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Sodium carbonateFrom Wikipedia, the free encyclopediaJump to: navigation, search This section needs additional citations for verification.Please help improve this article by adding reliable references. Unsourced material may be challenged and removed. (May 2010)

Sodium carbonate

Other names[hide]Soda ashWashing sodaSoda crystals

Identifiers

CAS number497-19-8Y,5968-11-6 (monohydrate)6132-02-1 (decahydrate)

PubChem10340

ChemSpider9916

EC number207-838-8

RTECS numberVZ4050000

Properties

Molecular formulaNa2CO3

Molar mass105.9884 g/mol (anhydrous)124.00 g/mol (monohydrate)286.14 g/mol (decahydrate)

AppearanceWhite solid, hygroscopic

OdorOdorless

Density2.54 g/cm3 (anhydrous)2.25 g/cm3 (monohydrate)1.46 g/cm3 (decahydrate)

Melting point851 C (anhydrous)100 C (decomp, monohydrate)34 C (decomp, decahydrate)

Boiling point1600 C (anhydrous)

Solubility in water22 g/100 ml (20 C)

7 g/100 g (0 C)21.6 g/100 g (20 C)45 g/100 g (100 C)[1]

Solubilityinsoluble in alcohol, ethanol

Basicity (pKb)3.67

Refractive index (nD)1.495 (anhydrous)1.420 (monohydrate)

Structure

Coordinationgeometrytrigonal planar

Hazards

MSDSSafety Data Sheet External MSDS

EU Index011-005-00-2

EU classificationIrritant (Xi)

R-phrasesR36

S-phrases(S2), S22, S26

NFPA 704011

Flash pointNon-flammable

Related compounds

Other anionsSodium bicarbonate

Other cationsLithium carbonatePotassium carbonateRubidium carbonateCaesium carbonate

Related compoundsAmmonium carbonateNatronSodium percarbonate

Y(what is this?)(verify)Except where noted otherwise, data are given for materials in their standard state (at 25C, 100kPa)

Infobox references

Sodium carbonate (also known as washing soda or soda ash), Na2CO3 is a sodium salt of carbonic acid. It most commonly occurs as a crystalline heptahydrate, which readily effloresces to form a white powder, the monohydrate. Sodium carbonate is domestically well known for its everyday use as a water softener. It has a cooling alkaline taste, and can be extracted from the ashes of many plants. It is synthetically produced in large quantities from table salt in a process known as the Solvay process.The manufacture of glass is one of the most important uses of sodium carbonate. When it is combined with silica (SiO2) and calcium carbonate (CaCO3) and heated to very high temperatures, then cooled very rapidly, glass is produced. This type of glass is known as soda lime glass.Sodium carbonate is also used as a relatively strong base in various settings. For example, sodium carbonate is used as a pH regulator to maintain stable alkaline conditions necessary for the action of the majority of developing agents.[citation needed] It is a common additive in municipal pools used to neutralize the acidic effects of chlorine and raise pH.[2] In cooking, it is sometimes used in place of sodium hydroxide for lying, especially with German pretzels and lye rolls. These dishes are treated with a solution of an alkaline substance in order to change the pH of the surface of the food and thus improve browning.In taxidermy, sodium carbonate added to boiling water will remove flesh from the skull or bones of trophies to create the "European skull mount" or for educational display in biological and historical studies.In chemistry, it is often used as an electrolyte. This is because electrolytes are usually salt-based, and sodium carbonate acts as a very good conductor in the process of electrolysis. Additionally, unlike chloride ions which form chlorine gas, carbonate ions are not corrosive to the anodes. It is also used as a primary standard for acid-base titrations because it is solid and air-stable, making it easy to weigh accurately.In domestic use, it is used as a water softener during laundry. It competes with the ions magnesium and calcium in hard water and prevents them from bonding with the detergent being used. Without using washing soda, additional detergent is needed to soak up the magnesium and calcium ions. Called Washing Soda, Soda crystals or Sal Soda[3] in the detergent section of stores, it effectively removes oil, grease, and alcohol stains. Sodium carbonate is also used as a descaling agent in boilers such as found in coffee pots, espresso machines, etc.[citation needed]In dyeing with fiber-reactive dyes, sodium carbonate (often under a name such as soda ash fixative or soda ash activator) is used to ensure proper chemical bonding of the dye with the fibers, typically before dyeing (for tie dyes), mixed with the dye (for dye painting), or after dyeing (for immersion dyeing).[4][edit] Other applicationsSodium carbonate is a food additive (E500) used as an acidity regulator, anti-caking agent, raising agent and stabilizer. It is one of the components of kansui, a solution of alkaline salts used to give ramen noodles their characteristic flavor and texture.[5][6] Sodium carbonate is also used in the production of sherbet powder. The cooling and fizzing sensation results from the endothermic reaction between sodium carbonate and a weak acid, commonly citric acid, releasing carbon dioxide gas, which occurs when the sherbet is moistened by saliva.As a food additive (E500), it is used in the production of snus (Swedish style snuff) to stabilize the pH of the final product.[7] In Sweden, snus is regulated as a food product because it is put into the mouth, requiring pasteurization and only ingredients that are approved as food additives.Sodium carbonate is used by the brick industry as a wetting agent to reduce the amount of water needed to extrude the clay.[citation needed]In casting, it is referred to as "bonding agent" and is used to allow wet alginate to adhere to gelled alginate.[8]Sodium carbonate is used in toothpastes, where it acts as a foaming agent, an abrasive, and to temporarily increase mouth pH.Sodium carbonate is used to create the photo process known as reticulation.Sodium carbonate may be used for safely cleaning silver. First, aluminium foil is added to a glass or ceramic container, and covered with very hot water and some sodium carbonate. Silver items are dipped into this "bath" to clean them, making sure the silver makes contact with the aluminium foil. Finally, the silver is rinsed in water and left to dry.[9][edit] OccurrenceSodium carbonate is soluble in water, but can occur naturally in arid regions, especially in the mineral deposits (evaporites) formed when seasonal lakes evaporate. Deposits of the mineral natron, natural sodium carbonate decahydrate, have been mined from dry lake bottoms in Egypt since ancient times, when natron was used in the preparation of mummies and in the early manufacture of glass. Sodium carbonate has three known forms of hydrates: sodium carbonate decahydrate (natron), sodium carbonate heptahydrate (not known in mineral form) and sodium carbonate monohydrate (mineral thermonatrite). The anhydrous mineral form of sodium carbonate is quite rare and called natrite. Sodium carbonate also erupts from Tanzania's unique volcano Ol Doinyo Lengai [10], and probably erupted from other volcanoes in the past [11]. All three mineralogical forms of sodium carbonate, as well as trona (trisodium hydrogendicarbonate dihydrate), are also known from ultra-alkaline pegmatitic rocks, i.e. from the Kola Peninsula.[edit] Production[edit] MiningTrona, trisodium hydrogendicarbonate dihydrate (Na3HCO3CO32H2O), is mined in several areas of the United States and provides nearly all the domestic sodium carbonate. Large natural deposits found in 1938, such as the one near Green River, Wyoming, have made mining more economical than industrial production in North America.It is also mined out of certain alkaline lakes such as Lake Magadi in Kenya by using a basic dredging process and it is also self-regenerating so will never run out in its natural source.[edit] Barilla and kelpSeveral "halophyte" (salt tolerant) plant species and seaweed species can be processed to yield an impure form of sodium carbonate, and these sources predominated in Europe and elsewhere until the early 19th Century. The land plants (typically glassworts or saltworts) or the seaweed (typically Fucus species) were harvested, dried, and burned. The ashes were then "lixiviated" (washed with water) to form an alkali solution. This solution was boiled dry to create the final product, which was termed "soda ash"; this very old name refers to the archetypal plant source for soda ash, which was the small annual shrub Salsola soda ("barilla plant").The sodium carbonate concentration in soda ash varied very widely, from 2-3% for the seaweed-derived form ("kelp"), to 30% for the best barilla produced from saltwort plants in Spain. Plant and seaweed sources for soda ash, and also for the related alkali "potash", became increasingly inadequate by the end of the 18th Century, and the search for commercially-viable routes to synthesizing soda ash from salt and other chemicals intensified.[12][edit] Leblanc processMain article: Leblanc processIn 1791, the French chemist Nicolas Leblanc patented a process for producing sodium carbonate from salt, sulfuric acid, limestone, and coal. First, sea salt (sodium chloride) was boiled in sulfuric acid to yield sodium sulfate and hydrogen chloride gas, according to the chemical equation2 NaCl + H2SO4 Na2SO4 + 2 HClNext, the sodium sulfate was blended with crushed limestone (calcium carbonate) and coal, and the mixture was burnt, producing calcium sulfide.Na2SO4 + CaCO3 + 2 C Na2CO3 + 2 CO2 + CaSThe sodium carbonate was extracted from the ashes with water, and then collected by allowing the water to evaporate.The hydrochloric acid produced by the Leblanc process was a major source of air pollution, and the calcium sulfide byproduct also presented waste disposal issues. However, it remained the major production method for sodium carbonate until the late 1880s.[12][13][edit] Solvay process

