preparation of sodium chloride in the laboratory

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From Wikipedia

Sodium hypochlorite

Sodium hypochlorite is a chemical compound with the formula NaOCl. Sodium hypochlorite solution, commonly known as bleach, is frequently used as a disinfectant or a bleaching agent.

Production

Hypochlorite was first produced in 1789 by Claude Louis Berthollet in his laboratory on the quay Javel in Paris, France, by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of sodium hypochlorite. However, this process was not very efficient and alternate production methods were sought. One such method involved the extraction of chlorinated lime (known as bleaching powder) with sodium carbonate to yield low levels of available chlorine. This method was commonly used to produce hypochlorite solutions for use as a hospital antiseptic which was sold under the trade names "Eusol" and "Dakin's solution".

Near the end of the nineteenth century, E. S. Smith patented a method of hypochlorite production involving hydrolysis of brine to produce caustic soda and chlorine gas which then mixed to form hypochlorite. Both electric power and brine solution were in cheap supply at this time, and various enterprising marketers took advantage of this situation to satisfy the market's demand for hypochlorite. Bottled solutions of hypochlorite were sold under numerous trade names; one such early brand produced by this method was called Parozone.

Today, an improved version of this method, known as the Hooker process, is the only large scale industrial method of sodium hypochlorite production. In this process sodium hypochlorite (NaOCl) and sodium chloride (NaCl) are formed when chlorine is passed into cold and dilute sodium hydroxide solution. It is prepared industrially by electrolysis with minimal separation between the anode and the cathode. The solution must be kept below 40 °C (by cooling coils) to prevent the undesired formation of sodium chlorate.

Cl2 + 2 NaOH → NaCl + NaOCl + H2O

Sodium hydroxide and chlorine are commercially produced by the chloralkali process, and there is no need to isolate them to prepare sodium hypochlorite.

Hence, chlorine is simultaneously reduced and oxidized; this process is known as disproportionation.

The commercial solutions always contain significant amounts of sodium chloride (common salt) as the main by-product, as seen in the equation above.

Sodium hypochlorite can be also made by electrolyzing saturated sodium chloride solution and the product can be tested by dropping hydrochloric acid to determine if it is successfully synthesized.

Packaging and sale

Household bleach sold for use in laundering clothes is a 3-6% solution of sodium hypochlorite at the time of manufacture. Strength varies from one formulation to another and gradually decreases with long storage.

A 12% solution is widely [http://www.awwa.org/publications/AWWAJournalArticle.cfm?itemnumber=40152&showLogin=N] used in waterworks for the chlorination of water and a 15% solution is more commonly used for disinfection of waste water in treatment plants. High-test hypochlorite (HTH) is sold for chlorination of swimming pools and contains approximately 30% calcium hypochlorite. The crystalline salt is also sold for the same use; this salt usually contains less than 50% of calcium hypochlorite. However, the level of active chlorine may be much higher.

It can also be found on store shelves in Daily Sanitizing Sprays, as the sole active ingredient at 0.0095%.

Reactions

Sodium hypochlorite reacts with metals gradually, such as zinc, to produce the metal oxide or hydroxide:

NaOCl + Zn → ZnO + NaCl

It reacts with hydrochloric acid to release chlorine gas:

NaOCl + 2 HCl → Cl2 + H2O + NaCl

It reacts with other acids, such as acetic acid, to release hypochlorous acid:

NaOCl + CH3COOH → HClO + CH3COONa

It decomposes when heated or evaporated to form sodium chlorate and sodium chloride:

3 NaOCl → NaClO3 + 2 NaCl

In reaction with hydrogen peroxide it gives off molecular oxygen:

NaOCl + H2O2→ H2O + NaCl + O2↑

Uses

Bleaching

In household form, sodium hypochlorite is used for removal of stains from laundry. It is particularly effective on cotton fiber, which stains easily but bleaches well. Usually 50 to 250 mL of bleach per load is recommended for a standard-size washer. The properties of household bleach that make it effective for removing stains also result in cumulative damage to organic fibers such as cotton, and the useful lifespan of these materials will be shortened with regular bleaching. The sodium hydroxide (NaOH) that is also found in household bleach (as noted later) causes fiber degradation as well. It is not volatile, and residual amounts of NaOH not rinsed out will continue slowly degrading organic fibers in the presence of humidity. For these reasons, if stains are localized, spot treatments should be considered whenever possible. With safety precautions, post-treatment with vinegar (or another weak acid) will neutralize the NaOH, and volatilize the chlorine from residual hypochlorite. Old t-shirts and cotton sheets that rip easily demonstrate the costs of laundering with household bleach. Hot water increases the activity of the bleach, owing to the increased kinetic energy of the molecules.

