KARBOL

Berikut kami akan memberikan cara membuat karbol yang biasa kita gunakan dalam kehidupan sehari-hari, semoga bermanfaat :

Bahan-bahan yang dibutuhkan :

• Texapone 1 Kg
• NaCL 2 Kg
• Pine oil 1 L
• H2O +- 15 L

Alat-alat yang dibutuhkan :

• 2 Buah wadah yang terbuat dari bahan plastik
• 1 Buah tongkat yang terbuat dari bahan kayu yang berfungsi sebagai alat pengaduk

Cara membuat :

• Aduk texapone di wadah 1 hingga berubah warna menjadi putih tanpa menggunakan bahan yang lain.
• Larutkan NaCL dengan +- 2 L H2O di wadah 2 aduk hingga larut.
• Campurkan larutan di wadah 2 ke dalam wadah 1 aduk hingga merata dan mengental
• Tambahkan +- 13 L H2O sedikit demi sedikit aduk hingga mengental dan rata
• Tambahkan Pine oil aduk hingga merata, reaksi yang terjadi adalah bahan yang sudah kita aduk akan menjadi kembali mencair, tapai jangan khawatir itu adalah reaksi yang biasa terjadi
• diamkan adukan +- 2 jam, karbol sudah siap digunakan

NB : Karbol yang anda buat masih terbilang pekat maka untuk menggunakannya sebaiknya dicampurkan air kembali sesuai dengan kebutuhan. Selamat mencoba !!!

LIQUID SOAP

Berikut kami berikan cara membuat liquid soap atau sabun cair, semoga bermanfaat.

Bahan-bahan yang dibutuhkan :

• Texapone gell 1 kg
• Natrium Sulfate 1kg
• H2o +- 15 liter
• Comperland 30 ml
• Parfume secukupnya
• Pewarna secukupnya

Alat-alat yang dibutuhkan :

• 2 buah wadah yang terbuat dari bahan plastik, ini berfungsi untuk wadah pengadukan
• 1 buah tongkat yang terbuat dari ahan kayu, ini berfungsi untuk alat pengaduk

Cara membuat :

• Aduk texapone di wadah 1 tanpa menggunakan air hingga berubah warna menjadi putih
• Larutkan natrium sulfate dengan 2 liter air aduk hingga merata di wadah 2
• Campurkan larutan wadah 2 dengan bahan di wadah 1 aduk hingga mengental dan merata
• Tambahkan 13 liter air ke dalam bahan campuran sedikit demi sedikit aduk hingga mengental dan merata.
• Campurkan parfume, pewarna, dan comperland ke dalam bahan campuran, aduk hingga rata
• Apabila sabun sudah mengental hentikan adukan karena dapat menyebabkan sabun yang sudah dibuat akan kembali mencair.
• Diamkan sabun +- 2 jam, sabun sudah siap digunakan.

TRICHLOROETHYLENE (TCE)

The chemical compound trichloroethylene is a chlorinated hydrocarbon commonly used as an industrial solvent. It is a clear non-flammable liquid with a sweet smell.

Its IUPAC name is trichloroethene. In industry, it is informally referred to by the abbreviations TCE, Trike and tri, and it is sold under a variety of trade names. In addition to its industrial uses, trichloroethylene was used from about 1930 as a volatile anesthetic and analgesic in millions of patients, before its toxic properties were realized.

History

Pioneered by Imperial Chemical Industries in Britain, its development was hailed as a revolution: lacking the great hepatotoxic liability of chloroform and the unpleasant pungency and flammability of ether, it nonetheless had several pitfalls, including the sensitization of the myocardium to epinephrine, potentially acting in an arrhythmogenic manner. Its low volatility demanded the employment of carefully regulated heat in its vaporization. Research demonstrating its transient elevation of serum hepatic enzymes raised concerns regarding its hepatotoxic potential. Several deaths occurred as a result, though the incidence was comparable to that of halothane hepatitis. Incompatibility with soda lime (the CO2 adsorbent utilized in closed-circuit, low-flow rebreathing systems) was also a concern. TCE was readily decomposed into 1,2-dichloroacetylene, a neurotoxic compound potentially responsible for its hepatotoxic potential, though its metabolite trichloroacetic acid is more probably the etiological source of the latter. Halothane usurped a great portion of its market in 1956, with its total abandonment not achieved until the 1980s, when its use as an analgesic in obstetrics was implicated in fetal death. Concerns of its carcinogenic potential were raised simultaneously.

Due to concerns about its toxicity, the use of trichloroethylene in the food and pharmaceutical industries has been banned in much of the world since the 1970s. Legislation has forced the substitution of trichloroethylene in many process in Europe as the chemical was classified as a carcinogen carrying an R45 risk phrase. Many alternatives are being promoted such as Ensolv and Leksol, however each of these is based on nPropyl Bromide which carries an R60 risk phrase and they would not be a legally acceptable substitute.

Production

Prior to the early 1970s, most trichloroethylene was produced in a two-step process from acetylene. First, acetylene was treated with chlorine using a ferric chloride catalyst at 90 °C to produce 1,1,2,2-tetrachloroethane according to the chemical equation
HC≡CH + 2 Cl2 → Cl2CHCHCl2

The 1,1,2,2-tetrachloroethane is then dehydrochlorinated to give trichloroethylene. This can either be accomplished with an aqueous solution of calcium hydroxide
2 Cl2CHCHCl2 + Ca(OH)2 → 2 ClCH=CCl2 + CaCl2 + 2 H2O

or in the vapor phase by heating it to 300-500°C on a barium chloride or calcium chloride catalyst
Cl2CHCHCl2 → ClCH=CCl2 + HCl

Today, however, most trichloroethylene is produced from ethylene. First, ethylene is chlorinated over a ferric chloride catalyst to produce 1,2-dichloroethane.
CH2=CH2 + Cl2 → ClCH2CH2Cl

When heated to around 400 °C with additional chlorine, 1,2-dichloroethane is converted to trichloroethylene
ClCH2CH2Cl + 2 Cl2 → ClCH=CCl2 + 3 HCl

This reaction can be catalyzed by a variety of substances. The most commonly used catalyst is a mixture of potassium chloride and aluminum chloride. However, various forms of porous carbon can also be used. This reaction produces tetrachloroethylene as a byproduct, and depending on the amount of chlorine fed to the reaction, tetrachloroethylene can even be the major product. Typically, trichloroethylene and tetrachloroethylene are collected together and then separated by distillation.

