Polyacrylamide

Polyacrylamide

Polyacrylamide (IUPAC poly(2-propenamide) or poly(1-carbamoylethylene), abbreviated as PAM) is a polymer (-CH₂CHCONH₂-) formed from acrylamide subunits. It can be synthesized as a simple linear-chain structure or cross-linked, typically using N,N‍ '​-methylenebisacrylamide. In the cross-linked form, the possibility of the monomer being present is reduced even further. It is highly water-absorbent, forming a soft gel when hydrated, used in such applications as polyacrylamide gel electrophoresis and in manufacturing soft contact lenses. In the straight-chain form, it is also used as a thickener and suspending agent. More recently, it has been used as a subdermal filler for aesthetic facial surgery (see Aquamid).

1 - Uses of polyacrylamide

One of the largest uses for polyacrylamide is to flocculate solids in a liquid. This process applies to water treatment, and processes like paper making, Screen Printing. Polyacrylamide can be supplied in a powder or liquid form, with the liquid form being subcategorized as solution and emulsion polymer. Even though these products are often called 'polyacrylamide', many are actually copolymers of acrylamide and one or more other chemical species, such as an acrylic acid or a salt thereof. The main consequence of this is to give the 'modified' polymer a particular ionic character.

Another common use of polyacrylamide and its derivatives is in subsurface applications such as Enhanced Oil Recovery. High viscosity aqueous solutions can be generated with low concentrations of polyacrylamide polymers, and these can be injected to improve the economics of conventional waterflooding.

It has also been advertised as a soil conditioner called Krilium by Monsanto Company in the 1950s and today "MP", which is stated to be a "unique formulation of PAM (water-soluble polyacrylamide)". It is often used for horticultural and agricultural use under trade names such as Broadleaf P4, Swell-Gel and so on. The anionic form of cross-linked polyacrylamide is frequently used as a soil conditioner on farm land and construction sites for erosion control, in order to protect the water quality of nearby rivers and streams.

The polymer is also used to make Gro-Beast toys, which expand when placed in water, such as the Test Tube Aliens. Similarly, the absorbent properties of one of its copolymers can be utilized as an additive in body-powder.

The ionic form of polyacrylamide has found an important role in the potable water treatment industry. Trivalent metal salts like ferric chloride and aluminium chloride are bridged by the long polymer chains of polyacrylamide. This results in significant enhancement of the flocculation rate. This allows water treatment plants to greatly improve the removal of total organic content (TOC) from raw water.

Polyacrylamide is often used in molecular biology applications as a medium for electrophoresis of proteins and nucleic acids in a technique known as PAGE.

It was also used in the synthesis of the first Boger fluid.

- Soil Conditioner

The primary functions of Polyacrylamide Soil Conditioners are to increase soil tilth, aeration, and porosity and reduce compaction, dustiness and water run-off. Secondary functions are to increase plant vigor, color, appearance, rooting depth and emergence of seeds while decreasing water requirements, diseases, erosion and maintenance expenses. FC 2712 is used for this purpose.

2 - Stability

In dilute aqueous solution, such as is commonly used for Enhanced Oil Recovery applications, polyacrylamide polymers are susceptible to chemical, thermal, and mechanical degradation. Chemical degradation occurs when the labile amine moiety hydrolyzes at elevated temperature or pH, resulting in the evolution of ammonia and a remaining carboxyl group. Thus, the degree of anionicity of the molecule increases. Thermal degradation of the vinyl backbone can occur through several possible radical mechanisms, including the autooxidation of small amounts of iron and reactions between oxygen and residual impurities from polymerization at elevated temperature. Mechanical degradation can also be an issue at the high shear rates experienced in the near-wellbore region. However, cross-linked variants of polyacrylamide have shown greater resistance to all of these methods of degradation, and have proved much more stable.

3 - Environmental effects

Concerns have been raised that polyacrylamide used in agriculture may contaminate food with acrylamide, a known neurotoxin. While polyacrylamide itself is relatively non-toxic, it is known that commercially available polyacrylamide contains minute residual amounts of acrylamide remaining from its production, usually less than 0.05% w/w.

Additionally, there are concerns that polyacrylamide may de-polymerise to form acrylamide. In a study conducted in 2003 at the Central Science Laboratory in Sand Hutton, England, polyacrylamide was treated similarly as food during cooking. It was shown that these conditions do not cause polyacrylamide to de-polymerise significantly.

In a study conducted in 1997 at Kansas State University, the effect of environmental conditions on polyacrylamide were tested, and it was shown that degradation of polyacrylamide under certain conditions can cause the release of acrylamide. The experimental design of this study as well as its results and their interpretation have been questioned, and a 1999 study by the Nalco Chemical Company did not replicate the results.

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Aluminium chlorohydrate or Polyaluminium chloride

Aluminium chlorohydrate or Polyaluminium chloride

Aluminium chlorohydrate is a group of specific aluminium salts having the general formula Al_nCl_(3n-m)(OH)_m. It is used in deodorants and antiperspirants and as a coagulant in water purification.

In water purification, this compound is preferred in some cases because of its high charge, which makes it more effective at destabilizing and removing suspended materials than other aluminium salts such as aluminium sulfate, aluminium chloride and various forms of polyaluminium chloride and polyaluminium chlorisulfate, in which the aluminium structure results in a lower net charge than aluminium chlorohydrate. Further, the high degree of neutralization of the HCl results in minimal impact on treated water pH when compared to other aluminium and iron salts.

1 - Structure

Aluminium chlorohydrate is best described as an inorganic polymer and as such is difficult to structurally characterise. However, techniques such as gel permeation chromatography, X-ray crystallography and ²⁷Al-NMR have been used in research by various groups including that of Nazar and Laden to show that the material is based on Al₁₃ units with a Keggin ion structure and that this base unit then undergoes complex transformations to form larger poly-aluminium complexes.

2 - Synthesis

Aluminium chlorohydrate can be commercially manufactured by reacting aluminium with hydrochloric acid. A number of aluminium-containing raw materials can be used, including aluminium metal, alumina trihydrate, aluminium chloride, aluminium sulfate and combinations of these. The products can contain by-product salts, such as sodium/calcium/magnesium chloride or sulfate.