Disinfection

A weak solution of 1% household bleach in warm water is used to sanitize smooth surfaces prior to brewing of beer or wine. Surfaces must be rinsed to avoid imparting flavors to the brew; these chlorinated byproducts of sanitizing surfaces are also harmful.

US Government regulations (21 CFR Part 178) allow food processing equipment and food contact surfaces to be sanitized with solutions containing bleach, provided that the solution is allowed to drain adequately before contact with food, and that the solutions do not exceed 200 parts per million

Magnesium chloride

Magnesium chloride is the name for the chemical compounds with the formulas MgCl2 and its various hydrates MgCl2(H2O)x. These salts are typical ionic halides, being highly soluble in water. The hydrated magnesium chloride can be extracted from brine or sea water. Magnesium chloride as the natural mineral bischofite is also extracted (solution mining) out of ancient seabeds, for example the Zechstein seabed in northwest Europe. Anhydrous magnesium chloride is the principal precursor to magnesium metal, which is produced on a large scale.

Structure, preparation, and general properties

MgCl2 crystallizes in the cadmium chloride motif, which features octahedral Mg. A variety of hydrates are known with the formula MgCl2(H2O)x, and each loses water with increasing temperature: x = 12 (-16.4 °C), 8 (-3.4 °C), 6 (116.7 °C), 4 (181 °C), 2 (ca. 300 °C). In the hexahydrate, the Mg2+ remains octahedral, but is coordinated to six water ligands. The thermal dehydration of the hydrates MgCl2(H2O)x (x = 6, 12) does not occur straightforwardly.

As suggested by the existence of some hydrates, anhydrous MgCl2 is a Lewis acid, although a relatively weak one.

In the Dow process, magnesium chloride is regenerated from magnesium hydroxide using hydrochloric acid:

Mg(OH)2(s) + 2 HCl → MgCl2(aq) + 2 H2O(l)

It can also be prepared from magnesium carbonate by a similar reaction.

In most of its derivatives, MgCl2 forms octahedral complexes. Derivatives with tetrahedral Mg2+ are less common. Examples include salts of (tetraethylammonium)2MgCl4 and adducts such as MgCl2(TMEDA).

Applications

Magnesium chloride serves as precursor to other magnesium compounds, for example by precipitation:

MgCl2(aq) + Ca(OH)2(aq) → Mg(OH)2(s) + CaCl2(aq)

It can be electrolysed to give magnesium metal:

MgCl2(l) → Mg(l) + Cl2(g)

This process is practiced on a substantial scale.

Magnesium chloride is used for a variety of other applications besides the production of magnesium: the manufacture of textiles, paper, fireproofing agents, cements and refrigeration brine, and dust and erosion control. Mixed with hydrated magnesium oxide, magnesium chloride forms a hard material called Sorel cement.

Magnesium ion Mg2+ (usually added as the chloride) is an important component in the polymerase chain reaction, a procedure used to amplify DNA fragments. It is generally used in experimental biology whenever RNA and DNA and their enzymes are to function in vitro, since Mg2+ is a necessary associate ion for nucleotides in biology, such as ATP.

Magnesium chloride is also used in several medical and topical (skin related) applications. It has been used in pills as supplemental sources of magnesium, where it serves as a soluble compound which is not as laxitive as magnesium sulfate, and more bioavailable than magnesium hydroxide and magnesium oxide, since it does not require stomach acid to produce soluble Mg2+ ion. It can also be used as an effective anaesthetic for cephalopods, some species of crustaceans, and several species of bivalve, including oysters.

Culinary use

Magnesium chloride is an important coagulant (E511 ) used in the preparation of tofu from soy milk. In Japan it is sold as nigari (��り, derived from the Japanese word for "bitter"), a white powder produced from seawater after the sodium chloride has been removed, and the water evaporated. In China it is called lushui (�水). Nigari or lushui consists mostly of magnesium chloride, with some magnesium sulfate and other trace elements. It is also an ingredient in baby formula milk.

Use as

Sodium acetate

Sodium acetate, NaOAc, also sodium ethanoate, is the sodiumsalt of acetic acid. This colourless salt has a wide range of uses.

Applications

Industrial

Sodium acetate is used in the textile industry to neutralize sulfuric acid waste streams, and as a photoresist while using aniline dyes. It is also a pickling agent in chrome tanning, and it helps to retard vulcanization of chloroprene in synthetic rubber production.

Food

Sodium acetate may be added to foods as a seasoning. It may be used in the form of sodium diacetate— a 1:1 complex of sodium acetate and acetic acid, given the E-numberE262. A frequent use of this form is in salt and vinegar chips in the United States. Many US brands, including national manufacturer Frito-Lay, sell "salt and vinegar flavoured" chips that use this chemical, with lactose and smaller percentages of other chemicals, in lieu of a real salt and vinegar preparation.