Main article: Solvay processIn 1861, the Belgian industrial chemist Ernest Solvay developed a method to convert sodium chloride to sodium carbonate using ammonia. The Solvay process centered around a large hollow tower. At the bottom, calcium carbonate (limestone) was heated to release carbon dioxide:CaCO3 CaO + CO2At the top, a concentrated solution of sodium chloride and ammonia entered the tower. As the carbon dioxide bubbled up through it, sodium bicarbonate precipitated:NaCl + NH3 + CO2 + H2O NaHCO3 + NH4ClThe sodium bicarbonate was then converted to sodium carbonate by heating it, releasing water and carbon dioxide:2 NaHCO3 Na2CO3 + H2O + CO2Meanwhile, the ammonia was regenerated from the ammonium chloride byproduct by treating it with the lime (calcium hydroxide) left over from carbon dioxide generation:CaO + H2O Ca(OH)2Ca(OH)2 + 2 NH4Cl CaCl2 + 2 NH3 + 2 H2OBecause the Solvay process recycles its ammonia, it consumes only brine and limestone, and has calcium chloride as its only waste product. This made it substantially more economical than the Leblanc process, and it soon came to dominate world sodium carbonate production. By 1900, 90% of sodium carbonate was produced by the Solvay process, and the last Leblanc process plant closed in the early 1920s.[edit] Hou's processDeveloped by Chinese chemist Hou Debang in 1930s, the first few steps are the same as the Solvay process. However, instead of treating the remaining solution with lime, carbon dioxide and ammonia are pumped into the solution, then sodium chloride is added until the solution saturates at 40 C. Next, the solution is cooled to 10 C. Ammonium chloride precipitates and is removed by filtration, and the solution is recycled to produce more sodium carbonate. Hou's process eliminates the production of calcium chloride and the byproduct ammonium chloride can be refined or used as a fertilizer.Natrium bikarbonat adalah senyawa kimia dengan rumus NaHCO3. Dalam penyebutannya kerap disingkat menjadi bicnat. Senyawa ini termasuk kelompok garam dan telah digunakan sejak lama.Senyawa ini disebut juga baking soda (soda kue), Sodium bikarbonat, natrium hidrogen karbonat, dan lain-lain. Senyawa ini merupakan kristal yang sering terdapat dalam bentuk serbuk. Natrium bikarbonat larut dalam air. Senyawa ini digunakan dalam roti atau kue karena bereaksi dengan bahan lain membentuk gas karbon dioksida, yang menyebabkan roti "mengembang".Senyawa ini juga digunakan sebagai obat antasid (penyakit maag atau tukak lambung). Karena bersifat alkaloid (basa), senyawa ini juga digunakan sebagai obat penetral asam bagi penderita asidosis tubulus renalis (ATR) atau rhenal tubular acidosis (RTA).NaHCO3 umumnya diproduksi melalui proses Solvay, yang memerlukan reaksi natrium klorida, amonia, dan karbon dioksida dalam air. NaHCO3 diproduksi sebanyak 100 000 ton/tahun (2001).[1]Soda kue juga diproduksi secara komesial dari soda abu (diperoleh melalui penambangan bijih trona, yang dilarutkan dalam air lalu direaksikan dengan karbon dioksida. Lalu NaHCO3 mengendap sesuai persamaan berikutNa2CO3 + CO2 + H2O 2 NaHCO3

Adapun cara-cara yang digunakan untuk menghasilkan soda Ash adalah dengan cara seperti flowsheet diatas. Dengan penggunaan dari Kalsium Karbonat maka akan dihasilkan soda ash yang dapat digunakan untuk menetralisis air yang telah tercemar ataupun untuk penggunaan yang lain.

http://en.wikipedia.org/wiki/Sodium_carbonatePotassium sulfateFrom Wikipedia, the free encyclopedia(Redirected from Potassium sulphate)Jump to: navigation, search This article relies largely or entirely upon a single source. Please help improve this article by introducing appropriate citations to additional sources. (April 2010)

Potassium sulfate

Other names[hide]Potassium sulphate

Identifiers

CAS number7778-80-5Y

PubChem24507

RTECS numberTT5900000

Properties

Molecular formulaK2SO4

Molar mass174.259 g/mol

AppearanceWhite crystalline solid

Density2.66 g/cm3 [1]

Melting point1069C, 1342K, 1956F

Boiling point1689C, 1962K, 3072F

Solubility in water111 g/L (20 C)120 g/L (25 C)240 g/L (100 C)

Solubilityslightly soluble in glycerolinsoluble in acetone, alcohol, CS2

Structure

Crystal structureorthorhombic

Hazards

MSDSExternal MSDS

EU IndexNot listed

R-phrasesR22

S-phrasesS36


LD506600 mg/kg

Related compounds

Other anionsPotassium selenatePotassium tellurate

Other cationsLithium sulfateSodium sulfateRubidium sulfateCaesium sulfate

Related compoundsPotassium hydrogen sulfatePotassium sulfitePotassium bisulfitePotassium persulfate


Infobox references

Potassium sulfate (K2SO4) (in British English potassium sulphate, also called sulphate of potash, arcanite, or archaically known as potash of sulfur) is a non-flammable white crystalline salt which is soluble in water. The chemical is commonly used in fertilizers, providing both potassium and sulfur.Contents[hide] 1 History 2 Natural resources 3 Manufacture 4 Properties 5 Uses 6 Potassium hydrogen sulfate 7 References 8 See also

[edit] HistoryPotassium sulfate (K2SO4) has been known since early in the 14th century, and it was studied by Glauber, Boyle and Tachenius. In the 17th century, it was named arcanuni or sal duplicatum, as it was a combination of an acid salt with an alkaline salt. It was also known as vitriolic tartar and Glaser's salt or sal polychrestum Glaseri after the pharmaceutical chemist Christopher Glaser who prepared it and used medicinally.[2][3][edit] Natural resourcesThe mineral form of potassium sulfate, namely arcanite, is relatively rare. Natural resources of potassium sulfate are minerals abundant in the Stassfurt salt. These are cocrystalisations of potassium sulfate and sulfates of magnesium calcium and sodium. The minerals are Kainite, MgSO4KClH2O Schnite, K2SO4MgSO46H2O Leonite, K2SO4MgSO44H2O Langbeinite, K2SO42MgSO4 Glaserite, K3Na(SO4)2 Polyhalite, K2SO4MgSO42CaSO42H2OFrom some of the minerals like kainite, the potassium sulfate can be separated, because the corresponding salt is less soluble in water.With potassium chloride kieserite MgSO42H2O can be transformed and then the potassium sulfate can be dissolved in water.[edit] ManufactureThe process for manufacturing potassium sulfate is similar to that used for the manufacture of sodium sulfate.Potassium sulfate can be synthesised by reaction of potassium chloride with sulfuric acid according to the Leblanc process. Potassium sulfate is produced according to the following reaction:2 KCl + H2SO4 2 HCl + K2SO4The Hargreaves process uses sulfur dioxide, oxygen and water and potassium chloride as the starting materials to produce potassium sulfate. Hydrochloric acid evaporates off. SO2 is produced through the burning of sulfur.[edit] PropertiesThe anhydrous crystals form a double six-sided pyramid, but are in fact classified as rhombic. They are transparent, very hard and have a bitter, salty taste. The salt is soluble in water, but insoluble in solutions of potassium hydroxide (sp. gr. 1.35), or in absolute ethanol. It melts at 1078 C.[edit] UsesThe principal use of potassium sulfate is as a fertilizer. The crude salt is also used occasionally in the manufacture of glass.[edit] Potassium hydrogen sulfatePotassium hydrogen sulfate or bisulfate, KHSO4, is readily produced by mixing K2SO4 with an equivalent number of moles of sulfuric acid. It forms rhombic pyramids, which melt at 197 C. It dissolves in three parts of water at 0 C. The solution behaves much as if its two congeners, K2SO4 and H2SO4, were present side by side of each other uncombined; an excess of ethanol the precipitates normal sulfate (with little bisulfate) with excess acid remaining.The behavior of the fused dry salt is similar when heated to several hundred degrees; it acts on silicates, titanates, etc., the same way as sulfuric acid that is heated beyond its natural boiling point does. Hence it is frequently used in analytical chemistry as a disintegrating agent. For information about other salts that contain sulfate, see sulfate.http://en.wikipedia.org/wiki/Potassium_sulphateTeknologi Proses Produksi Pupuk ZK (Bagian 2)Wednesday, August 5th, 2009

Panjing Hengxing Chemicals Co., Ltd. Salah satu produsen pupuk ZK yang berdomisili di CinaFurnace MainnheimProses ini menggunakan furnace Mannheim yang berupa bejana silindris yang memiliki 2 ruang bakar, yaitu combustion chamber dan reaction chamber. Temperatur operasi furnace Mannheim adalah sebesar 800C. Karakteristik dari proses ini yaitu:1. Temperatur tinggi2. Banyak problem pada material (tingkat korosi, dll)3. Diperoleh by-product HClReaksi yang terjadi adalah:KCl + H2SO4 -> KHSO4 + HClKCl + KHSO4 -> K2SO4 + HClReaksi tahap pertama bersifat eksotermis dan terjadi pada temperatur yang rendah, sedangkan reaksi tahap kedua bersifat endotermis dan berlangsung pada temperatur 550 600C. Produk ZK selanjutnya didinginkan di cooling drum. Residu H2SO4 dinetralkan dengan penambahan Ca(OH)2 dan CaCO3 sedangkan by-product HCl yang terbentuk didinginkan di graphite heat exchanger dan selanjutnya dilakukan absorbsi 2 tahap dengan air.

Diagram alir proses MannheimSpesifikasi produk yang dihasilkan adalah sebagai berikut:

Emisi yang dihasilkan dikontrol dengan batasan HCl maksimum 5 ppm dan SO2 maksimum 800 ppm. Beberapa negara di dunia yang telah mendirikan pabrik ZK dengan proses Mannheim antara lain Belgia, Amerika Serikat, Indonesia, dan Cina.

Sodium chromateFrom Wikipedia, the free encyclopediaJump to: navigation, search Sodium chromate

IUPAC name[hide]Sodium chromate

Other names[hide]Chromic acid, (Na2CrO4), disodium saltChromium disodium oxideRachromate

Identifiers


PubChem24488

EC number231-889-5

UN number3288

RTECS numberGB2955000

Properties

Molecular formulaNa2CrO4


Appearanceyellow crystals

Density2.698 g/cm3

Melting point762 C

Solubility in water53 g/100 ml (20 C)

Structure

Crystal structureorthorhombic (hexagonal above 413 C)

Thermochemistry

Std enthalpy offormation fHo2981329 kJ/mol

Hazards

MSDSICSC 1370

EU Index024-018-00-3

EU classificationCarc. Cat. 2Muta. Cat. 2Repr. Cat. 2Very toxic (T+)Harmful (Xn)Corrosive (C)Dangerous for the environment (N)

R-phrasesR45, R46, R60, R61, R21, R25, R26, R34, R42/43, R48/23, R50/53

S-phrasesS53, S45, S60, S61

NFPA 704030OX


Related compounds

Other anionsSodium dichromateSodium molybdateSodium tungstate

Other cationsPotassium chromateCalcium chromateBarium chromate


Infobox references

Sodium chromate is a yellow solid chemical compound used as a corrosion inhibitor in the petroleum industry,[1] a dyeing auxiliary in the textile industry,[1] as a wood preservative[2]. and as a diagnostic pharmaceutical in determining red blood cell volume[3].It is obtained from the reaction of sodium dichromate with sodium hydroxide. It is hygroscopic and can form tetra-, hexa-, and decahydrates. Sodium chromate, like other hexavalent chromium compounds, can be carcinogenic.[1]The substance is a strong oxidant. It is soluble in water[4], producing a weakly basic solution[5].http://en.wikipedia.org/wiki/Sodium_chromate

Phosphoric acidFrom Wikipedia, the free encyclopediaJump to: navigation, search This article is about orthophosphoric acid. For other acids commonly called "phosphoric acid", see Phosphoric acids and phosphates.Phosphoric acid