Because of the explosion hazard related to hydrogen produced by the reaction of aluminium metal with hydrochloric acid, the most common industrial practice is to prepare a solution of aluminium chlorohydrate (ACH) by reacting aluminium hydroxide with hydrochloric acid. The ACH product is reacted with aluminium ingots at 100 °C using steam in an open mixing tank. The Al to ACH ratio and the time of reaction allowed determines the polymer form of the PAC n to m ratio.

3 - Uses

Aluminium chlorohydrate is one of the most common active ingredients in commercial antiperspirants. The variation most commonly used in deodorants and antiperspirants is Al₂Cl(OH)₅.

Aluminium chlorohydrate is also used as a flocculant in water and waste water treatment processes to remove dissolved organic matter and colloidal particles present in suspension.

4 - Safety

The Food and Drug Administration considers the use of aluminium chlorohydrate in antiperspirants to be safe and it is permitted in concentrations up to 25%.

- Alzheimer's disease

There have been studies that have found an association between exposure to and long-term use of antiperspirants and Alzheimer's disease, however the studies also have shown that the association is negligible (less than 1%). There is no adequate evidence that exposure to aluminium in antiperspirants leads to progressive dementia and Alzheimer's disease.

Heather M. Snyder, PhD, the senior associate director of medical and scientific relations for the Alzheimer's Association, has stated "There was a lot of research that looked at the link between Alzheimer's and aluminium, and there hasn't been any definitive evidence to suggest there is a link".

- Breast cancer

The International Journal of Fertility and Women's Medicine found no evidence that certain chemicals used in underarm cosmetics increase the risk of breast cancer. Ted S. Gansler, MD, MBA, the director of medical content for the American Cancer Society, stated "There is no convincing evidence that antiperspirant or deodorant use increases cancer risk".

The European Journal of Cancer Prevention stated "underarm shaving with antiperspirant/deodorant use may play a role in breast cancer." The journal Breast Cancer Research proposed a link between breast cancer and the application of cosmetic chemicals in the underarm, including aluminium, with oestrogenic and/or genotoxic properties. Personal care products are potential contributors to the body burden in aluminium and newer evidence has shown that more aluminium is deposited in the outer regions than the inner regions of the breast. But whether differences in the distribution of aluminium are related to higher incidence of tumours in the outer upper region of the breast remains unknown.

5 - Synonyms

Aluminium hydroxychloride

Aluminium chlorhydroxide

Aluminium chloride basic

Aluminium chlorohydrol

Polyaluminium chloride

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Citric acid

Citric acid

Citric acid is a weak organic acid with the formula  C₆H₈O₇. It is a natural preservative which is present in citrus fruits. It is also used to add an acidic or sour taste to foods and drinks. In biochemistry, the conjugate base of citric acid, citrate, is important as an intermediate in the citric acid cycle, which occurs in the metabolism of all aerobic organisms. It consists of 3 carboxyl (R-COOH) groups.

Citric acid is a commodity chemical, and more than a million tons are produced every year by fermentation. It is used mainly as an acidifier, as a flavoring, and as a chelating agent.

1 - Properties

Citric acid crystal under polarized light, enlarged 200x

Concentration     pH
1M                      ≈1.57
0.5M                   ≈1.72
0.1M                   ≈2.08

At room temperature, citric acid is a white hygroscopic crystalline powder. It can exist either in an anhydrous (water-free) form or as a monohydrate. The anhydrous form crystallizes from hot water, while the monohydrate forms when citric acid is crystallized from cold water. The monohydrate can be converted to the anhydrous form by heating above 78 °C. Citric acid also dissolves in absolute (anhydrous) ethanol (76 parts of citric acid per 100 parts of ethanol) at 15 °C.

In chemical structure, citric acid shares the properties of other carboxylic acids. When heated above 175 °C, it decomposes through the loss of carbon dioxide and water (see decarboxylation).

Citric acid is a slightly stronger acid than typical carboxylic acids because the anion can be stabilized by intramolecular hydrogen-bonding from other protic groups on citric acid.

2 - Discovery and production

Lemons, oranges, limes, and other citrus fruits possess high concentrations of citric acid

Medieval scholars in Europe were aware of the acidity of lemon and lime juices; such knowledge is recorded in the 13th century encyclopedia Speculum Maius (The Great Mirror), compiled by Vincent of Beauvais.[citation needed] Citric acid was first isolated in 1784 by the chemist Carl Wilhelm Scheele, who crystallized it from lemon juice.

Industrial-scale citric acid production first began in 1890 based on the Italian citrus fruit industry, where the juice was treated with hydrated lime (calcium hydroxide) to precipitate calcium citrate, which was isolated and converted back to the acid using diluted sulfuric acid.

3Ca(OH)_2(s) + 2C₆H₈O_7(l) → Ca₃(C₆H₅O₇)_2(s) + 3H₂O_(l) 

3H₂SO_4(l) + Ca₃(C₆H₅O₇)_2(s) → 2C₆H₈O_7(l) + 3CaSO_4(s)

In 1893, C. Wehmer discovered Penicillium mold could produce citric acid from sugar. However, microbial production of citric acid did not become industrially important until World War I disrupted Italian citrus exports.

In 1917, American food chemist James Currie discovered certain strains of the mold Aspergillus niger could be efficient citric acid producers, and the pharmaceutical company Pfizer began industrial-level production using this technique two years later, followed by Citrique Belge in 1929.

In this production technique, which is still the major industrial route to citric acid used today, cultures of A. niger are fed on a sucrose or glucose-containing medium to produce citric acid. The source of sugar is corn steep liquor, molasses, hydrolyzed corn starch or other inexpensive sugary solutions. After the mold is filtered out of the resulting solution, citric acid is isolated by precipitating it with calcium hydroxide to yield calcium citrate salt, from which citric acid is regenerated by treatment with sulfuric acid, as in the direct extraction from citrus fruit juice.

In 1977, a patent was granted to Lever Brothers for the chemical synthesis of citric acid starting either from aconitic or isocitrate/alloisocitrate calcium salts under high pressure conditions. This produced citric acid in near quantitative conversion under what appeared to be a reverse non-enzymatic Krebs cycle reaction.

In 2007, worldwide annual production stood at approximately 1,600,000 tons. More than 50% of this volume was produced in China. More than 50% was used as acidity regulator in beverages, some 20% in other food applications, 20% for detergent applications and 10% for related applications other than food, such as cosmetics, pharmaceutics and in the chemical industry.