Buffer solution

As the conjugate base of a weak acid, a solution of sodium acetate and acetic acid can act as a buffer to keep a relatively constant pH. This is useful especially in biochemical applications where reactions are pH dependent.

Heating pad

Sodium acetate is also used in consumer heating pads or hand warmers and is also used in hot ice. Sodium acetate trihydrate crystals melt at 58°C, dissolving in their water of crystallization. When they are heated to around 100°C, and subsequently allowed to cool, the aqueous solution becomes supersaturated. This solution is capable of cooling to room temperature without forming crystals. By clicking on a metal disc in the heating pad, a nucleation centre is formed which causes the solution to crystallize into solid sodium acetate trihydrate again. The bond-forming process of crystallization is exothermic, hence heat is emitted. The latent heat of fusion is about 264–289 kJ/kg. Unlike some other types of heat packs that depend on irreversible chemical reactions, sodium acetate heat packs can be easily recharged by boiling until all crystals are dissolved; they can be reused indefinitely.

Preparation

For laboratory use, sodium acetate is inexpensive, and is usually purchased instead of being synthesized. It is sometimes produced in a laboratory experiment by the reaction of acetic acid with sodium carbonate, sodium bicarbonate, or sodium hydroxide. These reactions produce aqueous sodium acetate and water. Carbon dioxide is produced in the reaction with sodium carbonate and bicarbonate, and it leaves the reaction vessel as a gas (unless the reaction vessel is pressurized). This is the well-known "volcano" reaction between baking soda and vinegar.

CH3COOH + NaHCO3→ CH3COONa + H2O + CO2

Industrially, sodium acetate is prepared from glacial acetic acid and sodium hydroxide.

C2H4O2 + NaOH → NaO2CCH3 + H2O

Reactions

Sodium acetate can be used to form an ester with an alkyl halide such as bromoethane:

NaO2CCH3 + BrCH2CH3→ C2H5O2CCH3 + NaBr

Caesium salts catalyze this reaction.


Sodium hydride

Sodium hydride is the chemical compound with the empirical formulaNaH. It is primarily used as a strong base in organic synthesis. NaH is representative of the saline hydrides, meaning it is a salt-like hydride, composed of Na+ and H− ions, in contrast to the more molecular hydrides such as borane, methane, ammonia and water. It is an ionic material that is insoluble in organic solvents (although soluble in molten Na), consistent with the fact that H− remains an unknown anion in solution. Because of the insolubility of NaH, all reactions involving NaH occur at the surface of the solid.

Basic properties and structure

NaH is produced by the direct reaction of hydrogen and liquid sodium. Pure NaH is colorless, although samples generally appear grey. NaH is ca. 40% denser than Na (0.968 g/cm³).

NaH, like LiH, KH, RbH, and CsH, adopts the NaClcrystal structure. In this motif, each Na+ ion is surrounded by six H− centers in an octahedral geometry. The ionic radii of H− (146 pm in NaH) and F− (133 pm) are comparable, as judged by the Na−H and Na−F distances.

Applications in organic synthesis

As a strong base

First and foremost, NaH is a base of wide scope and utility in organic chemistry. It is capable of deprotonating a range of even weak Brønsted acids to give the corresponding sodium derivatives. Typical "easy" substrates contain O-H, N-H, S-H bonds, including alcohols, phenols, pyrazoles, and thiols.

NaH most notably is employed to deprotonate carbon acids such as 1,3-dicarbonyls and analogues such as malonic esters. The resulting sodium derivatives can be alkylated. NaH is widely used to promote condensation reactions of carbonyl compounds via the Dieckmann condensation, Stobbe condensation, Darzens condensation, and Claisen condensation. Other carbon acids susceptible to deprotonation by NaH include sulfonium salts and DMSO. NaH is used to make sulfurylides, which in turn are used to convert ketones into epoxides.

As a reducing agent

NaH reduces certain main group compounds, but analogous reactivity is unknown in organic chemistry. Notably boron trifluoride reacts to give diborane and sodium fluoride:

6 NaH + 2 BF3→ B2H6 + 6 NaF

Si-Si and S-S bonds in disilanes and disulfides are also reduced.

Drying agent

Because of its rapid and irreversible reaction with water, NaH can be used to dry some organic solvents. Other drying agents are far more widely used, such as calcium hydride.

Hydrogen storage

The use of sodium hydride has been proposed for hydrogen storage for use in fuel cell vehicles, the hydride being encased in plastic pellets which are crushed in the presence of water to release the hydrogen.

Practical considerations

Sodium hydride is sold by many chemical suppliers usually as a mixture of 60% sodium hydride (w/w) in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH. The pure white solid is prepared by rinsing the oil with pentane or THF, care being taken because the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as nitrogen or argon gas. Typically NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds.