IUPAC name[hide]trihydroxidooxidophosphorusphosphoric acid

Other names[hide]Orthophosphoric acid

Identifiers

CAS number7664-38-2Y,16271-20-8 (hemihydrate)

ChemSpider979

EC number231-633-2

UN number1805

RTECS numberTB6300000

Properties

Molecular formulaH3PO4


Appearancewhite solid or colourless, viscous liquid (>42 C)

Density1.885 g/mL (liquid)1.685 g/mL (85% solution)2.030 g/mL (crystal at 25 C)

Melting point42.35 C (anhydrous)29.32 C (hemihydrate)

Boiling point158 C (decomp)

Solubility in water5.48 g/mL

Acidity (pKa)2.148, 7.198, 12.375

Viscosity2.49.4 cP (85% aq. soln.)147 cP (100%)

Hazards

MSDSICSC 1008

EU Index015-011-00-6

EU classificationCorrosive (C)

R-phrasesR34

S-phrases(S1/2) S26 S45

NFPA 704020COR


Related compounds

Related phosphorus oxoacidsHypophosphorous acidPhosphorous acidPyrophosphoric acidTriphosphoric acidPerphosphoric acidPermonophosphoric acid


Infobox references

Phosphoric acid, also known as orthophosphoric acid or phosphoric(V) acid, is a mineral (inorganic) acid having the chemical formula H3PO4. Orthophosphoric acid molecules can combine with themselves to form a variety of compounds which are also referred to as phosphoric acids, but in a more general way. The term phosphoric acid can also refer to a chemical or reagent consisting of phosphoric acids, usually orthophosphoric acid.Contents[hide] 1 Orthophosphoric acid chemistry 1.1 pH and composition of a phosphoric acid aqueous solution 2 Chemical reagent 2.1 Preparation of hydrogen halides 2.2 Rust removal 2.3 Processed food use 2.3.1 Biological effects on bone calcium and kidney health 2.4 Medical use 3 Preparation 3.1 Thermal phosphoric acid 3.2 Wet phosphoric acid 3.3 Kiln Phosphoric Acid 4 Other applications 5 See also 6 References 7 External links

[edit] Orthophosphoric acid chemistryPure anhydrous phosphoric acid is a white solid that melts at 42.35 C to form a colorless, viscous liquid.Most people and even chemists refer to orthophosphoric acid as phosphoric acid, which is the IUPAC name for this compound. The prefix ortho is used to distinguish the acid from other phosphoric acids, called polyphosphoric acids two(ii) . Orthophosphoric acid is a non-toxic, inorganic, rather weak triprotic acid, which, when pure, is a solid at room temperature and pressure. The chemical structure of orthophosphoric acid is shown above in the data table. Orthophosphoric acid is a very polar molecule; therefore it is highly soluble in water. The oxidation state of phosphorus (P) in ortho- and other phosphoric acids is +5; the oxidation state of all the oxygen atoms (O) is 2 and all the hydrogen atoms (H) is +1. Triprotic means that an orthophosphoric acid molecule can dissociate up to three times, giving up an H+ each time, which typically combines with a water molecule, H2O, as shown in these reactions:H3PO4(s) + H2O(l) H3O+(aq) + H2PO4(aq) Ka1= 7.25103H2PO4(aq)+ H2O(l) H3O+(aq) + HPO42(aq) Ka2= 6.31108HPO42(aq)+ H2O(l) H3O+(aq) + PO43(aq) Ka3= 3.981013The anion after the first dissociation, H2PO4, is the dihydrogen phosphate anion. The anion after the second dissociation, HPO42, is the hydrogen phosphate anion. The anion after the third dissociation, PO43, is the phosphate or orthophosphate anion. For each of the dissociation reactions shown above, there is a separate acid dissociation constant, called Ka1, Ka2, and Ka3 given at 25 C. Associated with these three dissociation constants are corresponding pKa1=2.12 , pKa2=7.21 , and pKa3=12.67 values at 25 C. Even though all three hydrogen (H ) atoms are equivalent on an orthophosphoric acid molecule, the successive Ka values differ since it is energetically less favorable to lose another H+ if one (or more) has already been lost and the molecule/ion is more negatively-charged.Because the triprotic dissociation of orthophosphoric acid, the fact that its conjugate bases (the phosphates mentioned above) cover a wide pH range, and, because phosphoric acid/phosphate solutions are, in general, non-toxic, mixtures of these types of phosphates are often used as buffering agents or to make buffer solutions, where the desired pH depends on the proportions of the phosphates in the mixtures. Similarly, the non-toxic, anion salts of triprotic organic citric acid are also often used to make buffers. Phosphates are found pervasively in biology, especially in the compounds derived from phosphorylated sugars, such as DNA, RNA, and adenosine triphosphate (ATP). There is a separate article on phosphate as an anion or its salts.Upon heating orthophosphoric acid, condensation of the phosphoric units can be induced by driving off the water formed from condensation. When one molecule of water has been removed for each two molecules of phosphoric acid, the result is pyrophosphoric acid (H4P2O7). When an average of one molecule of water per phosphoric unit has been driven off, the resulting substance is a glassy solid having an empirical formula of HPO3 and is called metaphosphoric acid.[1] Metaphosphoric acid is a singly anhydrous version of orthophosphoic acid and is sometimes used as a water- or moisture-absorbing reagent. Further dehydrating is very difficult, and can be accomplished only by means of an extremely strong desiccant (and not by heating alone). It produces phosphoric anhydride, which has an empirical formula P2O5, although an actual molecule has a chemical formula of P4O10. Phosphoric anhydride is a solid, which is very strongly moisture-absorbing and is used as a desiccant.[edit] pH and composition of a phosphoric acid aqueous solutionFor a given total acid concentration [A] = [H3PO4] + [H2PO4] + [HPO42] + [PO43] ([A] is the total number of moles of pure H3PO4 which have been used to prepare 1 liter of solution), the composition of an aqueous solution of phosphoric acid can be calculated using the equilibrium equations associated with the three reactions described above together with the [H+][OH] = 1014 relation and the electrical neutrality equation. Possible concentrations of polyphosphoric molecules and ions is neglected. The system may be reduced to a fifth degree equation for [H+] which can be solved numerically, yielding:[A] (mol/L)pH[H3PO4]/[A] (%)[H2PO4]/[A] (%)[HPO42]/[A] (%)[PO43]/[A] (%)

11.0891.78.296.201061.601017

1011.6276.123.96.201055.551016

1022.2543.156.96.201042.331014

1033.0510.689.36.201031.481012

1044.011.3098.66.191021.341010

1055.000.13399.30.6121.30108

1065.971.3410294.55.501.11106

1076.741.8010374.525.53.02105

10107.008.2410461.738.38.18105

For large acid concentrations, the solution is mainly composed of H3PO4. For [A] = 102, the pH is closed to pKa1, giving an equimolar mixture of H3PO4 and H2PO4. For [A] below 103, the solution is mainly composed of H2PO4 with [HPO42] becoming non negligible for very dilute solutions. [PO43] is always negligible.[edit] Chemical reagentPure 7585% aqueous solutions (the most common) are clear, colourless, odourless, non-volatile, rather viscous, syrupy liquids, but still pourable. Phosphoric acid is very commonly used as an aqueous solution of 85% phosphoric acid or H3PO4. Because it is a concentrated acid, an 85% solution can be corrosive, although nontoxic when diluted. Because of the high percentage of phosphoric acid in this reagent, at least some of the orthophosphoric acid is condensed into polyphosphoric acids in a temperature-dependent equilibrium, but, for the sake of labeling and simplicity, the 85% represents H3PO4 as if it were all orthophosphoric acid. Other percentages are possible too, even above 100%, where the phosphoric acids and water would be in an unspecified equilibrium, but the overall elemental mole content would be considered specified. When aqueous solutions of phosphoric acid and/or phosphate are dilute, they are in or will reach an equilibrium after a while where practically all the phosphoric/phosphate units are in the ortho- form.[edit] Preparation of hydrogen halidesPhosphoric acid reacts with halides to form the corresponding hydrogen halide gas (steamy fumes are observed on warming the reaction mixture). This is a common practice for the laboratory preparation of hydrogen halides.NaCl(s)+ H3PO4(l) NaH2PO4(s)+ HCl(g)

NaBr(s)+ H3PO4(l) NaH2PO4(s)+ HBr(g)

NaI(s)+ H3PO4(l) NaH2PO4(s)+ HI(g)