3 - Occurrence

Citric acid exists in greater than trace amounts in a variety of fruits and vegetables, most notably citrus fruits. Lemons and limes have particularly high concentrations of the acid; it can constitute as much as 8% of the dry weight of these fruits (about 47 g/L in the juices). The concentrations of citric acid in citrus fruits range from 0.005 mol/L for oranges and grapefruits to 0.30 mol/L in lemons and limes. Within species, these values vary depending on the cultivar and the circumstances in which the fruit was grown.

4 - Biochemistry

- Citric acid cycle
Main article: Citric acid cycle

Citrate, the conjugate base of citric acid is one of a series of compounds involved in the physiological oxidation of acetate from fats, proteins, and carbohydrates. The acetate from these macronutrients are converted into the intracellular energy of ATP, as well as the common byproducts carbon dioxide, and water.

This series of chemical reactions is central to nearly all metabolic reactions, and is the source of two-thirds of the food-derived energy in higher organisms. Hans Adolf Krebs received the 1953 Nobel Prize in Physiology or Medicine for the discovery. The series of reactions is known by various names, including the "citric acid cycle", the "Krebs cycle" or "Szent-Györgyi — Krebs cycle", and the "tricarboxylic acid (TCA) cycle".

- Other biological roles

Citrate is a vital component of bone, helping to regulate the size of calcium crystals.

5 - Applications

The dominant use of citric acid is as a flavoring and preservative in food and beverages, especially soft drinks. Within the European Union it is denoted by E number E330. Citrate salts of various metals are used to deliver those minerals in a biologically available form in many dietary supplements. The buffering properties of citrates are used to control pH in household cleaners and pharmaceuticals. In the United States the purity requirements for citric acid as a food additive are defined by the Food Chemicals Codex, which is published by the United States Pharmacopoeia (USP).

- Foods, other

Citric acid can be added to ice cream as an emulsifying agent to keep fats from separating, to caramel to prevent sucrose crystallization, or in recipes in place of fresh lemon juice. Citric acid is used with sodium bicarbonate in a wide range of effervescent formulae, both for ingestion (e.g., powders and tablets) and for personal care (e.g., bath salts, bath bombs, and cleaning of grease). Citric acid is also often used in cleaning products and sodas or fizzy drinks.

Citric acid sold in a dry powdered form is commonly sold in markets and groceries as "sour salt", due to its physical resemblance to table salt. It has use in culinary applications where an acid is needed for either its chemical properties or for its sour flavor, but a dry ingredient is needed and additional flavors are unwanted (e.g., instead of vinegar or lemon juice).

- Cleaning and chelating agent

Citric acid is an excellent chelating agent, binding metals. It is used to remove limescale from boilers and evaporators. It can be used to soften water, which makes it useful in soaps and laundry detergents. By chelating the metals in hard water, it lets these cleaners produce foam and work better without need for water softening. Citric acid is the active ingredient in some bathroom and kitchen cleaning solutions. A solution with a 6% concentration of citric acid will remove hard water stains from glass without scrubbing. In the industry, it is used to dissolve rust from steel. Citric acid can be used in shampoo to wash out wax and coloring from the hair.

Illustrative of its chelating abilities, citric acid was the first successful eluant used for total ion-exchange separation of the lanthanides, during the Manhattan Project in the 1940s. In the 1950s, it was replaced by the far more efficient EDTA. It can be used to substantially slow setting of Portland cement.

- Cosmetics and pharmaceuticals

Citric acid is widely used as a pH adjusting agent in creams and gels of all kinds. In this role, it is classified in most jurisdictions as a processing aid and so does not need to be listed on ingredient lists.

Citric acid is an alpha hydroxy acid and used as an active ingredient in chemical peels.

Citric acid is commonly used as a buffer to increase the solubility of brown heroin. Single-use citric acid sachets have been used as an inducement to get heroin users to exchange their dirty needles for clean needles in an attempt to decrease the spread of HIV and hepatitis.[15] Other acidifiers used for brown heroin are ascorbic acid, acetic acid, and lactic acid; in their absence, a drug user will often substitute lemon juice or vinegar.

Citric acid is used as one of the active ingredients in the production of antiviral tissues

- Dyeing

Citric acid can be used in food coloring to balance the pH level of a normally basic dye. It is used as an odorless alternative to white vinegar for home dyeing with acid dyes.
Qualitative analysis

Sodium citrate, the sodium salt of citric acid, is used as a chelating agent and is present in the Benedict's reagent, used for identification both qualitatively and quantitatively, of reducing sugars.

- Industrial and construction

Citric acid can be used as a successful alternative to nitric acid in passivation of stainless steel.

- Photography

Citric acid can be used as a lower-odor stop bath as part of the process for developing photographic film. Photographic developers are alkaline, so a mild acid is used to neutralize and stop their action quickly, but commonly used acetic acid leaves a strong vinegar odor in the darkroom.

- Synthesize solid materials from small molecules

In materials science, the Citrate-gel method is similar process to sol-gel method which is a method for producing solid materials from small molecules. During the synthetic process, metal salts or alkoxides are introduced into a citric acid solution. The formation of citric complexes is believed to balance the difference in individual behaviour of ions in solution, which results in a better distribution of ions and prevents the separation of components at later process stages. The polycondensation of ethylene glycol and citric acid starts above 100ºС, resulting in polymer citrate gel formation.

6 - Safety

Although a weak acid, exposure to pure citric acid can cause adverse effects: inhalation may cause cough, shortness of breath, or sore throat; ingestion may cause abdominal pain and sore throat; exposure to skin or eyes may cause redness or pain. Long-term or repeated consumption may cause erosion of tooth enamel.

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Sodium Bicarbonate or Sodium Hydrogen Carbonate

Sodium Bicarbonate or Sodium Hydrogen Carbonate

Sodium bicarbonate (IUPAC name: sodium hydrogen carbonate) is a chemical compound with the formula  NaHCO₃. Sodium bicarbonate is a white solid that is crystalline but often appears as a fine powder. It has a slightly salty, alkaline taste resembling that of washing soda (sodium carbonate). The natural mineral form is nahcolite. It is a component of the mineral natron and is found dissolved in many mineral springs. It is among the food additives encoded by European Union, identified by the initials E 500. Since it has long been known and is widely used, the salt has many related names such as baking soda, bread soda, cooking soda, and bicarbonate of soda. The word saleratus, from Latin sal æratus meaning aerated salt, was widely used in the 19th century for both sodium bicarbonate and potassium bicarbonate.