Uses

Trichloroethylene is an effective solvent for a variety of organic materials. When it was first widely produced in the 1920s, its major use was to extract vegetable oils from plant materials such as soy, coconut, and palm. Other uses in the food industry included coffee decaffeination and the preparation of flavoring extracts from hops and spices. It was also used as a dry cleaning solvent, although tetrachloroethylene (also known as perchloroethylene) surpassed it in this role in the 1950s.

Trichloroethylene has been widely used as a degreaser for metal parts. In the late 1950s, the demand for trichloroethylene as a degreaser began to decline in favor of the less toxic 1,1,1-trichloroethane. Another problem with trichloroethylene is that it’s just too good a solvent in many mechanical applications, as it easily will strip many paints almost instantly and dissolves some plastics. However, 1,1,1-trichloroethane production has been phased out in most of the world under the terms of the Montreal Protocol, and as a result trichloroethylene has experienced a resurgence in use. It has also been used for drying out the last bit of water for production of 100% ethanol.

Trichloroethylene (Trimar and Trilene) was used as a volatile gas anesthetic from the 1930s through the 1960s in Europe and North America. Supplanting chloroform and ether for a significant period of time, trichloroethylene demonstrated superior efficacy in induction times and cost-effectiveness. It retained use in other locations well into the 1990s. It was known for its favorable analgesic properties. Induction of general anesthesia was accomplished by administering up to 1%(v/v) vapor. Equilibration would often result in patient levels of 0.1 to 0.5% vapor. Many patients were given Trilene inhalers to self administer analgesia, especially in obstetrical labor. The number of patients exposed to these high levels of trichloroethylene is difficult to know, but is certainly well into the millions.

Chemical instability

Although it has proven useful as a metal degreaser, trichloroethylene itself is unstable in the presence of metal over prolonged exposure. As early as 1961, this phenomenon was clearly recognized by the manufacturing industry, since an additive was instilled in the commercial formulation of trichloroethylene. The reactive instability is accentuated by higher temperatures, so that the search for stabilizing additives is conducted by heating trichloroethylene to its boiling point in a reflux condenser and observing decomposition. The first widely used stabilizing additive was dioxane; however, its use was patented by Dow Chemical Company and could not be used by other manufacturers. Considerable research took place in the 1960s to develop alternative stabilizers for trichlorethylene. The principal family of chemicals that showed promise was the ketone family, such as methyl ethyl ketone. Considerable research was conducted at Frontier Chemical Company, Wichita, Kansas on this class of ketones using reflux condensation experiments.

Physiological effects

When inhaled, trichloroethylene, as with any anesthetic gas, depresses the central nervous system. Its symptoms are similar to those of alcohol intoxication, beginning with headache, dizziness, and confusion and progressing with increasing exposure to unconsciousness . Respiratory and circulatory depression from any anesthetic can result in death if administration is not carefully controlled. As mentioned above, cardiac sensitization to catecholamines such as epinephrine can result in dangerous cardiac arrhythmias. Caution should be exercised anywhere a high concentration of trichloroethylene vapors may be present; the drug can desensitize the nose to its scent, and it is possible to unknowingly inhale harmful or lethal amounts of the vapor.

Much of what is known about the human health effects of trichloroethylene is based on occupational exposures. Beyond the effects to the central nervous system, workplace exposure to trichloroethylene has been associated with toxic effects in the liver and kidney . Over time, occupational exposure limits on trichloroethylene have tightened, resulting in more stringent ventilation controls and personal protective equipment use by workers. The tightening of occupational exposure limits and increased need for worker protection in part contributed to the substitution of other lower toxicity chemicals for trichloroethylene in solvent cleaning and degreasing.

The carcinogenicity of trichloroethylene was first evaluated in laboratory animals in the 1970s.Cancer bioassays performed by the National Cancer Institute (later the National Toxicology Program) showed that exposure to trichloroethylene is carcinogenic in animals, producing liver cancer in mice, and kidney cancer in rats . Numerous epidemiological studies have been conducted on trichloroethylene exposure in the workplace, with differing opinions regarding the strength of evidence between trichloroethylene and human cancer.[citation needed] Recent studies on the mechanisms of carcinogenicity have shown that metabolism of trichloroethylene in the liver produces metabolites (such as trichloroacetic acid and dichloroacetic acid, which are responsible for liver tumors in mice) that are the ultimate carcinogens in laboratory animals.Other studies using physiologically-based pharmacokinetic (PBPK) modeling, have examined the similarities and differences in metabolism between humans and laboratory animals,to better understand the relationship between carcinogenicity observed in laboratory animals and human cancer risks. The National Toxicology Program’s 11th Report on Carcinogens categorizes trichloroethylene as “reasonably anticipated to be a human carcinogen”, based on limited evidence of carcinogenicity from studies in humans and sufficient evidence of carcinogenicity from studies in experimental animals.

One recent review of the epidemiology of kidney cancer rated cigarette smoking and obesity as more important risk factors for kidney cancer than exposure to solvents such as trichloroethylene. In contrast, the most recent overall assessment of human health risks associated with trichloroethylene states, “[t]here is concordance between animal and human studies, which supports the conclusion that trichloroethylene is a potential kidney carcinogen”. The evidence appears to be less certain at this time regarding the relationship between humans and liver cancer observed in mice, with the NAS suggesting that low-level exposure might not represent a significant liver cancer risk in the general population. However the NAS also concluded that higher levels of exposure, such as workplace exposure, or locations with significant environmental contamination, might be associated with a liver cancer risk in humans.

Recent studies in laboratory animals and observations in human populations suggest that exposure to trichloroethylene might be associated with congenital heart defects (J Am Coll Cardiol. 1990 Jul;16(1):155-64.; J Am Coll Cardiol. 1993 May;21(6):1466-72; Toxicol Sci. 2000 Jan;53(1):109-17; Birth Defects Res A Clin Mol Teratol. 2003 Jul;67(7):488-95; Environ Health Perspect. 2006 Jun;114(6):842-7). While it is not clear what levels of exposure are associated with cardiac defects in humans, there is consistency between the cadiac defects observed in studies of communities exposed to trichloroethylene contamination in groundwater, and the effects observed in laboratory animals. Trichloroethylene can also affect the fertility of males and females in laboratory animals, but the relevance of these findings to humans is not clear.