[edit] Rust removalPhosphoric acid may be used as a "rust converter", by direct application to rusted iron, steel tools, or surfaces. The phosphoric acid converts reddish-brown iron(III) oxide (rust) to black ferric phosphate, FePO4."Rust converter" is sometimes a greenish liquid suitable for dipping (in the same sort of acid bath as is used for pickling metal), but it is more often formulated as a gel, commonly called naval jelly. It is sometimes sold under other names, such as "rust remover" or "rust killer". As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces.After treatment, the black ferric-phosphate coating can be scrubbed off, leaving a fresh metal surface. Multiple applications of phosphoric acid may be required to remove all rust. The black phosphate coating can also be left in place, where it will provide moderate further corrosion resistance. (Such protection is also provided by the superficially similar Parkerizing and blued electrochemical conversion coating processes.)[edit] Processed food useFood-grade phosphoric acid (additive E338) is used to acidify foods and beverages such as various colas, but not without controversy regarding its health effects. It provides a tangy or sour taste and, being a mass-produced chemical, is available cheaply and in large quantities. The low cost and bulk availability is unlike more expensive seasonings that give comparable flavors, such as citric acid which is obtainable from lemons and limes. (However most citric acid in the food industry is not extracted from citrus fruit, but fermented by Aspergillus niger mold from scrap molasses, waste starch hydrolysates and phosphoric acid.)[edit] Biological effects on bone calcium and kidney healthPhosphoric acid, used in many soft drinks (primarily cola), has been linked to lower bone density in epidemiological studies. For example, a study[2] using dual-energy X-ray absorptiometry rather than a questionnaire about breakage, provides reasonable evidence to support the theory that drinking cola results in lower bone density. This study was published in the American Journal of Clinical Nutrition. A total of 1672 women and 1148 men were studied between 1996 and 2001. Dietary information was collected using a food frequency questionnaire that had specific questions about the number of servings of cola and other carbonated beverages and that also made a differentiation between regular, caffeine-free, and diet drinks. The paper cites significant statistical evidence to show that women who consume cola daily have lower bone density. Total phosphorus intake was not significantly higher in daily cola consumers than in nonconsumers; however, the calcium-to-phosphorus ratios were lower. The study also suggests that further research is needed to confirm the findings.On the other hand, a study funded by Pepsi suggests that insufficient intake of phosphorus leads to lower bone density. The study does not examine the effect of phosphoric acid, which binds with magnesium and calcium in the digestive tract to form salts that are not absorbed, but rather studies general phosphorus intake.[3]However, a well-controlled clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion.[4] The study compared the impact of water, milk, and various soft drinks (two with caffeine and two without; two with phosphoric acid and two with citric acid) on the calcium balance of 20- to 40-year-old women who customarily consumed ~3 or more cups (680 mL) of a carbonated soft drink per day. They found that, relative to water, only milk and the two caffeine-containing soft drinks increased urinary calcium, and that the calcium loss associated with the caffeinated soft drink consumption was about equal to that previously found for caffeine alone. Phosphoric acid without caffeine had no impact on urine calcium, nor did it augment the urinary calcium loss related to caffeine. Because studies have shown that the effect of caffeine is compensated for by reduced calcium losses later in the day,[5] Heaney and Rafferty concluded that the net effect of carbonated beveragesincluding those with caffeine and phosphoric acidis negligible, and that the skeletal effects of carbonated soft drink consumption are likely due primarily to milk displacement.Other chemicals such as caffeine (also a significant component of popular common cola drinks) were also suspected as possible contributors to low bone density, due to the known effect of caffeine on calciuria. One other study, involving 30 women over the course of a week, suggests that phosphoric acid in colas has no such effect, and postulates that caffeine has only a temporary effect, which is later reversed. The authors of this study conclude that the skeletal effects of carbonated beverage consumption are likely due primarily to milk displacement.[4] (Another possible confounding factor may be an association between high soft drink consumption and sedentary lifestyle.)Cola consumption has also been associated with chronic kidney disease and kidney stones through medical research.[6] The preliminary results suggest that cola consumption may increase the risk of chronic kidney disease.[edit] Medical usePhosphoric acid is used in dentistry and orthodontics as an etching solution, to clean and roughen the surfaces of teeth where dental appliances or fillings will be placed. Phosphoric acid is also an ingredient in over-the-counter anti-nausea medications that also contain high levels of sugar (glucose and fructose). This acid is also used in many teeth whiteners to eliminate plaque that may be on the teeth before application.[edit] PreparationPhosphoric acid can be prepared by three routes the thermal process, the wet process and the dry kiln process.[edit] Thermal phosphoric acidThis very pure phosphoric acid is obtained by burning elemental phosphorus to produce phosphorus pentoxide and dissolving the product in dilute phosphoric acid. This produces a very pure phosphoric acid, since most impurities present in the rock have been removed when extracting phosphorus from the rock in a furnace. The end result is food-grade, thermal phosphoric acid; however, for critical applications, additional processing to remove arsenic compounds may be needed.[edit] Wet phosphoric acidSee also: Nitrophosphate processWet process phosphoric acid is prepared by adding sulfuric acid to tricalcium phosphate rock, typically found in nature as apatite.The reaction is:Ca5(PO4)3X + 5 H2SO4 + 10 H2O 3 H3PO4 + 5 CaSO42H2O + HXwhere X may include OH, F, Cl, and BrThe initial phosphoric acid solution may contain 2333% P2O5, but can be concentrated by the evaporation of water to produce commercial- or merchant-grade phosphoric acid, which contains about 54% P2O5. Further evaporation of water yields superphosphoric acid with a P2O5 concentration above 70%.[7][8]Digestion of the phosphate ore using sulfuric acid yields the insoluble calcium sulfate (gypsum), which is filtered and removed as phosphogypsum. Wet-process acid can be further purified by removing fluorine to produce animal-grade phosphoric acid, or by solvent extraction and arsenic removal to produce food-grade phosphoric acid.[edit] Kiln Phosphoric AcidKiln phosphoric acid (KPA) process technology is the most recent technology. Called the Improved Hard Process,[9] this technology will both make low grade phosphate rock reserves commercially viable and will increase the P2O5 recovery from existing phosphate reserves. This may significantly extend the commercial viability of phosphate reserves.[edit] Other applications Phosphoric acid is used as the electrolyte in phosphoric-acid fuel cells. It is also used as an external standard for phosphorus-31 Nuclear magnetic resonance (NMR). Phosphoric acid is used as a cleaner by construction trades to remove mineral deposits, cementitious smears, and hard water stains. It is also used as a chelant in some household cleaners aimed at similar cleaning tasks. Hot phosphoric acid is used in microfabrication to etch silicon nitride (Si3N4). It is highly selective in etching Si3N4 instead of SiO2, silicon dioxide.[10] Phosphoric acid is used as a flux by hobbyists (such as model railroaders) as an aid to soldering. Phosphoric acid is also used in hydroponics pH solutions to lower the pH of nutrient solutions. While other types of acids can be used, phosphorus is a nutrient used by plants, especially during flowering, making phosphoric acid particularly desirable. General Hydroponics pH Down liquid solution contains phosphoric acid in addition to citric acid and ammonium bisulfate with buffers to maintain a stable pH in the nutrient reservoir. Phosphoric acid is used as an electrolyte in copper electropolishing for burr removal and circuit board planarization. Phosphoric acid is used with distilled water (23 drops per gallon) as an electrolyte in oxyhydrogen (HHO) generators. Phosphoric acid is used as a pH adjuster in cosmetics and skin-care products.[11] Phosphoric acid is used as a chemical oxidizing agent for activated carbon production, as used in the Wentworth Process.[12] Phosphoric acid is also used for high-performance liquid chromatography. Phosphoric acid can be used as a dispersing agent in detergents and leather treatment. Phosphoric acid can be used for deliming hides in leather tanning. Phosphoric acid can be used as an additive to stabilize acidic aqueous solutions within a wanted and specified pH range Phosphoric acid is the key ingredient that gives the bite taste in widely consumed Coca-Cola and Pepsi sodas.[edit] See also Phosphate fertilizers, such as ammonium phosphate fertilizers In compound semiconductor processing, phosphoric acid is a common wet etching agent: for example, in combination with hydrogen peroxide and water it is used to etch InGaAs selective to InP.[13]http://en.wikipedia.org/wiki/Phosphoric_acidDescriptionArafura Resources (ASX:ARU) is an emerging rare earths producer that listed on the Australian Securities Exchange in 2003.Arafura is currently developing its Nolans rare earths-phosphate-uranium project in Australias Northern Territory. The project is underpinned by a 30 million tonne resource that can sustain a 20-year mine life. Arafura has developed a processing flow sheet, demonstrated the recovery of rare earths, phosphoric acid and uranium at a pre-production scale pilot plant, and is making rapid progress in completing a bankable feasibility study. The Company is on track for first production from Nolans in 2013.Arafura has an exploration and development program aimed at enhancing its position in the rare earths market.

BeneficiationThe beneficiation process, comprising crushing, screening, gravity separation and flotation to remove impurities, will generate about 700,000 tonnes of mineral concentrate each year. This mineral concentrate will go through a grinding process, then be dried ready for transport to Arafuras Rare Earths Complex at Whyalla in South Australia.This treatment process removes about 300,000 tonnes a year as waste material (tailings), which will remain at the mine permanently in specially designed tailings storage facilities.All waste materials, including overburden and tailings, will be stored according to accepted engineering and environmental principles. This will ensure that erosion is minimised and natural water courses remain free flowing and are not adversely impacted.Some mine overburden rock may be used as road base or to construct the tailings storage facilities. Wherever possible, water will be recycled to the beneficiation plant and also used for dust suppression at the mine.Phosphoric AcidThe Nolans project is delivering products for the future as well as materials for well established, active and transparent markets.

With the capability to produce phosphoric acid and calcium chloride, Nolans provides multiple market opportunities.

Phosphoric acid is a fundamental building block in the production of phosphate fertilisers, detergents, food stuffs andseveral other chemicals that contribute to todays improving standard of living, particularly in emerging economies like Brazil, Russia, India and China.

With the potential high growth of ethanol as an alternative fuel, much of which is extracted from corn, phosphatefertilisers will be essential to this industry.

Arafura Resources will, in the next few years, become a long-term supplier of rare earth oxides, phosphoric acid, uranium oxide and gypsum from its 100%-owned Nolans Project.

Sodium chlorideFrom Wikipedia, the free encyclopediaJump to: navigation, search "NaCl" redirects here. For the Google technology, see Google Native Client.This article is about the chemical compound. For sodium chloride in the diet, see Salt. For sodium chloride as a mineral, see Halite.Sodium chloride

IUPAC name[hide]Sodium chloride

Other names[hide]Common salt, halite, table salt, rock salt, saline, hyposaline, sodium monochloride, sodium chloric, saltex[1]

Identifiers


PubChem5234

ChemSpider5044

RTECS numberVZ4725000

Properties

Molecular formulaNaCl


AppearanceColorless/white crystalline solid

OdorOdorless

Density2.165 g/cm3



Solubility in water356 g/L (0 C)359 g/L (25 C)391 g/L (100 C)

Solubilitysoluble in glycerol, ethylene glycol, formic acidinsoluble in HCl

Solubility in methanol14.9 g/L

Solubility in ammonia21.5 g/L

Acidity (pKa)6.77.3

Refractive index (nD)1.5442 (589 nm)

Structure

Crystal structureFace-centered cubic(see text), cF8

Space groupFm3m, No. 225

Lattice constanta=564.02 pm

CoordinationgeometryOctahedral (Na+)Octahedral (Cl)

Hazards

MSDSExternal MSDS

EU IndexNot listed

NFPA 704000


LD5030008000 mg/kg (oral in rats, mice, rabbits)[2]

Related compounds

Other anionsSodium fluorideSodium bromideSodium iodide

Other cationsLithium chloridePotassium chlorideRubidium chlorideCaesium chloride

Supplementary data page

Structure andpropertiesn, r, etc.