1 - History

The ancient Egyptians used natural deposits of natron, a mixture consisting mostly of sodium carbonate decahydrate, and sodium bicarbonate. The natron was ground up, solvated, and used as paint for hieroglyphics.

In 1791, a French chemist, Nicolas Leblanc, produced sodium carbonate, also known as soda ash. In 1846, two New York bakers, John Dwight and Austin Church, established the first factory to develop baking soda from sodium carbonate and carbon dioxide.

This compound, referred to as saleratus, is mentioned in the novel Captains Courageous by Rudyard Kipling as being used extensively in the 1800s in commercial fishing to prevent freshly-caught fish from spoiling.

2 - Production
Main article: Solvay processNaHCO₃ is mainly prepared by the Solvay process, which is the reaction of sodium chloride, ammonia, and carbon dioxide in water. Calcium carbonate is used as the source of CO₂ and the resultant calcium oxide is used to recover the ammonia from the ammonium chloride. The product shows a low purity (75%). Pure product is obtained from sodium carbonate, water and carbon dioxide as reported in one of the following reactions. It is produced on the scale of about 100,000 tonnes/year (as of 2001).NaHCO₃ may be obtained by the reaction of carbon dioxide with an aqueous solution of sodium hydroxide. The initial reaction produces sodium carbonate:

CO₂ + 2 NaOH → Na₂CO₃ + H₂O

Further addition of carbon dioxide produces sodium bicarbonate, which at sufficiently high concentration will precipitate out of solution:

Na₂CO₃ + CO₂ + H₂O → 2 NaHCO₃

Commercial quantities of baking soda are also produced by a similar method: soda ash, mined in the form of the ore trona, is dissolved in water and treated with carbon dioxide. Sodium bicarbonate precipitates as a solid from this method: 

Na₂CO₃ + CO₂ + H₂O → 2 NaHCO₃

3 - Mining

Naturally occurring deposits of nahcolite (NaHCO₃) are found in the Eocene-age (55.8–33.9 Mya) Green River Formation, Piceance Basin in Colorado. Nahcolite was deposited as beds during periods of high evaporation in the basin. It is commercially mined using in situ leach techniques involving dissolution of the nahcolite by heated water pumped through the nahcolite beds and reconstituted through a natural cooling crystallisation process.

4 - Chemistry

Sodium bicarbonate is an amphoteric compound. Aqueous solutions are mildly alkaline due to the formation of carbonic acid and hydroxide ion:

HCO−
3 + H₂O → H
2CO
3 + OH⁻

Sodium bicarbonate can be used as a wash to remove any acidic impurities from a "crude" liquid, producing a purer sample. Reaction of sodium bicarbonate and an acid produce a salt and carbonic acid, which readily decomposes to carbon dioxide and water:

NaHCO₃ + HCl → NaCl + H₂CO₃

H₂CO₃ → H₂O + CO₂(g)

Sodium bicarbonate reacts with acetic acid (found in vinegar), producing sodium acetate, water, and carbon dioxide: 

NaHCO₃ + CH₃COOH → CH₃COONa + H₂O + CO₂(g)

Sodium bicarbonate reacts with bases such as sodium hydroxide to form carbonates:

NaHCO₃ + NaOH → Na₂CO₃ + H₂O

Sodium bicarbonate reacts with carboxyl groups in proteins to give a brisk effervescence from the formation of CO₂. This reaction is used to test for the presence of carboxylic groups in protein.

- Thermal decomposition

Above 50 °C, sodium bicarbonate gradually decomposes into sodium carbonate, water and carbon dioxide. The conversion is fast at 200 °C:

2 NaHCO₃ → Na₂CO₃ + H₂O + CO₂

Most bicarbonates undergo this dehydration reaction. Further heating converts the carbonate into the oxide (at over 850 °C):

Na₂CO₃ → Na₂O + CO₂

These conversions are relevant to the use of NaHCO₃ as a fire-suppression agent ("BC powder") in some dry powder fire extinguishers.

5 - Applications

A list of applications are below.

- Pest Control

Used to kill cockroaches. Once consumed, it causes internal organs of cockroaches to burst due to gas collection. Paint and Corrosion Removal

Sodium bicarbonate is used in a process for removing paint and corrosion called sodablasting; the process is particularly suitable for cleaning aluminium panels which can be distorted by other types of abrasive.

- PH Balancer

It can be administered to pools, spas, and garden ponds to raise pH levels.

- Mild Disinfectant

It has weak disinfectant properties, and it may be an effective fungicide against some organisms. Because baking soda will absorb musty smells, it has become a reliable method for used-book sellers when making books less malodorous.

- Fire extinguisher

Sodium bicarbonate can be used to extinguish small grease or electrical fires by being thrown over the fire, as heating of sodium bicarbonate releases carbon dioxide. However, it should not be applied to fires in deep fryers; the sudden release of gas may cause the grease to splatter. Sodium bicarbonate is used in BC dry chemical fire extinguishers as an alternative to the more corrosive ammonium phosphate in ABC extinguishers. The alkali nature of sodium bicarbonate makes it the only dry chemical agent, besides Purple-K, that was used in large-scale fire suppression systems installed in commercial kitchens. Because it can act as an alkali, the agent has a mild saponification effect on hot grease, which forms a smothering soapy foam.

- Cooking
Main article: Leavening agent

-- With acids

Sodium bicarbonate, referred to as "baking soda", is primarily used in cooking (baking), as a leavening agent. It reacts with acidic components in batters, releasing carbon dioxide, which causes expansion of the batter and forms the characteristic texture and grain in pancakes, cakes, quick breads, soda bread, and other baked and fried foods. Acidic compounds that induce this reaction include phosphates, cream of tartar, lemon juice, yogurt, buttermilk, cocoa, vinegar, etc. Natural acids in sourdough can be leavened with the addition of small amounts as well. Sodium bicarbonate can be substituted for baking powder provided sufficient acid reagent is also added to the recipe. Many forms of baking powder contain sodium bicarbonate combined with calcium acid phosphate, sodium aluminium sulphate or cream of tartar.