The health risks of trichloroethylene have been studied extensively. The U.S. Environmental Protection Agency (EPA) sponsored a “state of the science” review of the health effects associated with exposure to trichloroethylene. Based on this review, the EPA published a risk assessment that concluded trichloroethylene posed a more significant human health risk than previous studies had indicated.EPA’s report provoked considerable debate about the quality of evidence describing the health risks of trichloroethylene, and the methods used to assess that evidence.[citation needed] In 2004, an interagency group composed of the EPA, Department of Defense, Department of Energy, and the National Aeronautics and Space Administration requested the National Academy of Sciences (NAS) to provide independent guidance on the scientific issues related regarding trichloroethylene health risks.[citation needed] The NAS report concluded that evidence on the carcinogenic risk and other potential health hazards from exposure to TCE has strengthened since EPA released their toxicological assessment of TCE, and encourages federal agencies to finalize the risk assessment for TCE using currently available information, so that risk management decisions for this chemical can be expedited.

Human exposure

Human exposure to trichloroethylene is potentially widespread.It is a common contaminant in soil and groundwater at hundreds of waste sites across the United States.[citation needed] Some are exposed to trichloroethylene through contaminated drinking water. Others are potentially exposed through inhalation of vapor from contaminated soil or groundwater entering nearby buildings.Tens of thousands of workers are potentially exposed to trichloroethylene used as a degreasing and cleaning chemical.Other exposures have occurred through the long-term use of trichloroethylene as a surgical anesthetic.

TCE was first detected in groundwater in 1977, and is one of the most frequently detected contaminants in groundwater in the U.S. Up to 34 percent of the drinking water supply sources tested in the U.S. may have some TCE contamination, though EPA has reported that most water supplies are in compliance with the Maximum Contaminant Level (MCL) of 5 ug/L.[citation needed] In addition, a growing concern in recent years at sites with TCE contamination in soil or groundwater has been vapor intrusion in buildings, which has resulted in indoor air exposures.[citation needed] Trichloroethylene has been detected in 852 Superfund sites across the United States, according to the Agency for Toxic Substances and Disease Registry (ATSDR).

Existing regulation

Until recent years, the US Agency for Toxic Substances and Disease Registry (ATSDR) contended that trichloroethylene had little-to-no carcinogenic potential, and was probably a co-carcinogen—that is, it acted in concert with other substances to promote the formation of tumors.

Half a dozen state, federal, and international agencies now classify trichloroethylene as a probable carcinogen. The International Agency for Research on Cancer considers trichloroethylene a Group 2A carcinogen, indicating that it considers it is probably carcinogenic to humans. California EPA regulators consider it a known carcinogen and issued a risk assessment in 1999 that concluded that it was far more toxic than previous scientific studies had shown.

Proposed U.S. federal regulation

In 2001, a draft report of the Environmental Protection Agency (EPA) laid the groundwork for tough new standards to limit public exposure to trichloroethylene. The assessment set off a fight between the EPA and the Department of Defense (DoD), the Department of Energy, and NASA, who appealed directly to the White House. They argued that the EPA had produced junk science, its assumptions were badly flawed, and that evidence exonerating the chemical was ignored.

The DoD has about 1,400 military properties nationwide that are polluted with trichloroethylene. The chemical has contaminated 23 sites in the Energy Department’s nuclear weapons complex — including Lawrence Livermore National Laboratory in the San Francisco Bay area, and NASA centers, including the Jet Propulsion Laboratory in La Cañada Flintridge.

High-level political appointees in the EPA — notably research director Paul Gilman — sided with the Pentagon and agreed to pull back the risk assessment. In 2004, the National Academy of Sciences was given a a $680,000 contract to study the matter, releasing its report in the summer of 2006. The report has raised greater concern about the adverse health effects of TCE, opening up the debate for better regulation.

Reduced production and remediation

In recent times, there has been a substantial reduction in the production output of trichloroethylene; alternatives for use in metal degreasing abound, chlorinated aliphatic hydrocarbons being phased out in a large majority of industries due to the potential for irreversible health effects and the legal liability that ensues as a result.

The U.S. military has virtually eliminated its use of the chemical, purchasing only 11 gallons in 2005. About 100 tons of it is used annually in the U.S. as of 2006.

Recent research has focused on aerobic degradation pathways in order to reduce environmental pollution through the use of genetically modified bacteria. Limited success has been attained thus far; the intended application is for treatment and detoxification of industrial wastewater.

Cases of TCE contaminated water

• Shannon, Quebec and Val-Belair, Quebec
• Woburn, Massachusetts (the town from the book/movie A Civil Action)
• Camp Lejeune, North Carolina
• Lisle, Illinois
• Scottsdale, Arizona (from the Indian Bend Wash Area Superfund site, 2007 and 2008)
• Cambridge, Ontario: http://news.therecord.com/News/Local/article/271823
• Salem, Ohio Nease Chemical, Supplied U.S. Army with chemicals for Agent Orange: http://cfpub.epa.gov/supercpad/SiteProfiles/index.cfm?fuseaction=second.Contams&id=0504619
• Tuscon, Arizona
• Walkersville, Maryland
• Victor, New York

• Salina, Kansas

OXALIC ACID

Oxalic acid is the chemical compound with the formula H2C2O4. This dicarboxylic acid is better described with the formula HOOCCOOH. It is a relatively strong organic acid, being about 10,000 times stronger than acetic acid. The di-anion, known as oxalate, is also a reducing agent as well as a ligand in coordination chemistry. Many metal ions form insoluble precipitates with oxalate, a prominent example being calcium oxalate, which is the primary constituent of the most common kind of kidney stone.

Preparation

Although it can be readily purchased, oxalic acid can be prepared in the laboratory by oxidizing sucrose using nitric acid in the presence of a small amount of vanadium pentoxide as a catalyst. On a large scale, sodium oxalate is manufactured by absorbing carbon monoxide under pressure in hot sodium hydroxide.

Typically oxalic acid is obtained as the dihydrate. This solid can be dehydrated with heat or by azeotropic distillation. Anhydrous oxalic acid exists as two polymorphs; in one the hydrogen-bonding results in a chain-like structure whereas the hydrogen bonding pattern in the other form defines a sheet-like structure.

Reactions

Oxalic acid is a relatively strong weak acid with pKa1=1.27 and pKa2=4.28. Oxalic acid exhibits many of the reactions characteristic of other carboxylic acids. It forms esters such as dimethyloxalate (m.p. 52.5–53.5 °C). It forms an acid chloride called oxalyl chloride.