ThermodynamicdataPhase behaviourSolid, liquid, gas

Spectral dataUV, IR, NMR, MS


Infobox references

Sodium chloride, also known as salt, common salt, table salt, or halite, is an ionic compound with the formula NaCl. Sodium chloride is the salt most responsible for the salinity of the ocean and of the extracellular fluid of many multicellular organisms. As the major ingredient in edible salt, it is commonly used as a condiment and food preservative.Contents[hide] 1 Properties 2 Production and use 2.1 Synthetic uses 2.2 Biological uses 2.3 Optical uses 2.4 Optical data 2.5 Household uses 2.6 Firefighting uses 2.7 In weather 3 Biological functions 4 Crystal structure 5 Road salt 5.1 Additives 5.2 Environmental impact 6 See also 7 References 8 External links

[edit] PropertiesThermal conductivity of pure NaCl as a function of temperature has a maximum of 2.03 W/(cm K) at 8 K and decreases to 0.069 at 314 K (41 C). It also decreases with doping.[3][edit] Production and use

Modern rock salt mine near Mount Morris, New York, United StatesSalt is currently mass-produced by evaporation of seawater or brine from other sources, such as brine wells and salt lakes, and by mining rock salt, called halite. In 2009, world production was estimated at 260 million metric tons, the top five producers (in million tonnes) being China (60.0), United States (46.0), Germany (16.5), India (15.8) and Canada (14.0).[4]As well as the familiar uses of salt in cooking, salt is used in many applications, from manufacturing pulp and paper, to setting dyes in textiles and fabric, to producing soaps, detergents, and other bath products. It is the major source of industrial chlorine and sodium hydroxide, and used in almost every industry.Sodium chloride is sometimes used as a cheap and safe desiccant because it appears to have hygroscopic properties, making salting an effective method of food preservation historically; the salt draws water out of bacteria through osmotic pressure, keeping it from reproducing, a major source of food spoilage. Even though more effective desiccants are available, few are safe for humans to ingest.Solubility of NaCl in various solvents(g NaCl / 1kg of solvent at 25 C)[5]

H2O360

Liquid ammonia30.2

glycerin83

propylene glycol71

Methanol14

Ethanol0.65

1-propanol0.124

2-propanol0.03

1-butanol0.05

1-pentanol0.018

Sulfolane0.05

Formic acid52

Acetone0.00042

Formamide94

Acetonitrile0.003

Dimethylformamide0.4

[edit] Synthetic usesUses of chlorine include PVC, pesticides and epoxy resins. Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the chemical equation2 NaCl + 2 H2O Cl2 + H2 + 2 NaOHSodium metal is produced commercially through the electrolysis of liquid sodium chloride. This is now done in a Down's cell in which sodium chloride is mixed with calcium chloride to lower the melting point below 700 C. As calcium is more electropositive than sodium, no calcium will be formed at the cathode. This method is less expensive than the previous method of electrolyzing sodium hydroxide.Sodium chloride is used in other chemical processes for the large-scale production of compounds containing sodium or chlorine. In the Solvay process, sodium chloride is used for producing sodium carbonate and calcium chloride. In the Mannheim process and in the Hargreaves process, it is used for the production of sodium sulfate and hydrochloric acid.[edit] Biological usesMany micro organisms cannot live in an overly salty environment: water is drawn out of their cells by osmosis. For this reason salt is used to preserve some foods, such as smoked bacon or fish. It can also be used to detach leeches that have attached themselves to feed. It is also used to disinfect wounds.[edit] Optical usesPure NaCl crystal is an optical compound with a wide transmission range from 200nm to 20m. It was often used in the infrared spectrum range and it is still used sometimes.While inexpensive, NaCl crystal is soft and hygroscopic. When exposed to free air, NaCl optics gradually covers with "frost". This limits application of NaCl to protected environments or for short-term uses such as prototyping.Today tougher crystals like zinc selenide (ZnSe) are used instead of NaCl (for the IR spectral range).[edit] Optical data Transmitivity: 92% (from 400nm to 13 m) Refractive Index: 1.494 at 10 m Reflection Loss: 7.5% at 10 m (2 surfaces) dN/dT: 36.2106/C at 0.7 m[edit] Household usesSince at least medieval times, people have used salt as a cleansing agent rubbed on household surfaces. It is also used in many brands of shampoo, and popularly to de-ice driveways and patches of ice.[edit] Firefighting uses

A class D fire extinguisher for various metalsSodium chloride is the principal extinguishing agent in fire extinguishers (Met-L-X, Super D) used on combustible metal fires such as magnesium, potassium, sodium, and NaK alloys (Class D). Thermoplastic powder is added to the mixture, along with waterproofing (metal stearates) and anti-caking materials (tricalcium phosphate) to form the extinguishing agent. When it is applied to the fire, the salt acts like a heat sink, dissipating heat from the fire, and also forms an oxygen-excluding crust to smother the fire. The plastic additive melts and helps the crust maintain its integrity until the burning metal cools below its ignition temperature. This type of extinguisher was invented in the late 1940s in the cartridge-operated type shown here, although stored pressure versions are now popular. Common sizes are 30lb. portable and 350lb. wheeled.[edit] In weather

Clouds above the PacificSmall particles of sea salt are the dominant cloud condensation nuclei well out at sea, which allow the formation of clouds in otherwise non-polluted air.[6] Snow removal by addition of salt (salting) is done to make travel easier and safer, and decrease the long term impact of a heavy snowfall on human populations. This process is done by both individual households and by governments and institutions and utilizes salts to eliminate snow from road surfaces and sidewalks.[7][edit] Biological functionsFor further information about sodium chloride in the diet, see Salt.In humans, a high-salt intake has long been suspected to generally raise blood pressure. More recently, it was demonstrated to attenuate nitric oxide production. Nitric oxide (NO) contributes to vessel homeostasis by inhibiting vascular smooth muscle contraction and growth, platelet aggregation, and leukocyte adhesion to the endothelium.[8][9][edit] Crystal structure

The crystal structure of sodium chloride. Each ion has six nearest neighbors, with octahedral geometry.

Close up view of NaCl crystalsMain article: Cubic crystal systemSodium chloride forms crystals with face-centered cubic symmetry. In these, the larger chloride ions, shown to the right as green spheres, are arranged in a cubic close-packing, while the smaller sodium ions, shown to the right as silver spheres, fill all the cubic gaps between them. Each ion is surrounded by six ions of the other kind; the surrounding ions are located at the vertices of a regular octahedron.This same basic structure is found in many other minerals and is commonly known as the halite or rock-salt crystal structure. It can be represented as a face-centered cubic (fcc) lattice with a two atom basis. The first atom is located at each lattice point, and the second atom is located half way between lattice points along the fcc unit cell edge.It is held together by an ionic bond which is produced by electrostatic forces arising from the difference in charge between the ions.[edit] Road saltWhile salt was once a scarce commodity in history, industrialized production has now made salt plentiful. Approximately 51% of world output is now used by cold countries to de-ice roads in winter, both in grit bins and spread by winter service vehicles. Calcium chloride is preferred over sodium chloride, since CaCl2 releases energy upon forming a solution with water, heating any ice or snow it is in contact with. It also lowers the freezing point, depending on the concentration. NaCl does not release heat upon solution; however, it does lower the freezing point. Calcium chloride is thought to be more environmentally friendly than sodium chloride when used to de-ice roads, however a drawback is that it tends to promote corrosion (of vehicles) more so than sodium chloride. NaCl is also more readily available and does not have any special handling or storage requirements, unlike calcium chloride. The salinity (S) of water is measured as grams salt per kilogram of water, and the freezing temperatures are as follows.S (g/kg)0153045597590106123140157175193212231250269290311331353

T (C)00.81.72.73.64.65.56.67.89.110.411.813.214.616.217.819.421.117.311.12.7

[edit] AdditivesMost table salt sold for consumption today is not pure sodium chloride. In 1911, magnesium carbonate was first added to salt to make it flow more freely.[10] In 1924 trace amounts of iodine in form of sodium iodide, potassium iodide or potassium iodate were first added, to reduce the incidence of simple goiter.[11]Salt for de-icing in the United Kingdom predominantly comes from a single mine in Winsford in Cheshire. Prior to distribution it has an anti-caking agent added: sodium hexacyanoferrate(II) at less than 100 ppm. This treatment enables rock salt to flow freely out of the gritting vehicles despite being stockpiled prior to use. In recent years this additive has also been used in table salt.[edit] Environmental impactRoad salt ends up in fresh water bodies and could harm aquatic plants and animals by disrupting their osmoregulation ability.[12] An alternative is to spread rough sand on ice so the surface is not slippery.The omnipresence of salt posts a problem in any coastal coating application, as trapped salts cause great problems in adhesion. Costs can reach staggering amounts. Naval authorities and ship builders keep a close eye on salt concentrations on surfaces during construction. Maximum salt concentrations on surfaces are dependent on the authority and application. The IMO regulation is mostly used and sets salt levels to a maximum of 50mg/m2 soluble salts measured as sodium chloride. These measurements are done by means of a Bresle test.Israeli and Jordanian salt evaporation ponds at the south end of the Dead Sea.Mounds of salt, Salar de Uyuni, Bolivia.Evaporation lagoons, Aigues-Mortes, France.

http://en.wikipedia.org/wiki/Sodium_chlorideLead(II) acetateFrom Wikipedia, the free encyclopedia(Redirected from Pb(CH3COO)2)Jump to: navigation, search Lead(II) acetate

IUPAC name[hide]Lead(II) acetate

Systematic name[hide]Lead(II) ethanoate

Other names[hide]Plumbous acetate, Salt of Saturn, Sugar of Lead, lead diacetate

Identifiers

CAS number301-04-2Y,6080-56-4 (trihydrate)51404-69-4 (Basic)

Properties

Molecular formulaPb(C2H3O2)2

Molar mass325.29 g/mol (anhydrous)379.33g/mol (trihydrate)

AppearanceWhite powder or colorless, efflorescent crystals

Density3.25 g/cm3 (anhydrous)2.55 g/cm3 (trihydrate)1.69 g/cm3 (decahydrate)

Melting point280 C (anhydrous)75 C (trihydrate)22 C (decahydrate)

Solubility in water44.39 g/100 mL (20 C)211 g/100 mL (50 C)[1]

Solubilityanhydrous soluble in alcohol; hydrates insoluble in alcohol

Refractive index (nD)1.567 (trihydrate)

Structure

Crystal structureMonoclinic

Hazards

Main hazardsNeurotoxic, Probable Human Carcinogen

NFPA 704130

Flash pointNonflammable


Infobox references

Lead(II) acetate is a chemical compound, a white crystalline substance with a sweetish taste. It is made by treating lead(II) oxide with acetic acid. Like other lead compounds, it is toxic. Lead acetate is soluble in water and glycerin. With water it forms the trihydrate, Pb(CH3COO)23H2O, a colorless or white efflorescent monoclinic crystalline substance. Lead(II) acetate is also known as lead acetate, lead diacetate, plumbous acetate, sugar of lead, lead sugar, salt of Saturn, and Goulard's powder (after Thomas Goulard).The substance is used as a reagent to make other lead compounds and as a fixative for some dyes. In low concentrations, it is the principal active ingredient in progressive types of hair coloring dyes. Lead(II) acetate is also used as a mordant in textile printing and dyeing, as a drier in paints and varnishes, and in preparing other lead compounds.Contents[hide] 1 Historical use 2 Other uses 3 Biological hazards 4 References 5 External links