Sodium bicarbonate was sometimes used in cooking vegetables, to make them softer, although this has gone out of fashion, as most people now prefer firmer vegetables. However, it is still used in Asian and Latin American cuisine to tenderise meats. Baking soda may react with acids in food, including vitamin C (L-ascorbic acid). It is also used in breadings such as for fried foods to enhance crispness.

-- By heating

Heat causes sodium bicarbonate to act as a raising agent by releasing carbon dioxide when used in baking. The carbon dioxide production starts at temperatures above 80 °C. Since the reaction does not occur at room temperature, mixtures (cake batter, etc.) can be allowed to stand without rising until they are heated in the oven.

 2NaHCO₃ → Na2CO₃ + H₂O + CO₂

- Neutralisation of acids and bases

Sodium bicarbonate is amphoteric, reacting with acids and bases. It reacts violently with acids, releasing CO2 gas as a reaction product. However, sodium bicarbonate is not recommended for the clean up of acid spills. The heat produced increases the reactivity of the acid[citation needed], and the large amount of carbon dioxide produced may increase the area of the spill.

A wide variety of applications follows from its neutralisation properties, including reducing the spread of white phosphorus from incendiary bullets inside an afflicted soldier's wounds.

- Medical uses

Sodium bicarbonate mixed with water can be used as an antacid to treat acid indigestion and heartburn. It is used as the medicinal ingredient in gripe water for infants.

Intravenous sodium bicarbonate is an aqueous solution that is sometimes used for cases of acidosis, or when there are insufficient sodium or bicarbonate ions in the blood.[24] In cases of respiratory acidosis, the infused bicarbonate ion drives the carbonic acid/bicarbonate buffer of plasma to the left and, thus, raises the pH. It is for this reason that sodium bicarbonate is used in medically supervised cardiopulmonary resuscitation. Infusion of bicarbonate is indicated only when the blood pH is markedly (<7.1–7.0) low.

It is used for treatment of hyperkalemia as it will drive K+ back into cells during periods of hypocholermic metabolic alkalosis. Since sodium bicarbonate can cause alkalosis, it is sometimes used to treat aspirin overdoses. Aspirin requires an acidic environment for proper absorption, and the basic environment diminishes aspirin absorption in the case of an overdose. Sodium bicarbonate has also been used in the treatment of tricyclic antidepressant overdose. It can also be applied topically as a paste, with three parts baking soda to one part water, to relieve some kinds of insect bites and stings (as well as accompanying swelling).

Adverse reactions to the administration of sodium bicarbonate can include metabolic alkalosis, edema due to sodium overload, congestive heart failure, hyperosmolar syndrome, hypervolemic hypernatremia, and hypertension due to increased sodium. In patients consuming a high-calcium or dairy-rich diet, calcium supplements, or calcium-containing antacids such as calcium carbonate (e.g., Tums), the use of sodium bicarbonate can cause milk-alkali syndrome, which can result in metastatic calcification, kidney stones, and kidney failure.

Sodium bicarbonate can be used to treat an allergic reaction to plants such as poison -ivy -oak or -sumac to relieve some of the associated itching.

Bicarbonate of soda can also be useful in removing splinters from the skin.

Some alternative practitioners, such as Tullio Simoncini, have promoted baking soda as a cancer cure, which the American Cancer Society has warned against due to both its unproven effectiveness and potential danger in use.

- Personal hygiene

Toothpaste containing sodium bicarbonate has in several studies been shown to have a better whitening and plaque removal effect than toothpastes without it.

Sodium bicarbonate is also used as an ingredient in some mouthwashes. It has anti-caries and abrasive properties. It works as a mechanical cleanser on the teeth and gums, neutralises the production of acid in the mouth and also acts as an antiseptic to help prevent infections.

Sodium bicarbonate in combination with other ingredients can be used to make a dry or wet deodorant. It may also be used as a shampoo.

Sodium bicarbonate may be used as a buffering agent, combined with table salt, when creating a solution for nasal irrigation.

It is used in eye hygiene to treat blepharitis. This is done by addition of a tablespoon of sodium bicarbonate to cool water that was recently boiled, followed by gentle scrubbing of the eyelash base with a cotton swab dipped in the solution.

- In sports

Small amounts of sodium bicarbonate have been shown to be useful as a supplement for athletes in speed-based events, like middle distance running, lasting from about one to seven minutes. However, overdose is a serious risk because sodium bicarbonate is slightly toxic; and gastrointestinal irritation is of particular concern. Additionally, this practice causes a significant increase in dietary sodium.

- As a cleaning agent

A paste from baking soda can be very effective when used in cleaning and scrubbing. For cleaning aluminium objects, the use of sodium bicarbonate is discouraged as it attacks the thin unreactive protective oxide layer of this otherwise very reactive metal. A solution in warm water will remove the tarnish from silver when the silver is in contact with a piece of aluminium foil. A paste of sodium bicarbonate and water is useful in removing surface rust as the rust forms a water-soluble compound when in a concentrated alkaline solution. Cold water should be used as hot water solutions can corrode steel.

Baking soda is commonly added to washing machines as a replacement for softener and to remove odors from clothes. Sodium bicarbonate is also effective in removing heavy tea and coffee stains from cups when diluted with warm water.

During the Manhattan Project to develop the atomic bomb in the early 1940s, many scientists investigated the toxic properties of uranium. They found that uranium oxides stick very well to cotton cloth, but did not wash out with soap or laundry detergent. The uranium would wash out with a 2% solution of sodium bicarbonate (baking soda). Clothing can become contaminated with depleted uranium (DU) dust, and then normal laundering will not remove it. Those at risk of DU dust exposure should have their clothing washed with about 6 ounces (170 g) of baking soda in 2 gallons (7.5 l) of water).

- As a biopesticide

Sodium bicarbonate can be an effective way of controlling fungus growth, and in the United States is registered by the Environmental Protection Agency as a biopesticide.

- Cattle feed supplements

Sodium bicarbonate is sold as a cattle feed supplement, in particular as a buffering agent for the rumen.