Oxalate, the conjugate base of oxalic acid, is an excellent ligand for metal ions. It usually binds as a bidentate ligand forming a 5-membered MO2C2 ring. An illustrative complex is potassium ferrioxalate, K3[Fe(C2O4)3]. The drug Oxaliplatin exhibits improved water solubility relative to older platinum-based drugs, avoiding the dose-limiting side-effect of nephrotoxicity.

Occurrence in nature

Oxalic acid and oxalates are abundantly present in many plants, most notably fat hen (lamb’s quarters), sorrel, and Oxalis species. The root and/or leaves of rhubarb and buckwheat are listed being high in oxalic acid.

Other edible plants that contain significant concentrations of oxalic acid include—in decreasing order—star fruit (carambola), black pepper, parsley, poppy seed, amaranth, spinach, chard, beets, cocoa, chocolate, most nuts, most berries, and beans.

The gritty “mouth feel” one experiences when drinking milk with a rhubarb dessert is caused by precipitation of calcium oxalate. Thus even dilute amounts of oxalic acid can readily “crack” the casein found in various dairy products.

Leaves of the tea plant (Camellia sinensis) contain among the greatest measured concentrations of oxalic acid relative to other plants. However the infusion beverage typically contains only low to moderate amounts of oxalic acid per serving, due to the small mass of leaves used for brewing.

Physiological effects
Chemical structure of Oxalic acid.

The affinity of divalent metal ions is sometimes reflected in their tendency to form insoluble precipitates. Thus in the body, oxalic acid also combines with metals ions such as Ca2+, Fe2+, and Mg2+ to deposit crystals of the corresponding oxalates, which irritate the gut and kidneys. Because it binds vital nutrients such as calcium, long-term consumption of foods high in oxalic acid can be problematic. Healthy individuals can safely consume such foods in moderation, but those with kidney disorders, gout, rheumatoid arthritis, or certain forms of chronic vulvar pain (vulvodynia) are typically advised to avoid foods high in oxalic acid or oxalates. Conversely, calcium supplements taken along with foods high in oxalic acid can cause calcium oxalate to precipitate out in the gut and drastically reduce the levels of oxalate absorbed by the body (by 97% in some cases.) The calcium oxalate precipitate (better known as kidney stones) obstruct the kidney tubules.

Oxalic acid can also be produced by the metabolism of ethylene glycol (“antifreeze”), glyoxylic acid or ascorbic acid (vitamin C). Under certain conditions of concentration and pH, oxalic acid can precipitate in the kidneys as calcium oxalate crystals, forming an estimated 80% of kidney stones.

Some Aspergillus species produce oxalic acid, which reacts with blood or tissue calcium to precipitate calcium oxalate. There is some preliminary evidence that the administration of probiotics can affect oxalic acid excretion rates (and presumably oxalic acid levels as well.)

Methods to reduce the oxalate content in food are of current interest.

Other uses

• In household chemical products such as Bar Keeper’s Friend, some bleaches, and rustproofing treatments.
• In wood restorers where the acid dissolves away a layer of dry surface wood to expose fresh material underneath.
• As an additive to automotive wheel cleaners.
• As a mordant in dyeing processes.
• Vaporized oxalic acid, or a 6% solution of oxalic acid in sugar syrup, is used by some beekeepers as an insecticide against the parasitic Varroa mite.
• As a rust remover in such applications as automotive shops and for the restoration of antiques.
• As a recommended surface pretreatment for stainless steels (surface etch) before application of solid metal or polymer self-lubricating coatings.
• For polishing stones and marble.

• Used in the acid treatment for destroying warts.

Tests for oxalic acid

Titration with potassium permanganate can reveal the presence of oxalic acid. Ascorbate interferes with this test which is based on reducing power. For this reason, a second test for strong reductants using, for example, iodine can be done.

sodium sulfate

Sodium sulfate is the sodium salt of sulfuric acid. Anhydrous, it is a white crystalline solid of formula Na2SO4; the decahydrate Na2SO4·10H2O has been known as Glauber’s salt or, historically, sal mirabilis since the 17th century. With an annual production of 6 million tonnes, it is one of the world’s major commodity chemicals.

Sodium sulfate is mainly used for the manufacture of detergents and in the Kraft process of paper pulping. About two thirds of the world’s production is from mirabilite, the natural mineral form of the decahydrate, and the remainder from by-products of chemical processes such as hydrochloric acid production.

History

The hydrate of sodium sulfate is known as Glauber’s Salt after the Dutch/German apothecary Johann Rudolf Glauber (1604–1670), who discovered it in Hungarian spring water. He himself named it sal mirabilis (miraculous salt), because of its medicinal properties: the crystals were used as a general purpose laxative, until more sophisticated alternatives came about in the 1900s.[1][2]

In the 18th century, Glauber’s salt began to be used as a raw material for the industrial production of soda ash (sodium carbonate), by reaction with potash (potassium carbonate). Requirement for soda ash increased and supply of sodium sulfate had to increase in line. Therefore, in the nineteenth century, the Leblanc process, producing synthetic sodium sulfate as a key intermediate, became the principal method of soda ash production.

Physical and chemical properties

Sodium sulfate is chemically very stable, being unreactive toward most oxidising or reducing agents at normal temperatures. At high temperatures, it can be reduced to sodium sulfide.[4] It is a neutral salt, which forms aqueous solutions with pH of 7. The neutrality of such solutions reflects the fact that Na2SO4 is derived, formally speaking, from the strong acid sulfuric acid and a strong base sodium hydroxide. Sodium sulfate reacts with an equivalent amount of sulfuric acid to give an equilibrium concentration of the acid salt sodium bisulfate[5][6]:
Na2SO4(aq) + H2SO4(aq) ⇌ 2 NaHSO4(aq)

In fact, the equilibrium is very complex, depending on concentration and temperature, with other acid salts being present.