[edit] Historical useLike other lead(II) salts, lead acetate has a sweet taste, which has led to its use as a sugar substitute throughout history. The ancient Romans, who had few sweeteners besides honey, would boil must (grape juice) in lead pots to produce a reduced sugar syrup called defrutum, concentrated again into sapa. This syrup was used to sweeten wine, and to sweeten and preserve fruit. It is possible that lead acetate or other lead compounds leaching into the syrup might have caused lead poisoning in anyone consuming it.[2]Pope Clement II died in October 1047. A recent toxicologic examination[citation needed] of his remains confirmed centuries-old rumors that he had been poisoned with lead sugar. It is not clear if he was assassinated.In 1787 painter Albert Christoph Dies swallowed, by accident, three-quarters of an ounce (21 g) of lead acetate. His recovery from this poison was slow and incomplete. He lived with illnesses until his death in 1822.[citation needed]Although its use was already illegal at that time, composer Ludwig van Beethoven may potentially have died of lead poisoning caused by wines adulterated with lead acetate.[3][4]Mary Seacole applied it among other remedies against an epidemic of cholera in Panama.[5][6]Sugar of lead has also been used to treat poison ivy.Lead acetate is no longer used as a sweetener in most of the world because of its recognized toxicity. Modern chemistry can easily detect it, which has all but stopped the illegal use that continued decades after legal use was terminated.[edit] Other usesLead acetate, as well as white lead, have been used in cosmetics throughout history, though this practice has ceased in Western countries.[7] It is still used in men's hair coloring products[8] like Grecian Formula.Lead acetate paper is used to detect the poisonous gas hydrogen sulfide. The gas reacts with lead(II) acetate on the moistened test paper to form a grey precipitate of lead(II) sulfide.Lead acetate solution was a commonly used folk remedy for sore nipples.[9]An aqueous solution of lead acetate is the byproduct of the 50/50 mixture of hydrogen peroxide and white vinegar used in the cleaning and maintenance of stainless steel firearm suppressors (silencers) and compensators. The solution is agitated by the bubbling action of the hydrogen peroxide, and the main reaction is the dissolution of lead deposits within the suppressor by the acetic acid, which forms lead acetate. Because of its high toxicity, this chemical solution must be appropriately disposed by a chemical processing facility or hazardous materials center. Alternatively, the solution may be reacted with sulfuric acid to precipitate insoluble lead sulfate. The solid may then be removed by mechanical filtration and is safer to dispose of than aqueous lead acetate.[edit] Biological hazardsLead(II) acetate, among other lead salts, has been reported to cross the placenta and to the embryo leading to fetal mortality. Lead salts also have teratogenic effect in some animal species.http://en.wikipedia.org/wiki/Pb%28CH3COO%292Dalam skala indsutri, timbal asetat dibuat dengan mereaksikan asam asetat dengan kepingan tipis timbal atau dengan timbal oksida dengan asam asetat kuat.Proses pertama dalam pembuatan timbal asetat adalah dengan memasukkan bahan baku ke dalam tungku, kemudian ditambahkan batu kapur, soda dan air kedalamnya dan dipanaskan pada suhu yang sangat tinggi. Setelah itu dimasukkan ke tangki chassing dan dilakukan pembakaran pada suhu rendah. Proses selanjutnya adalah melakukan de-silvering pada ketel yang telah berisi bahan olahan dan ditambah zinc. Kemudian dilakukan de-zincing pada ketel untuk menetralkan unsur zinc tadi. Setelah itu, diperolehlah timbal asetat.

Gambar 4.3. Flowsheet pembuatan Pb(CH3COO)2 Keterangan gambar:1. sintering2. pembakaran timbal tanur tinggi3. chassing4. pembakaran suhu rendah5. de-silvering ketel6. de-zincing ketel7. lasting ketel8. Pencetakan

Mercury(II) chlorideFrom Wikipedia, the free encyclopedia(Redirected from Mercuric chloride)Jump to: navigation, search Mercury(II) chloride

IUPAC name[hide]Mercury(II) chlorideMercury dichloride

Other names[hide]Mercuric chlorideCorrosive sublimate

Identifiers


EC number231-299-8

UN number1624

RTECS numberOV9100000

Properties

Molecular formulaHgCl2


Appearancewhite solid

Density5.43 g/cm3



Solubility in water7.4 g/100 ml (20 C)

Solubilitysoluble in alcohol, ether, acetone, ethyl acetateslightly soluble in benzene, CS2

Acidity (pKa)3.2 (0.2M solution)

Structure

Crystal structureorthogonal

Coordinationgeometrylinear

Molecular shapelinear

Dipole momentzero

Hazards

MSDSICSC 0979

EU Index080-010-00-X

EU classificationVery toxic (T+)Corrosive (C)Dangerous for the environment (N)

R-phrasesR28, R34, R48/24/25, R50/53

S-phrases(S1/2), S36/37/39, S45, S60, S61

NFPA 704040


Related compounds

Other anionsMercury(II) fluorideMercury(II) bromideMercury(II) iodide

Other cationsZinc chlorideCadmium chlorideMercury(I) chloride


Infobox references

Mercury(II) chloride or mercuric chloride (formerly corrosive sublimate), is the chemical compound with the formula HgCl2. This white crystalline solid is a laboratory reagent. It is not used as commonly as once was the case because, due to its solubility in water, it is highly toxic. It is a molecular compound.Contents[hide] 1 Production and basic properties 2 Applications 2.1 As a chemical reagent 2.2 Historic use in photography 2.3 Historic use in preservation 2.4 Historic use in medicine 3 Toxicity 4 References 5 External links

[edit] Production and basic propertiesMercuric chloride is not a salt but a linear triatomic molecule, hence its tendency to sublime. In the crystal, each mercury atom is bonded to two close chloride ligands with Hg---Cl distance of 2.38 ; four more chlorides are more distant at 3.38 .[1]Mercuric chloride is obtained by the action of chlorine on mercury or mercury(I) chloride, by the addition of hydrochloric acid to a hot, concentrated solution of mercury(I) compounds such as the nitrate:HgNO3 + 2 HCl HgCl2 + H2O + NO2,Heating a mixture of solid mercury(II) sulfate and sodium chloride also affords volatile HgCl2, which sublimes and condenses in the form of small rhombic crystals.Its solubility increases from 6% at 20 C to 36% in boiling water. In the presence of chloride ions, it dissolves to give the tetrahedral coordination complex [HgCl4]2-.[edit] ApplicationsThe main application of mercuric chloride is as a catalyst for the conversion of acetylene to vinyl chloride, the precursor to polyvinylchloride:C2H2 + HCl CH2=CHClFor this application, the mercuric chloride is supported on carbon in concentrations of about 5 weight percent. This technology has been eclipsed by the thermal cracking of 1,2-dichloroethane. Other significant applications of mercuric chloride include its use as a depolarizer in batteries and as a reagent in organic synthesis and analytical chemistry (see below).[2] It is being used in plant tissue culture for surface sterilisation of explants such as leaf or stem nodes.[edit] As a chemical reagentMercuric chloride is occasionally used to form an amalgam with metals, such as aluminium. Upon treatment with an aqueous solution of mercuric chloride, aluminium strips quickly become covered by a thin layer of the amalgam. Normally, aluminium is protected by a thin layer of oxide making it inert. Once amalgamated, aluminium can undergo a variety of reactions. For example, it will dissolve in water (this can be dangerous, as hydrogen gas and heat are generated). Halocarbons react with amalgamated aluminium in the Barbier reaction. These alkylaluminium compounds are nucleophilic and can be used in a similar fashion to the Grignard reagent. Amalgamated aluminium is also used as a reducing agent in organic synthesis. Zinc is also commonly amalgamated using mercuric chloride.Mercuric chloride is used to remove dithiane groups attached to a carbonyl in an umpolung reaction. This reaction exploits the high affinity of Hg2+ for anionic sulfur ligands.[edit] Historic use in photographyMercury(II) chloride was used as a photographic intensifier to produce positive pictures in the collodion process of the 1800s. When applied to a negative, the mercury(II) chloride whitens and thickens the image, thereby increasing the opacity of the shadows and creating the illusion of a positive image.[3][edit] Historic use in preservationFor the preservation of anthropological and biological specimens during the late 19th and early 20th centuries, objects were dipped in or were painted with a "mercuric solution". Objects in drawers were protected by scattering crystalline mercuric chloride over them.[4] It finds minor use in tanning, and wood was preserved by kyanizing (soaking in mercuric chloride).[5] Mercuric chloride was one of the three chemicals used for railroad tie wood treatment between 1830 and 1856 in Europe and the United States. Limited railroad ties were treated in the United States until there were concerns over lumber shortages in the 1890s.[6] The process was generally abandoned because mercuric chloride was water soluble and not effective for the long term, as well as poisonous. Furthermore, alternative treatment processes, such as copper sulfate, zinc chloride, and ultimately creosote; were found to be less toxic. Limited kyanizing was used for some railroad ties in the 1890s and early 1900s.[7][edit] Historic use in medicineSyphilis was frequently treated with mercuric chloride before the advent of antibiotics. It was inhaled, ingested, injected, and applied topically. Poisoning was so common that its symptoms were confused with those of syphilis. This usage of 'salts of white mercury' is referred to in the English folk-song, The Unfortunate Rake.[8][edit] ToxicityMain article: Mercury poisoningMercuric chloride is highly toxic, not only acutely but as a cumulative poison.http://en.wikipedia.org/wiki/Mercuric_chlorideDalam skala industri, pembuatan merkuri klorida dilakukan dengan cara sebagai berikut, pertama sekali unsur merkuri (stream A) dipompa dari tangki dimana bereaksi dengan gas klorin yang berlebihan (stream B). Hasil produksi (stream C) disalurkan ke unit presipitasi yang menghasilkan produk kering (HgCl2) dan merkuri klorida (stream D) dibungkus dan disegel dalam drum untuk dikirim ke luar sebagai produk utama merkuri klorida (HgCl2). Dari reaktor (stream E) dikirim ke scrubber kaustik dimana merkuri yang tidak tereaksikan dipulihkan kembali dan kemudian siklus kembali dilakukan (stream F).