6 - In popular culture

- Film

Sodium bicarbonate, as 'bicarbonate of soda', was a frequent source of punch lines for Groucho Marx in Marx brothers movies. In Duck Soup, Marx plays the leader of a nation at war. In one scene, he receives a message from the battlefield that his general is reporting a gas attack, and Groucho tells his aide, "Tell him to take a teaspoonful of bicarbonate of soda and a half a glass of water." In A Night at the Opera, Groucho's character addresses the opening night crowd at an opera by saying of the lead tenor, "Signor Lassparri comes from a very famous family. His mother was a well-known bass singer. His father was the first man to stuff spaghetti with bicarbonate of soda, thus causing and curing indigestion at the same time."

7 - Difference between baking soda and baking powder

Quite simply, baking powder contains baking soda, as well as a powdered acid and cornstarch. In scientific terms, baking soda is a pure substance; baking powder is a mixture.

Baking soda is alkaline, so acid is used in baking powder to avoid a metallic taste when the chemical change during baking creates sodium carbonate. However, to avoid the over-flavouring of acidic taste, non-acid ingredients such as whole milk or Dutch-processed cocoa must be added.

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Sodium Carbonate or Soda Ash Light

Sodium Carbonate or Soda Ash Light

Sodium carbonate (also known as washing soda, soda ash and soda crystals), Na₂CO₃, is the sodium salt of carbonic acid (soluble in water).

It most commonly occurs as a crystalline heptahydrate, which readily effloresces to form a white powder, the monohydrate. Pure sodium carbonate is a white, odourless powder that is hygroscopic (absorbs moisture from the air), has an alkaline taste, and forms a strongly alkaline water solution. Sodium carbonate is well known domestically for its everyday use as a water softener. It can be extracted from the ashes of many plants growing in sodium-rich soils, such as vegetation from the Middle East, kelp from Scotland and seaweed from Spain. Because the ashes of these sodium-rich plants were noticeably different from ashes of timber (used to create potash), they became known as "soda ash". It is synthetically produced in large quantities from salt (sodium chloride) and limestone by a method known as the Solvay process.

1 - Uses

The manufacture of glass is one of the most important uses of sodium carbonate. Sodium carbonate acts as a flux for silica, lowering the melting point of the mixture to something achievable without special materials. This "soda glass" is mildly water-soluble, so some calcium carbonate is added to the pre-melt mixture to make the glass produced insoluble. This type of glass is known as soda lime glass: "soda" for the sodium carbonate and "lime" for the calcium carbonate. Soda lime glass has been the most common form of glass for centuries.

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 photographic film developing agents.

It is a common additive in swimming pools used to neutralize the corrosive effects of chlorine and raise the pH.

In cooking, it is sometimes used in place of sodium hydroxide for lyeing, especially with German pretzels and lye rolls. These dishes are treated with a solution of an alkaline substance to change the pH of the surface of the food and 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. Electrolytes are usually salt-based, and sodium carbonate acts as a very good conductor in the process of electrolysis. In addition, 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.

- Domestic use

It is used as a water softener in laundering: it competes with the magnesium and calcium ions in hard water and prevents them from bonding with the detergent being used. Sodium carbonate can be used to remove grease, oil and wine stains. Sodium carbonate is also used as a descaling agent in boilers such as those found in coffee pots and espresso machines.

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 cellulose (plant) fibers, typically before dyeing (for tie dyes), mixed with the dye (for dye painting), or after dyeing (for immersion dyeing).

- Sodium carbonate test

The sodium carbonate test (not to be confused with sodium carbonate extract test) is used to distinguish between some common metal ions, which are precipitated as their respective carbonates. The test can distinguish between Cu, Fe and Ca/Zn/Pb. Sodium carbonate solution is added to the salt of the metal. A blue precipitate indicates Cu²⁺ ion. A dirty green precipitate indicates Fe²⁺ ion. A yellow-brown precipitate indicates  Fe³⁺ ion. A white precipitate indicates Ca²⁺, Zn²⁺ or Pb²⁺ ion. The compounds formed are, respectively, copper(II) carbonate, iron(II) carbonate, iron(III) oxide, calcium carbonate, zinc carbonate and lead(II) carbonate. This test is used to precipitate the ion present as almost all carbonates are insoluble. While this test is useful for telling these cations apart, it fails if other ions are present, because most metal carbonates are insoluble and will precipitate. In addition, calcium, zinc and lead ions all produce white precipitates with carbonate, making it difficult to distinguish between them. Instead of sodium carbonate, sodium hydroxide may be added, this gives nearly the same colours, except that lead and zinc hydroxides are soluble in excess alkali, and can hence be distinguished from calcium. For the complete sequence of tests used for qualitative cation analysis, see qualitative inorganic analysis.

- Other applications

Sodium 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 (かん水?), a solution of alkaline salts used to give ramen noodles their characteristic flavor and texture. It is also used in the production of snus (Swedish-style snuff) to stabilize the pH of the final product. In Sweden, snus is regulated as a food product because it is put into the mouth, requires pasteurization, and contains only ingredients that are approved as food additives.

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.

In China, it is used to replace lye-water in the crust of traditional Cantonese moon cakes, and in many other Chinese steamed buns and noodles.

Sodium carbonate is used by the brick industry as a wetting agent to reduce the amount of water needed to extrude the clay.

In casting, it is referred to as "bonding agent" and is used to allow wet alginate to adhere to gelled alginate.

Sodium carbonate is used in toothpastes, where it acts as a foaming agent and an abrasive, and to temporarily increase mouth pH.

Sodium carbonate is used by the cotton industry to neutralize the sulfuric acid needed for acid de-linting of fuzzy cottonseed.

Sodium carbonate, in a solution with common salt, may be used for cleaning silver. In a non-reactive container (glass, plastic or ceramic) aluminium foil and the silver object are immersed in the hot salt solution. The elevated pH dissolves the aluminium oxide layer on the foil and enables an electrolytic cell to be established. Hydrogen ions produced by this reaction reduce the sulfide ions on the silver restoring silver metal. The sulfide can be released as small amounts of hydrogen sulfide. Rinsing and gently polishing the silver restores a highly polished condition.

Sodium carbonate is used in some aquarium water pH buffers to maintain a desired pH and carbonate hardness (KH).

Because of its ability to absorb CO₂, sodium carbonate is being investigated as a carbon-capturing material for power plants and in other industries that produce greenhouse gases.

