap car batteries, overcharging with excessive charging voltages will generate oxygen and hydrogen gas, forming an extremely explosive mix. This should be avoided. Caution must also be observed because of the extremely corrosive nature of sulphuric acid.

Automotive and other applications

A chemical compound in the form of tablets can be added to each cell to reduce sulfate build up, and improve battery condition, however the effectiveness of such treatments is subject to debate. :See Car battery Wet cells designed for deep discharge are commonly used in golf carts and other battery electric vehicles, large backup power supplies for telephone and computer centers and off-grid household electric power systems. Absorbed glass mat (AGM) cells are also used in battery electric vehicles. Gel cells are used in back-up power supplies for alarm and smaller computer systems, and for electric scooters and electrified bicycles.

References

- U.S. Department of Energy, Primer On Lead-Acid Storage Batteries [http://www.eh.doe.gov/techstds/standard/hdbk1084/hdbk1084.pdf] (pdf). Category:Electric batteries Category:Automotive technologies ja:鉛蓄電池

Pharmaceutical

Pharmacology (in Greek: pharmacon (φάρμακον) is drug, and logos (λόγος) is science) is the study of how chemical substances interact with living systems. If these substances have medicinal properties, they are referred to as pharmaceuticals. The field encompasses drug composition, drug properties, interactions, toxicology, and desirable effects that can be used in therapy of diseases. Development of medication is a vital concern to medicine, but also has strong economical and political implications. To protect the consumer and prevent abuse, many governments regulate the manufacture, sale, and administration of medication. In the United States, the main regulatory body is the Food and Drug Administration through its publication of the USP. Pharmacology as a science is practiced by pharmacologists. Subdisciplines are clinical pharmacology (the medical field of medication effects on humans), neuro- and psychopharmacology (effects of medication on behavior and nervous system functioning), and theoretical pharmacology.

Scientific background

The study of medicinal chemicals requires intimate knowledge of the biological system affected. With the knowledge of cell biology and biochemistry increasing, the field of pharmacology has also changed substantially. It has become possible, through molecular analysis of receptors, to design chemicals that act on specific cellular signalling or metabolic pathways by affecting sites directly on cell-surface receptors (which modulate and mediate cellular signalling pathways controlling cellular function). A chemical has, from the pharmacological point-of-view, various properties. Pharmacokinetics describes its behaviour in the body - particularly in the blood (e.g. its half-life and volume of distribution), and pharmacodynamics relates its behaviour in the blood to its effects (desired effects or toxic side-effects). When describing the pharmacokinetic properties of a chemical, pharmacologists are often interested in ADME:
- Absorption - How is the medication absorbed (through the skin, the intestine, the oral mucosa)?
- Distribution - How does it spread through the organism?
- Metabolism - Is the medication converted chemically inside the body, and into which substances. Are these active? Could they be toxic?
- Excretion - How is the medication eliminated (through the bile, urine, breath, skin)? Medication is said to have a narrow or wide therapeutic index or therapeutic window. This describes the ratio of desired effect to toxic effect. A compound with a narrow therapeutic index (close to 1) exerts its desired effect at a dose close to its toxic dose. A compound with a wide therapeutic index (greater than 5) exerts its desired effect at a dose substantially below its toxic dose. Those with a narrow window are more difficult to dose and administer, and may require therapeutic drug monitoring (examples are warfarin, some antiepileptics, aminoglycoside antibiotics). Most anti-cancer drugs have a narrow therapeutic margin: toxic side-effects are almost always encountered at doses used to kill tumours.

Classification

Medication can be usually classified in various ways, e.g. by its chemical properties, mode of administration, or biological system affected. An elaborate and widely used classification system is the Anatomical Therapeutic Chemical Classification System.

Types of medication

For pain & consciousness (Analgesic drugs)

The main classes of painkillers are NSAIDs, opioids and various orphans such as paracetamol, tricyclic antidepressants and anticonvulsants.

For musculo-skeletal disorders

NSAIDs (including COX-2 selective inhibitors), muscle relaxant, neuromuscular drug
anticholinesterase

For the eye

- General: adrenergic neurone blocker, astringent, ocular lubricant
- Diagnostic: topical anesthetics, sympathomimetics, parasympatholytics, mydriatics, cycloplegics
- Anti-bacterial: antibiotics, topical antibiotics, sulfa drugs, aminoglycosides, fluoroquinolones
- Anti-viral:
- Anti-fungal: imidazoles, polyenes
- Anti-inflammatory: NSAIDs, corticosteroids
- Anti-allergy: mast cell inhibitors
- Anti-glaucoma: adrenergic agonists, beta-blockers, carbonic anhydrase inhibitors/hyperosmotics, cholinergics, miotics, parasympathomimetics, prostaglandin agonists/prostaglandin inhibitors. nitroglycerin

For the ear, nose and oropharynx

sympathomimetic, antihistamine, anticholinergic, NSAIDs, steroid, antiseptic, local anesthetic, antifungal, cerumenolytic

For the respiratory system

bronchodilator, NSAIDs, anti-allergic, antitussive, mucol

Sulfuric acid

Physical properties

Forms of sulfuric acid

Polarity and conductivity

Chemical properties

Reaction with water

Other reactions of sulfuric acid

Environmental aspects

History of sulfuric acid

Manufacture

Uses

Emergencies involving sulfuric acid

Precautions

Comic rhyme

References

External links

British English

-ise versus -ize

See also

Sulfur

Notable characteristics

Applications

Biological role

Environmental Impact

History

Occurrence

Compounds

Isotopes

Precautions

Spelling

See also

References

External links

Mineral acid

Vitriol

Extraction

Uses

Manufacture of sulfuric acid

Agriculture

Iron-gall ink

See also

External links

Tonne

Definition

Multiples

Origin

Conversions

Explanation

See also

References

Fertilizer

Inorganic fertilizers

Organic fertilizers

Environmental effects of fertilizer use

Application

See also

Wastewater

Wastewater types

Industrial

surface runoff

Domestic drainage

Treatment

Reuse

See also

Concentration

Qualitative notation

Quantitative notation

Mass percentage

Mass-volume percentage

Volume-volume percentage

Molarity

Molality

Molinity

Normality

Mole fraction

Formal

"Parts-per" notation

Techniques used to determine concentration

Table of concentration measures

Lead-acid battery