Gambar 3.4 Flowsheet pembuatan HgCl2

PerakDari Wikipedia bahasa Indonesia, ensiklopedia bebasBelum DiperiksaLangsung ke: navigasi, cari Untuk kegunaan lain dari Perak, lihat Perak (disambiguasi).47paladium perak kadmium

CuAgAuTabel periodik

Keterangan Umum Unsur

Nama, Lambang, Nomor atomperak, Ag, 47

Deret kimialogam transisi

Golongan, Periode, Blok11, 5, d

Penampilanlogam putih mengkilap

Massa atom107,8682(2) g/mol

Konfigurasi elektron[Kr] 4d10 5s1

Jumlah elektron tiap kulit2, 8, 18, 18, 1

Ciri-ciri fisik

Fasepadat

Massa jenis (sekitar suhu kamar)10,49 g/cm

Massa jenis cair pada titik lebur9,320 g/cm

Titik lebur1234,93 K(961,78 C, 1763,2 F)

Titik didih2435 K(2162 C, 3924 F)

Kalor peleburan11,28 kJ/mol

Kalor penguapan258 kJ/mol

Kapasitas kalor(25 C) 25,350 J/(molK)

Tekanan uap

P/Pa1101001 k10 k100 k

pada T/K128314131575178220552433

Ciri-ciri atom

Struktur kristalkubus pusat muka

Bilangan oksidasi1(oksida amfoter)

Elektronegativitas1,93 (skala Pauling)

Energi ionisasipertama: 731,0 kJ/mol

ke-2: 2070 kJ/mol

ke-3: 3361 kJ/mol

Jari-jari atom160 pm

Jari-jari atom (terhitung)165 pm

Jari-jari kovalen153 pm

Jari-jari Van der Waals172 pm

Lain-lain

Sifat magnetikdiamagnetik

Resistivitas listrik(20 C) 15,87 nm

Konduktivitas termal(300 K) 429 W/(mK)

Difusivitas termal(300 K) 174 mm/s

Ekspansi termal(25 C) 18,9 m/(mK)

Kecepatan suara(pada wujud kawat)(suhu kamar) 2680 m/s

Modulus Young83 GPa

Modulus geser30 GPa

Modulus ruah100 GPa

Nisbah Poisson0,37

Skala kekerasan Mohs2,5

Kekerasan Vickers251 MPa

Kekerasan Brinell24,5 MPa

Isotop

isoNAwaktu paruhDMDE (MeV)DP

105Agsyn41,2 hari-105Pd

0,344; 0,280;0,644; 0,443-

106mAgsyn8,28 hari-106Pd

0,511; 0,717;1,045; 0,450-

107Ag51,839%Ag stabil dengan 60 neutron

108mAgsyn418 tahun-108Pd

IT0,109108Ag

0,433; 0,614;0,722-

109Ag48,161%Ag stabil dengan 62 neutron

111Agsyn7,45 hari-1,036; 0,694111Cd

0,342-

Referensi

Di dalam industri, perak sering diproduksi dari bahan-bahan mineral logam. Maka biasanya bahan-bahan mineral logam biasanya akan dilakukan bleaching. Proses pemurnian bahan-bahan batuan mineral itu akanmenghasilkan renin yang dapat digunakan. Kemudian zat tersebut dicampurkan dengan zat asam maupun basa sehingga akhirnya akan terbentuk suatu endapan perak.Gambar 2.9 Flowchart pembentukan Perakhttp://id.wikipedia.org/wiki/PerakAmmonium chlorideFrom Wikipedia, the free encyclopediaJump to: navigation, search Ammonium chloride

IUPAC name[hide]Ammonium chloride

Other names[hide]Sal ammoniac

Identifiers


ChemSpider23807

UNII01Q9PC255DY

EC number235-186-4

RTECS numberBP4550000

SMILES[show][Cl-].[NH4+]

InChI[show]1/ClH.H3N/h1H;1H3

InChI key[show]NLXLAEXVIDQMFP-UHFFFAOYAI

Properties

Molecular formulaNH4Cl


AppearanceWhite solidhygroscopic

Odorodorless

Density1.5274 g/cm3

Melting point338 C (decomposes)

Solubility in water29.7 g/100 mL (0 C)37.2 g/100 mL (20 C)77.3 g/100 mL (100 C)

Solubility in alcohol0.6 g/100 mL (19 C)

Acidity (pKa)9.245

Refractive index (nD)1.642

Thermochemistry

Std enthalpy offormation fHo298314.55 kJ/mol[1]

Standard molarentropy So29894.85 JK1mol1 [1]

Hazards

MSDSICSC 1051

EU Index017-014-00-8

EU classificationHarmful (Xn)Irritant (Xi)

R-phrasesR22, R36

S-phrases(S2), S22

NFPA 704010


LD501650 mg/kg, oral (rat)

Related compounds

Other anionsAmmonium fluorideAmmonium bromideAmmonium iodide

Other cationsSodium chloridePotassium chlorideHydroxylamonium chloride


Infobox references

Ammonium chloride NH4Cl (also Sal ammoniac, salmiac, nushadir salt, sal armagnac, sal armoniac, salt armoniack) is, in its pure form, a clear white water-soluble crystalline salt of ammonia. The aqueous ammonium chloride solution is mildly acidic. Sal ammoniac is a name of natural, mineralogical form of ammonium chloride. The mineral is especially common on burning coal dumps (formed by condensation of coal-derived gases), but also on some volcanoes.Contents[hide] 1 Sources 2 Reactions 3 Applications 3.1 Chemistry and Physics 3.2 Biology and Agriculture 3.3 Pyrotechnics 3.4 Textile and Leather 3.5 Metalwork 3.6 Medicine 3.7 Food 3.8 Other Applications 4 References

[edit] SourcesThe substance occurs naturally in volcanic regions, forming on volcanic rocks near fume-releasing vents. The crystals deposit directly from the gaseous state, and tend to be short-lived, as they dissolve easily in water. It is a by-product of the Solvay process used to produce sodium carbonate.[2]Ammonium chloride is prepared commercially by reacting ammonia (NH3) with hydrogen chloride (HCl). As these chemicals are corrosive, this process has to be performed in vessels lined with nonreactive materials (e.g. glass, enamel, lead, or PVC).[2]NH3 + HCl NH4ClThis reaction can occur if poorly sealed bottles of household ammonia (ammonium hydroxide) and hydrochloric acid are stored in close proximity, leading to crystals forming around the openings of the bottles (mostly appear on those leaking more slowly).[edit] ReactionsAmmonium chloride appears to sublime but this process actually involves decomposition into ammonia and hydrogen chloride gas.[2]NH4Cl NH3 + HClAmmonium chloride may be reacted with a hydroxide base, e.g. sodium hydroxide, to release ammonia gas:NH4Cl + NaOH NH3 + NaCl + H2OIf test tubes of ammonia solution and hydrochloric acid are brought close together, a smoke composed of microcrystals of ammonium chloride will slowly rise out of the tube.[edit] Applications

Ammonium chloride crystal(s).[edit] Chemistry and PhysicsAmmonium chloride dissolved in water becomes an acid. Ammonium chloride is used to produce low temperatures in cooling baths. For example, the zero point of Fahrenheit temperature scale is determined by placing the thermometer in a mixture of ice, water, and ammonium chloride. Ammonium chloride solutions with ammonia are used as buffer solutions.[edit] Biology and AgricultureIn biological applications ammonium chloride acts as a nitrogen source and is used in fertilizers, as a feed supplement for cattle and as an ingredient in nutritive media for yeast microbiological organisms.

[edit] PyrotechnicsAmmonium chloride is an ingredient in fireworks and safety and contact explosives.[edit] Textile and LeatherAmmonium chloride is used in the textile and leather industry in dyeing, tanning textile printing and to luster cotton.[edit] MetalworkAmmonium chloride is used as a flux in preparing metals to be tin coated, galvanized or soldered. It works as a flux by cleaning the surface of workpieces by reacting with the metal oxides at the surface to form a volatile metal chloride. It is sold in blocks at hardware stores for use in cleaning the tip of a soldering iron and can also be included in solder as flux.[edit] MedicineAmmonium chloride is used as an expectorant in cough medicine. Its expectorant action is caused by irritative action on the bronchial mucosa. This causes the production of excess respiratory tract fluid which presumably is easier to cough up.Ammonium salts are an irritant to the gastric mucosa and may induce nausea and vomiting.Ammonium chloride is used as a systemic acidifying agent in treatment of severe metabolic alkalosis, in oral acid loading test to diagnose distal renal tubular acidosis, to maintain the urine at an acid pH in the treatment of some urinary-tract disordersAmmonium Chloride is also used as a diuretic in forced acid diuresis.[edit] FoodIn several countries ammonium chloride is known as sal ammoniac and used as food additive. The E number for ammonium chloride used as a food additive is E510.Sal ammoniac is used to spice up dark sweets called salty liquorice and in the flavouring Salmiakki Koskenkorva for vodkas.Sal ammoniac is also used in baking to give cookies a very crisp texture.[edit] Other ApplicationsAmmonium chloride is used in a ~5% aqueous solution to work on oil wells with clay swelling problems. It is also used as electrolyte in Zinccarbon batteries. Other uses include in hair shampoo, in the glue that bonds plywood, in cleaning products.http://en.wikipedia.org/wiki/Ammonium_chlorideAmmonium hydroxideFrom Wikipedia, the free encyclopediaJump to: navigation, search Main article: AmmoniaAmmonium hydroxide

Identifiers


ChemSpider14218

UNII5138Q19F1XY

SMILES[show][OH-].[NH4+]

InChI[show]1/H3N.H2O/h1H3;1H2

InChI key[show]VHUUQVKOLVNVRT-UHFFFAOYAI

Properties[1]

Molecular formulaNH4OH


Appearancevery volatile solution, colorless, bitter smell

Density0.91 g/cm3 (25%)0.88 g/cm3 (32%)

Melting point57.5 C (25%)91.5 C (32%)

Boiling point37.7 C (25%)24.7 C (32%)

Solubility in waterMiscible

Hazards[2]

EU classificationCorrosive (C)Dangerous to the environment (N)

R-phrasesR34, R50

S-phrases(S1/2), S26, S36/37/39, S45, S61

Related compounds

Other anionsAmmonium chlorideAmmonium cyanide

Other cationsTetramethylammonium hydroxide

Related compoundsAmmoniaHydroxylamine


Infobox references

Ammonium hydroxide, also known as ammonia water, ammonical liquor, ammonia liquor, aqua ammonia, aqueous ammonia, or simply ammonia, is a solution of ammonia in water. It can be denoted by the symbols NH3(aq). Although its name suggests a salt with composition [NH4+][OH], it is not actually possible to isolate samples of NH4OH it exists only in dilute aqueous solutions.[3]Contents[hide] 1 Basicity of ammonia in water 2 Saturated solutions 3 Applications 4 Laboratory use 5 See also 6 References