2 - Physical properties

The integral enthalpy of solution of sodium carbonate is −28.1 kJ/mol for a 10% w/w aqueous solution. The Mohs hardness of sodium carbonate monohydrate is 1.3.

3 - Occurrence

Sodium carbonate crystallizes from water to form three different hydrates:

    1. sodium carbonate decahydrate (natron)

    2. sodium carbonate heptahydrate (not known in mineral form)

    3. sodium carbonate monohydrate (thermonatrite).

Sodium carbonate is soluble in water, and can occur naturally in arid regions, especially in mineral deposits (evaporites) formed when seasonal lakes evaporate. Deposits of the mineral natron 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.

The anhydrous mineral form of sodium carbonate is quite rare and called natrite. Sodium carbonate also erupts from Ol Doinyo Lengai, Tanzania's unique volcano, and it is presumed to have erupted from other volcanoes in the past but, due to these minerals' instability at the earth's surface, are likely to be eroded. All three mineralogical forms of sodium carbonate, as well as trona, trisodium hydrogendicarbonate dihydrate, are also known from ultra-alkaline pegmatitic rocks, that occur for example in the Kola Peninsula in Russia.

4 - Production

- Mining

Trona, trisodium hydrogendicarbonate dihydrate (Na₃HCO₃CO₃·2H₂O), is mined in several areas of the US and provides nearly all the domestic consumption of 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. There are important reserves of Trona in Turkey; two million tons of soda ash have been extracted from the reserves near Ankara. It is also mined from some alkaline lakes such as Lake Magadi in Kenya by dredging. Hot saline springs continuously replenish salt in the lake so that, provided the rate of dredging is no greater than the replenishment rate, the source is fully sustainable.

- Barilla and kelp

Several "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 percent for the seaweed-derived form ("kelp"), to 30 percent 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.

- Leblanc process
Main article: Leblanc process

In 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 equation:

2 NaCl + H₂SO₄ → Na₂SO₄ + 2 HCl

Next, 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 + CaS

The 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.

- Solvay process
Main article: Solvay process

In 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 + CO₂

At 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 + NH₄Cl

The sodium bicarbonate was then converted to sodium carbonate by heating it, releasing water and carbon dioxide:

2 NaHCO₃ → Na₂CO₃ + H₂O + CO₂

Meanwhile, 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)₂

Ca(OH)₂ + 2 NH₄Cl → CaCl₂ + 2 NH₃ + 2 H₂O

Because 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.
 

- Hou's process

Developed by Chinese chemist Hou Debang in the 1930s. The earlier steam reforming byproduct carbon dioxide was pumped through a saturated solution of sodium chloride and ammonia to produce sodium bicarbonate via the following reactions:

NH₃ + CO₂ + H₂O → NH₄HCO₃

NH₄HCO₃ + NaCl → NH₄Cl + NaHCO₃

The sodium bicarbonate was collected as a precipitate due to its low solubility and then heated to yield pure sodium carbonate similar to last step of the Solvay process. More sodium chloride is added to the remaining solution of ammonium and sodium chlorides; also more ammonia is pumped at 30-40 °C to this solution. The solution temperature is then lowered to below 10 °C. Solubility of ammonium chloride is higher than that of sodium chloride at 30 °C and lower at 10 °C. Due to this temperature dependent solubility difference and the common-ion effect, ammonium chloride is precipitated in a sodium chloride solution.The Chinese name of Hou's process, lianhe zhijian fa (联合制碱法), means "Coupled Manufacturing Alkali Method": Hou's process is coupled to the Haber process and offers better atom economy by eliminating the production of calcium chloride since ammonia no longer needs to be regenerated. The byproduct ammonium chloride can be sold as a fertilizer.

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Peracetic acid or Proxitane 15%

Peracetic acid or Proxitane 15%

Peracetic acid (also known as peroxyacetic acid, or PAA), is an organic compound with the formula CH₃CO₃H. This organic peroxide is a colorless liquid with a characteristic acrid odor reminiscent of acetic acid. It can be highly corrosive.

Peracetic acid is a weaker acid than the parent acetic acid, with a pKa ( acid dissociation constant ) of 8.2.

1 - Production

Peracetic acid is produced industrially by the autoxidation of acetaldehyde:

O₂ + CH₃CHO → CH₃CO₃H

It forms upon treatment of acetic acid with hydrogen peroxide, with the equilibrium constant dependent on the concentrations and conditions of reaction:

H₂O₂ + CH₃CO₂H CH₃CO₃H + H₂O

As an alternative, acetyl chloride and acetic anhydride can be used to generate a solution of the acid with lower water content.

Peracetic acid is generated in situ by some laundry detergents. This route involves the reaction of tetraacetylethylenediamine (TAED) in the presence of an alkaline hydrogen peroxide solution. The peracetic acid is a more effective bleaching agent than hydrogen peroxide itself. PAA is also formed naturally in the environment through a series of photochemical reactions involving formaldehyde and photo-oxidant radicals.

Peracetic acid is always sold in solution with acetic acid and hydrogen peroxide to maintain the stability of the peracid. The concentration of the acid as the active ingredient can vary.

2 - Uses

The United States Environmental Protection Agency first registered peracetic acid as an antimicrobial in 1985 for indoor use on hard surfaces. Use sites include agricultural premises, food establishments, medical facilities, and home bathrooms. Peracetic acid is also registered for use in dairy/cheese processing plants, on food processing equipment, and in pasteurizers in breweries, wineries, and beverage plants. It is also applied for the disinfection of medical supplies, to prevent bio film formation in pulp industries, and as a water purifier and disinfectant. Peracetic acid can be used as a cooling tower water disinfect, where it prevents bio film formation and effectively controls Legionella bacteria. A trade name for peracetic acid as an antimicrobial is Nu-Cidex.

- Epoxidation

Although less active than more acidic peracids (e.g., MCPBA), peracetic acid in various forms is used for the epoxidation of various alkenes. Useful application are for unsaturated fats, synthetic and natural rubbers, and some natural products such as pinene. A variety of factors affect the amount of free acid or sulfuric acid (used to prepare the peracid in the first place).

3 - Niche uses

Peracetic acid will oxidize many metals, and is used for cleaning or creating a patina for artistic or protective purposes.