Sodium sulfate is a typical ionic sulfate, containing Na+ ions and SO42− ions. Aqueous solutions can produce precipitates when combined with salts of Ba2+ or Pb2+, which form insoluble sulfates
Na2SO4(aq) + BaCl2(aq) → 2 NaCl(aq) + BaSO4(s)

Sodium sulfate has unusual solubility characteristics in water.[7] Its solubility rises more than tenfold between 0 °C to 32.4 °C, where it reaches a maximum of 49.7 g Na2SO4 per 100 g water. At this point the solubility curve changes slope, and the solubility becomes almost independent of temperature. In the presence of NaCl, the solubility of sodium sulfate is markedly diminished. Such changes provide the basis for the use of sodium sulfate in passive solar heating systems, as well is in the preparation and purification of sodium sulfate. This nonconformity can be explained in terms of hydration, since 32.4 °C corresponds with the temperature at which the crystalline decahydrate (Glauber’s salt) changes to give a sulfate liquid phase and an anhydrous solid phase.

Sodium sulfate decahydrate is also unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 J·K-1·mol-1. This is ascribed to its ability to distribute water much more rapidly compared to most hydrates.[8]

Sodium sulfate displays a moderate tendency to form double salts. The only alums formed with common trivalent metals are NaAl(SO4)2 (unstable above 39 °C) and NaCr(SO4)2, in contrast to potassium sulfate and ammonium sulfate which form many stable alums.[9] Double salts with some other alkali metal sulfates are known, including Na2SO4.3K2SO4 which occurs naturally as the mineral glaserite. Formation of glaserite by reaction of sodium sulfate with potassium chloride has been used as the basis of a method for producing potassium sulfate, a fertiliser.[10] Other double salts include 3Na2SO4.CaSO4, 3Na2SO4.MgSO4 (vanthoffite) and NaF.Na2SO4.

Production

The world production of sodium sulfate, mostly in the form of the decahydrate amounts to approximately 5.5 to 6 million tonnes annually (Mt/a). In 1985, production was 4.5 Mt/a, half from natural sources, and half from chemical production. After 2000, at a stable level until 2006, natural production had increased to 4 Mt/a, and chemical production decreased to 1.5 to 2 Mt/a, with a total of 5.5 to 6 Mt/a.[12][13][14][15] For all applications, naturally produced and chemically produced sodium sulfate are practically interchangeable.

Natural sources

Two thirds of the world’s production of the decahydrate (Glauber’s salt) is from the natural mineral form mirabilite, for example as found in lake beds in southern Saskatchewan. In 1990, Mexico and Spain were the world’s main producers of natural sodium sulfate (each around 500,000 tonnes), with Russia, USA and Canada around 350,000 tonnes each.[13] Estimatedly, natural resources amount to over 1 billion tonnes.[12][13]

Major producers of 200–1500 Mt/a in 2006 include Searles Valley Minerals (California, USA), Airborne Industrial Minerals (Saskatchewan, Canada), Química del Rey (Coahuila, Mexico), Criaderos Minerales Y Derivados and Minera de Santa Marta, also known as Grupo Crimidesa (Burgos, Spain), FMC Foret (Toledo, Spain), Sulquisa (Madrid, Spain), and in China Chengdu Sanlian Tianquan Chemical (Sichuan), Hongze Yinzhu Chemical Group (Jiangsu), Nafine Chemical Industry Group (Shanxi), and Sichuan Province Chuanmei Mirabilite (Sichuan), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).[12][14]

Anhydrous sodium sulfate occurs in arid environments as the mineral thenardite. It slowly turns to mirabilite in damp air. Sodium sulfate is also found as glauberite, a calcium sodium sulfate mineral. Both minerals are less common than mirabilite.

Chemical industry

About one third of the world’s sodium sulfate is produced as by-product of other processes in chemical industry. Most of this production is chemically inherent to the primary process, and only marginally economical. By effort of the industry, therefore, sodium sulfate production as by-product is declining.

The most important chemical sodium sulfate production is during hydrochloric acid production, either from sodium chloride (salt) and sulfuric acid, in the Mannheim process, or from sulfur dioxide in the Hargreaves process.[16][17] The resulting sodium sulfate from these processes are known as salt cake.
Mannheim: 2 NaCl + H2SO4 → 2 HCl + Na2SO4
Hargreaves: 4 NaCl + 2 SO2 + O2 + 2 H2O → 4 HCl + 2 Na2SO4

The second major production of sodium sulfate are the processes where surplus sulfuric acid is neutralised by sodium hydroxide, as applied on a large scale in the production of rayon. This method is also a regularly applied and convenient laboratory preparation.
2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l)

Formerly, sodium sulfate was also a by-product of the manufacture of sodium dichromate, where sulfuric acid is added to sodium chromate solution forming sodium dichromate, or subsequently chromic acid. Alternatively, sodium sulfate is or was formed in the production of lithium carbonate, chelating agents, resorcinol, ascorbic acid, silica pigments, nitric acid, and phenol.[12]

Bulk sodium sulfate is usually purified via the decahydrate form, since the anhydrous form tends to attract iron compounds and organic compounds. The anhydrous form is easily produced from the hydrated form by gentle warming.

Major sodium sulfate by-product producers of 50–80 Mt/a in 2006 include Elementis Chromium (chromium industry, Castle Hayne, NC, USA), Lenzing AG (200 Mt/a, rayon industry, Lenzing, Austria), Addiseo (formerly Rhodia, methionine industry, Les Roches-Roussillon, France), Elementis (chromium industry, Stockton-on-Tees, UK), Shikoku Chemicals (Tokushima, Japan) and Visko-R (rayon industry, Russia).[12]

Applications
Commodity industries

With USA pricing at $30 per tonne in 1970, in 2006 up to $90 per tonne for salt cake quality and $130 for better grades, sodium sulfate is a very cheap material. The largest use is as filler in powdered home laundry detergents, consuming approx. 50% of world production. This use is waning as domestic consumers are increasingly switching to compact or liquid detergents that do not include sodium sulfate.[12]

Another formerly major use for sodium sulfate, notably in the USA and Canada, is in the Kraft process for the manufacture of wood pulp. Organics present in the “black liquor” from this process are burnt to produce heat, needed to drive the reduction of sodium sulfate to sodium sulfide. However, this process is being replaced by newer processes; use of sodium sulfate in the USA and Canadian pulp industry declined from 1.4 Mt/a in 1970 to only approx. 150,000 tonnes in 2006.[12]

The glass industry provides another significant application for sodium sulfate, as second largest application in Europe. Sodium sulfate is used as a fining agent, to help remove small air bubbles from molten glass. It fluxes the glass, and prevents scum formation of the glass melt during refining. The glass industry in Europe has been consuming from 1970 to 2006 a stable 110,000 tonnes annually.[12]

Sodium sulfate is important in the manufacture of textiles, particularly in Japan, where it is the largest application. Sodium sulfate helps in “levelling”, reducing negative charges on fibres so that dyes can penetrate evenly. Unlike the alternative sodium chloride, it does not corrode the stainless steel vessels used in dyeing. This application in Japan and USA consumed in 2006 approximately 100,000 tonnes.