Safety

NaH can ignite in air, especially upon contact with water to release hydrogen, which is also flammable. Hydrolysis converts NaH into sodium hydroxide (NaOH), a causticbase. In practice, most sodium hydride is dispensed as a dispersion in oil, which can be safely handled in air.



From Yahoo Answers

Question:Hello! The concentration is 0.5 M and the volume is 0.25 L. The question states: To design and carry out an experiment in which you prepare a standard solution of Sodium Chloride. It even says we should write the instructions of how to do so!! Please give me the steps!! THANKS!!

Answers:Molar mass NaCl = 58.4430 g/mol You want 0.5M and 0.25L volume: Mass = 58.443 * 0.50*0.25 = 7.305g Method: weigh out exactly 7.305g NaCl Dissolve in distilled water. Transfer solution to 250mL volumetric flask. Wash out dissolving vessel with distilled water and add to flask Make up to exactly 250mL Mix thoroughly.

Question:

Answers:Here's what I think the answer is: Acyl chloride's functional group is -COCl while carboxylic acid has -COOH Acyl chloride is very reactive because of the chlorine on the C=O (carbonyl carbon) which makes the structure very electronegative. Electronegativity is the ability of an atom to attract electron density towards itself (thus making it a very good electrophile). Carboxylic acids are polar and forms hydrogen bonding. If you compare the reactivity of this two, then acyl chloride is the better reagent because of the chlorine which in a way activates the the carbonyl group due to increased electronegativity. Hope this helps.. Ciao

Question:chem homework What is the molarity of a solution prepared by dissolving 27.2 g of sodium chloride in enough water to prepare 500.0 mL of solution?

Answers:Molarity = moles/L NaCl MW = 58.44277 g/mole 27.2 g NaCl = 27.2 g * mole/58.44277 g = 0.465 moles NaCl 500 ml * L/1000ml = 0.500 L M = 0.465 moles NaCl / 0.500 L = 0.931 M NaCl

Question:During the preparation of tert-butyl chloride from tert-butyl alcohol through nucleophilic substitution, why is anhydrous calcium chloride that has been used and not anhydrous sodium sulphate? Does it has any effects other than drying the product?

Answers:The sodium sulfate *may* react with the chloride in the reaction, making NaCl in the process. If these are all dissolved together and the solution being warm it is more likely. As small a chance as it may be, it is good lab practice to use like ions (chloride and chloride) to get the best yield and lowest contamination with unwanted products. Simply put, using chloride based drying agent won't mess up your reaction even if there are anion transfers from the drying product into the reactants because they are the same anions.

From Youtube

Preparation Of Oxygen Using Potassium Chlorate :Check us out at www.tutorvista.com Potassium chlorate is a compound containing potassium, chlorine and oxygen, with the chemical formula KClO3. In pure form, it is a white crystalline substance. It is the most common chlorate in industrial use, and is usually present in well-stocked laboratories. It is used as an oxidizing agent, to prepare oxygen, as a disinfectant, in safety matches, and in explosives and fireworks. Potassium chlorate ( KCLO3 ) decomposes upon heating to yield Potassium chloride and Oxygen Gas . Manganese dioxide ( MnO2 ) is mixed with Potassium chlorate and performs as the catalyst of this reaction . 2 KClO3 ----( Heat with MnO2 Catalyst )----2 KCl + 3 O2

05 - Make Copper (II) Chloride and Copper (II) Carbonate from Copper metal :In this video I show how to make two compounds of copper with copper in the +2 oxidation state. NOTE: I did not use stoichiometric amounts of reagents. Instructions for copper (II) chloride: Get some copper metal. I got mine from cables, so the purity is relatively high. Now add dilute hydrochloric acid (you can use conc. acid, but be aware that your container will get really hot and your mixture might splash out) and conc. hydrogen peroxide (30% w/v) to initiate the reaction between copper and acid. The peroxide itself is being consumed during the reaction. After the reaction has finished, boil down the mixture to obtain pure copper (II) chloride. Make sure that the reaction has finished; you don't want to boil hydrogen peroxide or this might result to a bad event ;). If your solid product is brown, this is due to dehydration of the monohydrate form. As long as you keep it in a moist place it will get its green color back. Instructions for copper (II) carbonate: Get your already prepared copper chloride and some sodium carbonate. This can be easily prepared by heating sodium bicarbonate to high temperature (over 200 oC). Sodium bicarbonate undergoes decomposition to sodium carbonate. Now dissolve both in two separate beakers and mix them. You 'll notice a blue precipitate of copper (II) carbonate. Just let it settle down, filter to get your product and let it dry. Thanks for watching! Let me know your questions!