4 - Safety

Peracetic acid is a strong oxidizing agent (E = 1.762 V vs Ag/AgCl) and a primary irritant. Exposure to peracetic acid can cause irritation to the skin, eyes and respiratory system and higher or long-term exposure can cause permanent lung damage. In addition, there have been cases of occupational asthma caused by peracetic acid. The ACGIH has published (spring 2014) a STEL TLV for peracetic acid of 0.4 ppm, calculated as a 15 minute time weighted average (TWA). Currently there is no OSHA Permissible Exposure Limit (PEL) for peracetic acid. In 2010, the US-EPA published Acute Exposure Guidelines (AEGL) for peracetic acid.

- 1 - The concentration at which the general population will experience transient and reversible problems, such as notable discomfort, irritation, or certain asymptomatic non-sensory effects - 0.52 mg/m3 - 0.17 ppm

- 2 - The concentration that results in irreversible or other serious, long-lasting adverse health effects or an impaired ability to escape - 1.6 mg/m3 - 0.52 ppm

- 3 - The concentration that results in life-threatening health effects or death - 4.1 mg/m3 - 1.3 ppm

In comparison, the OSHA PEL for hydrogen peroxide (8 hour TWA) is 1.0 ppm. Peracetic acid has found widespread use in healthcare, food processing, and water treatment because of its broad antimicrobial properties. In order for employers to meet the ACGIH STEL TLV and comply with the General Duty Clause to provide a safe work environment, it is recommended that facilities using peracetic acid solution employ continuous monitors to provide an alert to employees if the concentration exceeds safe levels. This monitor is designed and installed so that it can automatically turn on additional ventilation as needed.

Concentrated peroxyacetic acid, an organic peroxide, explodes at 110 °C.

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Trichloroisocyanuric acid or Chlorine TCCA 90%

Trichloroisocyanuric acid or Chlorine TCCA 90%

Trichloroisocyanuric acid is an organic compound with the formula (C₃Cl₃N₃O₃). It is used as an industrial disinfectant, bleaching agent and a reagent in organic synthesis. This white crystalline powder, which has a strong "chlorine odour," is sometimes sold in tablet or granule form for domestic and industrial use. Salts of trichloroisocyanuric acid are known as trichloroisocyanurates.

Applications

The compound is a disinfectant, algicide and bactericide mainly for swimming pools and dyestuffs, and is also used as a bleaching agent in the textile industry. It is widely used in civil sanitation for pools and spas, preventing and curing diseases in animal husbandry and fisheries, fruit and vegetable preservation, wastewater treatment, as an algicide for recycled water in industry and air conditioning, in anti shrink treatment for woolens, for treating seeds and in organic chemical synthesis.

Trichloroisocyanuric acid as used in swimming pools is easier to handle than chlorine. It dissolves slowly in water, but as it reacts, cyanuric acid concentration in the pool will build-up. At high cyanuric acid concentrations, normal chlorine levels can be rendered ineffective, requiring either dilution by draining and refilling the pool or by adding abnormally high doses of chlorine to overcome this effect.

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Calcium hypochlorite or Calcium oxychloride

Calcium hypochlorite or Calcium oxychloride

Calcium hypochlorite is an inorganic compound with formula Ca(OCl)
2. As a mixture with lime and calcium chloride, it is marketed as chlorine powder or bleach powder for water treatment and as a bleaching agent. This compound is relatively stable and has greater available chlorine than sodium hypochlorite (liquid bleach).It is a white solid, although commercial samples appear yellow. It strongly smells of chlorine, owing to its slow decomposition in moist air. It is not highly soluble in water and is more preferably used in soft to medium-hard water. It has two forms: dry and hydrated.

1 - Uses

- Sanitation

Calcium hypochlorite is commonly used to sanitize public swimming pools and disinfect drinking water. Generally the commercial substance is sold with a purity of a 68% (with other additives and contaminants varying based upon the product's intended purpose). For instance as a swimming pool chemical it is often mixed with cyanuric acid stabilizers and anti-scaling agents (in order to reduce the loss of chlorine from ultraviolet radiation and to prevent calcium hardening). Calcium hypochlorite is also used in kitchens to disinfect surfaces and equipment. Other common uses include bathroom cleansers, household disinfectant sprays, algaecides, herbicides, and laundry detergents.

- Organic chemistry

Calcium hypochlorite is a general oxidizing agent and therefore finds some use in organic chemistry. For instance the compound is used to cleave glycols, α-hydroxy carboxylic acids and keto acids to yield fragmented aldehydes or carboxylic acids. Calcium hypochlorite can also be used in the haloform reaction to manufacture chloroform.

2 - Production

Calcium hypochlorite is produced industrially by treating lime Ca(OH)
2 with chlorine gas. The reaction can be conducted in stages to give various compositions, each with different concentration of calcium hypochlorite, together with unconverted lime and calcium chloride. The full conversion is shown.

2 Cl
2 + 2 Ca(OH)
2 → Ca(OCl)
2 + CaCl
2 + 2 H
2O

Bleaching powder is made with slightly moist slaked lime. It is not a simple mixture of calcium hypochlorite, calcium chloride, and calcium hydroxide. Instead, it is a mixture consisting principally of calcium hypochlorite Ca(OCl)
2, dibasic calcium hypochlorite, Ca₃(OCl)₂(OH)₄, and dibasic calcium chloride, Ca₃Cl₂(OH)₄.

3 - Properties

Calcium hypochlorite reacts with carbon dioxide to form calcium carbonate and release dichlorine monoxide:

Ca(ClO)
2 + CO
2 → CaCO
3 + Cl
2O↑

A calcium hypochlorite solution is basic. This basicity is due to the hydrolysis performed by the hypochlorite ion, as hypochlorous acid is weak, but calcium hydroxide is a strong base. As a result, the hypochlorite ion is a strong conjugate base, and the calcium ion is a weak conjugate acid:

ClO− + H₂O → HClO + OH−

Similarly, calcium hypochlorite reacts with hydrochloric acid to form calcium chloride, water and chlorine:

Ca(OCl)₂ + 4 HCl → CaCl₂ + 2 H₂O + 2 Cl₂

4 - Safety

Calcium hypochlorite is stored dry and cold, away from any organic material and metals. The hydrated form is safer to handle.

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