Thermal storage

The high heat storage capacity in the phase change from solid to liquid, and the advantageous phase change temperature of 32 degrees Celsius (90 degrees Fahrenheit) makes this material especially appropriate for storing low grade solar heat for later release in space heating applications. In some application the material is incorporated into thermal tiles that are placed in an attic space while in other applications the salt is incorporated into cells surrounded by solar–heated water. The phase change allows a substantial reduction in the mass of the material required for effective heat storage (83 calories per gram stored across the phase change, versus one calorie per gram per degree Celsius using only water), with the further advantage of a consistency of temperature as long as sufficient material in the appropriate phase is available.

Small-scale applications

In the laboratory, anhydrous sodium sulfate is widely used as an inert drying agent, for removing traces of water from organic solutions.[18] It is more efficient, but slower-acting, than the similar agent magnesium sulfate. It is only effective below about 30 °C, but it can used with a variety of materials since it is chemically fairly inert. Sodium sulfate is added to the solution until the crystals no longer clump together; the two video clips (see above) demonstrate how the crystals clump when still wet, but some crystals flow freely once a sample is dry.

Glauber’s salt, the decahydrate, was historically used as a laxative. It is effective for the removal of certain drugs such as acetaminophen from the body, for example, after an overdose.[19][20]

In 1953, sodium sulfate was proposed for heat storage in passive solar heating systems. This takes advantage of its unusual solubility properties, and the high heat of crystallisation (78.2 kJ/mol).[21]

Other uses for sodium sulfate include de-frosting windows, in carpet fresheners, starch manufacture, and as an additive to cattle feed.

Lately, sodium sulfate has been found effective in dissolving very finely electroplated micrometre gold that is found in gold electroplated hardware on electronic products such as pins, and other connectors and switches. It is safer and cheaper than other reagents used for gold recovery, with little concern for adverse reactions or health effects.[citation needed]

At least one company makes a laptop computer chill mat using sodium sulfate decahydrate inside a quilted plastic pad. The material slowly turns to liquid as the heat from the laptop is transferred

Safety

Although sodium sulfate is generally regarded as non-toxic,[22] it should be handled with care. The dust can cause temporary asthma or eye irritation; this risk can be prevented by using eye protection and a paper mask. Transport is not limited, and no Risk Phrase or Safety Phrase apply

Sodium Benzoate

Sodium benzoate (E211), also called benzoate of soda, has chemical formula NaC6H5CO2. It is the sodium salt of benzoic acid and exists in this form when dissolved in water. It can be produced by reacting sodium hydroxide with benzoic acid.

Uses

Sodium benzoate is a preservative. It is bacteriostatic and fungistatic under acidic conditions. it is used most prevalently in acidic foods such as salad dressings (vinegar), carbonated drinks (carbonic acid), jams and fruit juices (citric acid), pickles (vinegar), and Chinese food sauces (soy, mustard, and duck).[citation needed] It is also found in alcohol-based mouthwash and silver polish. Sodium benzoate is declared on a product label as ‘sodium benzoate’ or E211. The taste of sodium benzoate cannot be detected by around 25 percent of the population, but for those who can taste the chemical, it tends to be perceived as sweet, sour, salty, or sometimes bitter.

It is also used in fireworks as a fuel in whistle mix, a powder which imparts a whistling noise when compressed into a tube and ignited.

It is found naturally in cranberries, prunes, greengage plums, cinnamon, ripe cloves, and apples. Concentration as a preservative is limited by the FDA in the U.S. to 0.1% by weight though organically-grown cranberries and prunes can conceivably contain levels exceeding this limit. The International Programme on Chemical Safety found no adverse effects in humans at doses of 647-825 mg/kg of body weight per day.

Cats have a significantly lower tolerance against benzoic acid and its salts than rats and mice. Sodium benzoate is, however, allowed as an animal food additive at up to 0.1%, according to AFCO’s official publication.

Mechanism of food preservation

The mechanism starts with the absorption of benzoic acid into the cell. If the intracellular pH changes to 5 or lower, the anaerobic fermentation of glucose through phosphofructokinase is decreased by 95%.

Safety and health
Main article: benzene in soft drinks

In combination with ascorbic acid (vitamin C, E300), sodium benzoate and potassium benzoate may form benzene, a known carcinogen. Heat, light and shelf life can affect the rate at which benzene is formed.

Professor Peter Piper of the University of Sheffield claims that sodium benzoate by itself can damage and inactivate vital parts of DNA in a cell’s mitochondria. “The mitochondria consumes the oxygen to give you energy and if you damage it – as happens in a number of diseased states – then the cell starts to malfunction very seriously. And there is a whole array of diseases that are now being tied to damage to this DNA – Parkinson’s and quite a lot of neuro-degenerative diseases, but above all the whole process of aging.

ADHD

Research published in 2007 for the UK’s Food Standards Agency suggests that sodium benzoate (E211) is linked to hyperactive behaviour and decreased intelligence in children. According to the report, a high consumption of sodium benzoate is associated with a reduction in IQ of close to 5.5 points. On 6 September 2007, the British Food Standards Agency issued revised advice on certain artificial food additives, including sodium benzoate (E211).

Professor Jim Stevenson from Southampton University, and author of the report, said: “This has been a major study investigating an important area of research. The results suggest that consumption of certain mixtures of artificial food colours and sodium benzoate preservative are associated with increases in hyperactive behaviour in children.

“However, parents should not think that simply taking these additives out of food will prevent hyperactive disorders. We know that many other influences are at work but this at least is one a child can avoid.”

Two mixtures of additives were tested in the research:

Mix A:
Sunset yellow (E110)
Tartrazine (E102)
Carmoisine (E122)
Ponceau 4R (E124)
Sodium benzoate (E211)

Mix B:
Sunset yellow (E110)
Quinoline yellow (E104)
Carmoisine (E122)
Allura red (E129)
Sodium benzoate (E211)

Sodium benzoate was included in both mixes, but the effects observed were not consistent. The Food Standards Agency therefore considers that, if real, the observed increases in hyperactive behaviour were more likely to be linked to one or more of the specific colours tested.