[edit] Basicity of ammonia in waterIn aqueous solution, ammonia deprotonates a small fraction of the water to give ammonium and hydroxide according to the following equilibrium:NH3 + H2O NH4+ + OH.In a 1M ammonia solution, about 0.42% of the ammonia is converted to ammonium, equivalent to a pH of 11.63. The base ionization constant isKb = [NH4+][OH-]/[NH3] = 1.8105[edit] Saturated solutionsLike other gases, ammonia exhibits decreasing solubility in solvent liquids as the temperature of the solvent increases. "Ammonium hydroxide" solutions decrease in density as the concentration of dissolved ammonia increases. At 15.6C (60.1F), the density of a saturated solution is 0.88 g/ml and contains 35% ammonia by mass, 308 g/l w/v, (308grams of ammonia per litre of solution) and has a molarity of approximately 18mol L1. At higher temperatures, the molarity of the saturated solution decreases and the density increases.When solutions that are saturated at cold temperatures are sealed in containers and subsequently warmed, the concentration of the solution decreases and the vapor pressure of ammonia gas increases. Unsealing such containers can lead to a burst of ammonia gas. In extreme cases, the containers could rupture.From a laboratory perspective, one should be aware that the concentration of a saturated solution is continually dropping as the container is handled in a warmer environment. Thus, old samples of ammonium hydroxide will deviate from 18 M, as can be verified by titration.[edit] ApplicationsHousehold ammonia is dilute ammonium hydroxide, which is also an ingredient of numerous other cleaning agents, including many window cleaning formulas. In addition to use as an ingredient in cleansers with other cleansing ingredients, ammonium hydroxide in water is also sold as a cleaning agent by itself, usually labelled as simply "ammonia". It may be sold plain, lemon-scented (and typically colored yellow), or pine-scented (green).In industry, ammonium hydroxide is used as a precursor to some alkyl amines, although anhydrous ammonia is usually preferred. Hexamethylenetetramine forms readily from aqueous ammonia and formaldehyde. Ethylenediamine forms from 1,2-dichloroethane and aqueous ammonia.[4]In furniture-making, ammonium hydroxide was traditionally used to darken or stain wood containing tannic acid. Tannic acid with ammonium hydroxide or iron salts creates a brown stain which can be applied to wood. [5]Ammonium hydroxide is used in the meat packing industry. Some companies treat their beef "with a pH enhancement process that forms ammonium hydroxide in the finished product."[edit] Laboratory useAqueous ammonia is used in traditional qualitative inorganic analysis as a complexant and base. Like many amines, it gives a deep blue coloration with copper(II) solutions. Ammonia solution can dissolve silver residues, such as that formed from Tollens' reagent.When ammonium hydroxide is mixed with dilute hydrogen peroxide in the presence of a metal ion, such as Cu2+, the peroxide will undergo rapid decomposition.http://en.wikipedia.org/wiki/Ammonium_hydroxideNatrium hidroksidaDari Wikipedia bahasa Indonesia, ensiklopedia bebas(Dialihkan dari NaOH)Belum DiperiksaLangsung ke: navigasi, cari Natrium Hidroksida

Nama IUPAC[sembunyikan]Natrium Hidroksida

Nama lain[sembunyikan]Soda kaustik

Identifikasi

Nomor CAS[1310-73-2]

Sifat

Rumus molekulNaOH

Massa molar39,9971 g/mol

Penampilanzat padat putih

Densitas2,1 g/cm, padat

Titik leleh318C (591 K)

Titik didih1390C (1663 K)

Kelarutan dalam air111 g/100 ml (20C)

Kebasaan (pKb)-2,43

Bahaya

MSDSExternal MSDS

NFPA 704031

Titik nyalaTidak mudah terbakar.

Senyawa terkait

Alkali hidroksida terkaitLitium hidroksidaKalium hidroksidaRubidium hidroksidaSesium hidroksida

Kecuali dinyatakan sebaliknya, data di atas berlakupada temperatur dan tekanan standar (25C, 100kPa)Sangkalan dan referensi

Natrium hidroksida (NaOH), juga dikenal sebagai soda kaustik atau sodium hidroksida, adalah sejenis basa logam kaustik. Natrium Hidroksida terbentuk dari oksida basa Natrium Oksida dilarutkan dalam air. Natrium hidroksida membentuk larutan alkalin yang kuat ketika dilarutkan ke dalam air. Ia digunakan di berbagai macam bidang industri, kebanyakan digunakan sebagai basa dalam proses produksi bubur kayu dan kertas, tekstil, air minum, sabun dan deterjen. Natrium hidroksida adalah basa yang paling umum digunakan dalam laboratorium kimia.Natrium hidroksida murni berbentuk putih padat dan tersedia dalam bentuk pelet, serpihan, butiran ataupun larutan jenuh 50%. Ia bersifat lembab cair dan secara spontan menyerap karbon dioksida dari udara bebas. Ia sangat larut dalam air dan akan melepaskan panas ketika dilarutkan. Ia juga larut dalam etanol dan metanol, walaupun kelarutan NaOH dalam kedua cairan ini lebih kecil daripada kelarutan KOH. Ia tidak larut dalam dietil eter dan pelarut non-polar lainnya. Larutan natrium hidroksida akan meninggalkan noda kuning pada kain dan kertas.http://id.wikipedia.org/wiki/NaOHNitric acidFrom Wikipedia, the free encyclopedia(Redirected from HNO3)Jump to: navigation, search Nitric acid

Preferred IUPAC name[hide]Nitric acid

Systematic name[hide]Oxoazinic acid

Other names[hide]Aqua fortisSalpetre acidSpirit of nitre

Identifiers

CAS number7697-37-2Y, 43625-06-5(2H)Y, 13587-52-5(2H)Y

PubChem944Y, 10313048(15N)Y, 12025424(2H)Y

ChemSpider919Y, 8488513(15N)Y, 21171471(2H)Y

EC number231-714-2

UN numberUN 2031

KEGGC00244

MeSHNitric+Acid

ChEBI48107

RTECS numberQU5775000

SMILES[show]ON(=O)=O

InChI[show]1S/HNO3/c2-1(3)4/h(H,2,3,4)

InChI key[show]GRYLNZFGIOXLOG-UHFFFAOYSA-N

Gmelin Reference1576

3DMetB00068

Properties

Molecular formulaHNO3


AppearanceClear, colorless liquid

Density1.5129 g/cm3

Melting point-42C, 231K, -44F

Boiling point83C, 356K, 181F (bp of pure acid. 68% solution boils at 120.5C)

Solubility in watercompletely miscible

Acidity (pKa)-1.4

Refractive index (nD)1.397 (16.5 C)

Dipole moment2.17 0.02 D

Hazards

MSDSICSC 0183PCTL Safety Website

EU Index007-004-00-1

EU classificationToxic (T)Corrosive (C)Oxidant (O)

R-phrasesR8 R35

S-phrases(S1/2) S23 S26 S36 S45

NFPA 704040OX


Related compounds

Other anionsNitrous acid

Other cationsSodium nitratePotassium nitrateAmmonium nitrate

Related compoundsDinitrogen pentoxide


Infobox references

Nitric acid (HNO3), also known as aqua fortis and spirit of nitre, is a highly corrosive and toxic strong acid.Colorless when pure, older samples tend to acquire a yellow cast due to the accumulation of oxides of nitrogen. If the solution contains more than 86% nitric acid, it is referred to as fuming nitric acid. Fuming nitric acid is characterized as white fuming nitric acid and red fuming nitric acid, depending on the amount of nitrogen dioxide present. At concentrations above 95% at room temperature, nitric acid tends to rapidly develop a yellow color due to decomposition. Nitric acid is also commonly used as a strong oxidizing agent.Contents[hide] 1 Properties 1.1 Acidic properties 1.2 Oxidizing properties 1.2.1 Reactions with metals 1.2.2 Passivation 1.2.3 Reactions with non-metals 1.3 Xanthoproteic test 2 Grades 3 Industrial production 4 Laboratory synthesis 5 Uses 5.1 Rocket fuel 5.2 Chemical reagent 5.3 Woodworking 6 Other uses 7 Safety 8 History 9 References 10 External links

[edit] PropertiesPure anhydrous nitric acid (100%) is a colorless mobile liquid with a density of 1.522 g/cm3 which solidifies at 42 C to form white crystals and boils at 83 C. When boiling in light, and slowly even at room temperature, there is a partial decomposition with the formation of nitrogen dioxide following the reaction:4 HNO3 2 H2O + 4 NO2 + O2which means that anhydrous nitric acid should be stored below 0 C to avoid decomposition. The nitrogen dioxide (NO2) remains dissolved in the nitric acid coloring it yellow, or red at higher temperatures. While the pure acid tends to give off white fumes when exposed to air, acid with dissolved nitrogen dioxide gives off reddish-brown vapors, leading to the common name "red fuming acid" or "fuming nitric acid". Fuming nitric acid is also referred to as 16 molar nitric acid. It is the most concentrated form of nitric acid at Standard Temperature and Pressure (STP).Nitric acid is miscible with water and distillation gives a maximum-boiling azeotrope with a concentration of 68% HNO3 and a boiling temperature of 120.5 C at 1 atm, which is the ordinary concentrated nitric acid of commerce. Two solid hydrates are known; the monohydrate (HNO3H2O) and the trihydrate (HNO33H2O).Nitrogen oxides (NOx) are soluble in nitric acid and this property influences more or less all the physical characteristics depending on the concentration of the oxides. These mainly include the vapor pressure above the liquid and the boiling temperature, as well as the color mentioned above.Nitric acid is subject to thermal or light decomposition with increasing concentration and this may give rise to some non-negligible variations in the vapor pressure above the liquid because the nitrogen oxides produced dissolve partly or completely in the acid.[edit] Acidic propertiesBeing a typical strong acid, nitric acid reacts with alkalis, basic oxides, and carbonates to form salts, such as ammonium nitrate. Due to its oxidizing nature, nitric acid generally does not donate its proton (that is, it does not liberate hydrogen) on reaction with metals and the resulting salts are usually in the higher oxidized states. For this reason, heavy corrosion can be expected and should be guarded agains

literatur

Documents

place of sodium hydroxide

calcium carbonate caco3

gcm3 decahydrate

sodium salt of carbonic

monohydrate34 c decomp

c45 g100 g

soda lime glass

water22 g100