On 10 April 2008, the Foods Standard Agency called for a voluntary removal of the colours (but not sodium benzoate) by 2009. In addition, it recommended that there should be action to phase them out in food and drink in the European Union (EU) over a specified period.

SILICONE OIL

Silicone oils (polymerized siloxanes) are silicon analogues of carbon based organic compounds, and can form (relatively) long and complex molecules based on silicon rather than carbon. Chains are formed of alternating silicon-oxygen atoms (…Si-O-Si-O-Si…) or siloxane, rather than carbon atoms (…C-C-C-C…). Other species attach to the tetravalent silicon atoms, not to the divalent oxygen atoms which are fully committed to forming the siloxane chain. A typical example is polydimethylsiloxane, where two methyl groups attach to each silicon atom to form (H3C)[SiO(CH3)2]nSi(CH3). The carbon analogue would be an alkane, e.g. dimethylpropane C5H12 or (H3C)[C(CH3)2](CH3)

Silicone oils could be a basis for silicon-based organic life, but their more prosaic, primary uses are as lubricants or hydraulic fluids. They are excellent electrical insulators and, unlike their carbon analogues, are non flammable. Their temperature-stability and good heat-transfer characteristics make them widely used in laboratories for heating baths (“oil baths”) placed on top of hotplate stirrers. Silicone oil is also commonly used as working fluid in diffusion pumps

Some silicone oils such as simethicone are potent anti-foaming agents. They are used in industrial applications such as distillation or fermentation where excessive amounts of foam can be problematic. They are sometimes added to cooking oils to prevent excessive frothing during deep frying. Consumer products to control flatus (antiflatulents) often contain silicone oil. Silicone oils used as lubricants can be inadvertent defoamers (contaminants) in processes where foam is desired, such as in the manufacture of polyurethane foam. Silicone oils have been used as a vitreous fluid substitute to treat difficult case of retinal detachment, such as those complicated with proliferative vitreoretinopathy, giant retinal tears, and penetrating ocular trauma [1]. Silicone oil is also one of the two main ingredients in Silly Putty, along with borax. Silicone oil plays a useful role in gas powered airsoft guns where it is used to lubricate the rubber gas seals in gas blowback guns without degrading them as carbon based oils would as well as lubricating the moving parts of the guns.

TOLUENE

Toluene, also known as methylbenzene or phenylmethane, is a clear, water-insoluble liquid with the typical smell of paint thinners, redolent of the sweet smell of the related compound benzene. It is an aromatic hydrocarbon that is widely used as an industrial feedstock and as a solvent. Like other solvents, toluene is also used as an inhalant drug for its intoxicating properties.

History

The name toluene was derived from the older name toluol, which refers to tolu balsam, an aromatic extract from the tropical Colombian tree Myroxylon balsamum, from which it was first isolated. It was originally named by Jöns Jakob Berzelius.

Chemical properties

Toluene reacts as a normal aromatic hydrocarbon towards electrophilic aromatic substitution. The methyl group makes it around 25 times more reactive than benzene in such reactions. It undergoes smooth sulfonation to give p-toluenesulfonic acid, and chlorination by Cl2 in the presence of FeCl3 to give ortho and para isomers of chlorotoluene. It undergoes nitration to give ortho and para nitrotoluene isomers, but if heated it can give dinitrotoluene and ultimately the explosive trinitrotoluene (TNT).

With other reagents the methyl side chain in toluene may react, undergoing oxidation. Reaction with potassium permanganate leads to benzoic acid, whereas reaction with chromyl chloride leads to benzaldehyde (Étard reaction). Halogenation can be performed under free radical conditions. For example, N-bromosuccinimide (NBS) heated with toluene in the presence of AIBN leads to benzyl bromide.

Catalytic hydrogenation of toluene to methylcyclohexane requires a high pressure of hydrogen to go to completion, because of the stability of the aromatic system. pka is approximately 45.

Preparation

Toluene occurs naturally at low levels in crude oil and is usually produced in the processes of making gasoline via a catalytic reformer, in an ethylene cracker or making coke from coal. Final separation (either via distillation or solvent extraction) takes place in a BTX plant.

Uses

Toluene is a common solvent, able to dissolve: paints, paint thinners, silicone sealants, many chemical reactants, rubber, printing ink, adhesives (glues), lacquers, leather tanners, and disinfectants. It can also be used as a fullerene indicator, and is a raw material for toluene diisocyanate (used in the manufacture of polyurethane foam) and TNT. Industrial uses of toluene include dealkylation to benzene and disproportionation to a mixture of benzene and xylene. When oxidized it yields benzaldehyde and benzoic acid, two important intermediates in chemistry. It is also used as a carbon source for making Multi-Wall Carbon Nanotubes. Toluene can be used to break open red blood cells in order to extract hemoglobin in biochemistry experiments.

Toluene can be used as an octane booster in gasoline fuels used in internal combustion engines. Toluene at 86% by volume fueled all the turbo Formula 1 teams in the 1980s.

Toxicology and metabolism
Main article: Toluene (toxicology)

Inhalation of toluene fumes can be intoxicating, but in larger doses nausea-inducing. Toluene may enter the human system not only through vapour inhalation from the liquid evaporation, but also following soil contamination events, where human contact with soil, ingestion of contaminated groundwater or soil vapour off-gassing can occur.

The toxicity of toluene can be explained mostly by its metabolism. As toluene has very low water solubility, it cannot exit the body via the normal routes (urine, feces, or sweat).[citation needed] It must be metabolized in order to be excreted. The methyl group of toluene is more easily oxidized by cytochrome P450 than the benzene ring. Therefore, in the metabolism of toluene, 95% is oxidized to become benzyl alcohol. The toxic metabolites are created by the remaining 5% that are oxidized to benzaldehyde and cresols. Most of the reactive products are detoxified by conjugation to glutathione but the remainder may severely damage cells.

HYDROFLUORIC ACID

Hydrofluoric acid is a solution of hydrogen fluoride in water. Together with hydrogen fluoride, hydrofluoric acid is a valued source of fluorine, being the precursor to numerous pharmaceuticals, diverse polymers (e.g. Teflon), and most other synthetic materials that contain fluorine. Hydrofluoric acid is best known to the public for its ability to dissolve glass by reacting with SiO2, the major component of most glasses. This dissolution process can be described as follows:
SiO2(s) + 4HF(aq) → SiF4(g) + 2H2O(l)
SiO2(s) + 6HF(aq) → H2[SiF6](aq) + 2H2O(l)

Because of its high reactivity toward glass, hydrofluoric acid is typically stored in polyethylene or Teflon containers. It is also unique in its ability to dissolve many metal and semimetal oxides. It is corrosive, as explained below.

Acidity

Hydrogen fluoride dissociates in aqueous solution in a similar fashion to other common acids:
HF + H2O → H3O+ + F−

When the concentration of HF approaches 100%, the acidity increases dramatically due to the following equilibrium:
2HF → H+ + FHF−

The FHF− anion is stabilized by the very strong hydrogen – fluorine hydrogen bond. Hydrofluoric acid is the only one of the hydrohalic acids that is not considered a strong acid due to its lack of ionization in aqueous solution.

Production
Main article: hydrogen fluoride

Industrially, hydrofluoric acid is produced by treatment of the mineral fluorite (CaF2) with concentrated sulfuric acid. When combined at 250 °C, these two substances react to produce hydrogen fluoride according to the following chemical equation:
CaF2 + H2SO4 → 2HF + CaSO4

Uses

Because of its ability to dissolve metal oxides, hydrofluoric acid is used in the purification of both aluminium and uranium. It is also used to etch glass, to remove surface oxides from silicon in the semiconductor industry, as a catalyst for the alkylation of isobutane and butene (olefinic C4) in oil refineries, and to remove oxide impurities from stainless steel in a process called pickling. Dilute hydrofluoric acid is sold as a household rust stain remover. Recently it has even been used in car washes in “wheel cleaner” compounds. Due to its ability to dissolve silicate compounds, hydrofluoric acid is often used to dissolve rock samples (usually powdered) prior to analysis.

Hydrofluoric acid is also used in the synthesis of many fluorine-containing organic compounds, including Teflon, fluoropolymers, perfluorocarbons, and refrigerants such as freon. Additionally, hydrofluoric acid is commonly used in refinery alkylation processes to produce a high-octane gasoline blending component called alkylate from FCCU C3 and C4 olefins and isobutane.

Affinity for magnesium and calcium

Hydrofluoric acid attacks many metal oxides, forming the corresponding fluoro derivatives. In the body, hydrofluoric acid reacts with the ubiquitous biologically important ions Ca2+ and Mg2+. In some cases, exposures can lead to hypocalcemia. Thus, hydrofluoric acid exposure is often treated with calcium gluconate, a source of Ca2+ that sequesters the fluoride ions.

ydrofluoric acid is corrosive and a contact poison. It should be handled with extreme care, beyond that accorded to other mineral acids, in part because of its low dissociation constant, which allows HF to penetrate tissue more quickly. Symptoms of exposure to hydrofluoric acid may not be immediately evident. HF interferes with nerve function and burns may not initially be painful. Accidental exposures can go unnoticed, delaying treatment and increasing the extent and seriousness of the injury. HF is known to etch bone, and since it penetrates the skin it essentially breaks the person’s bones without destroying the skin. Hydrogen fluoride is released upon combustion of fluorine-containing compounds such as products containing Viton and Teflon parts. Hydrogen fluoride converts immediately to hydrofluoric acid upon contact with moisture.

HF chemical burns can be treated with a water wash and 2.5% calcium gluconate gel or special rinsing solutions.

NACL

Sodium chloride, also known as common salt, table salt, or halite. It is a chemical 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. In one gram of sodium chloride, there are approximately 0.3933 grams of sodium, and 0.6067 grams of chlorine.

Production and use

Salt 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 2002, world production was estimated at 210 million metric tonnes, the top five producers being the United States (40.3 million tonnes), China (32.9), Germany (17.7), India (14.5), and Canada (12.3).[1]

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 and detergents. In cold countries, large quantities of rock salt are used to help clear highways of ice during winter, although “Road Salt” loses its melting ability at temperatures below -15°C to -20°C (5°F to -4°F). Sodium chloride is sometimes used as a cheap and safe desiccant due to its hygroscopic properties, making salting an effective method of food preservation historically. Even though more effective desiccants are available, few are safe for humans to ingest.

Synthetic uses

Salt is also the raw material used to produce chlorine which itself is required for the production of many modern materials including PVC and pesticides. 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 equation
2NaCl + 2H2O → Cl2 + H2 + 2NaOH

Sodium metal is produced commercially through the electrolysis of liquid sodium chloride. This is 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.

Biological uses

Many microorganisms 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 and can also be used to detach leeches that have attached themselves to feed. It has also been used to disinfect wounds. In medieval times salt would be rubbed into household surfaces as a cleansing agent.

Biological functions

In humans, a high-salt intake 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

Crystal structure
Sodium chloride forms crystals with 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 blue spheres, fill the octahedral gaps between them.

Each ion is surrounded by six ions of the other kind. This same basic structure is found in many other minerals, and is known as the halite structure. This arrangement is known as cubic close packed (ccp). It can be represented as two interpenetrating face-centered cubic (fcc) lattices, or one fcc lattice with a two atom basis. It is most commonly known as the rocksalt crystal structure.

It is held together with an ionic bond and electrostatic force

Road salt

While salt was once a scarce commodity in history, industrialized production has now made salt plentiful. About 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. This works because salt and water form an eutectic mixture. Adding salt to water will lower the freezing temperature of the water, depending on the concentration. The salinity (S) of water is measured as grams salt per kilogram (1000g) water, and the freezing temperatures are as follows.

Additives

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.[2] 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.[3]

Salt for de-icing in the UK typically contains sodium hexacyanoferrate (II) at less than 100ppm as an anti-caking agent. In recent years this additive has also been used in table salt.

Common chemicals

Chemicals used in de-icing salts are mostly found to be sodium chloride (NaCl) or calcium chloride (CaCl2). Both are similar and are effective in de-icing roads. When these chemicals are produced, they are mined/made, crushed to fine granules, then treated with an anti-caking agent. Adding salt lowers the freezing point of the water, which allows the liquid to be stable at lower temperatures and allows the ice to melt. Alternative de-icing chemicals have also been used. Chemicals such as calcium magnesium acetate and potassium formate are being produced. These chemicals have few of the negative chemical effects on the environment commonly associated with NaCl and CaCl2.