JEE Syllabus
Preparation and properties of the following compounds:
Oxides,
peroxides,
hydroxides,
carbonates,
bicarbonates,
chlorides and
sulphates of
sodium,
potassium,
magnesium and
calcium;
Aluminium: alumina, aluminium chloride and alums;
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MAIN TOPICS IN TMH BOOK
SODIUM AND POTASSIUM
MAGNESIUM AND CALCIUM
ALUMINIUM
IRON
ZINC
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Compounds of Sodium
Sodium oxide
Sodium oxide has formula Na2O.
It is also called sodium(I) oxide, disodium oxide, sodium monoxide, and soda.
It is used in ceramics as a glaze additive.
It is also a constituent of glass at around 15% sodium oxide
Sodium oxide has the formula weight of 61.979 u.
It is formed when sodium is burned with limited oxygen to the following equation:
4Na + O-2 → 2Na-2O
Sodium oxide is a basic compound, thus on reaction with water will create sodium hydroxide (NaOH).
Na-2O + H-2O → 2NaOH
sodium peroxide
A nearly white compound (Na 2 O 2 ), having vigorous oxidizing properties, and used in bleaching mechanical paper pulps and as a final stage in the bleaching of chemical paper pulps in some multi-stage bleaching sequences.
It is dangerous in use because, when in contact with organic matter, it reacts so vigorously with atmospheric moisture that sufficient heat can be generate to cause organic matter to burn or even explode.
sodium hydroxide
sodium hydroxide chemical compound, NaOH, is a white crystalline substance that readily absorbs carbon dioxide and moisture from the air.
It is very soluble in water, alcohol, and glycerin.
It is a caustic and a strong base
It is commonly known as caustic soda, lye, or sodium hydrate.
The principal method for its manufacture is electrolytic dissociation of sodium chloride; chlorine gas is a coproduct.
Small amounts of sodium hydroxide are produced by the soda-lime process in which a concentrated solution of sodium carbonate (soda) is reacted with calcium hydroxide (slaked lime); calcium carbonate precipitates, leaving a sodium hydroxide solution.
The major use of sodium hydroxide is as a chemical and in the manufacture of other chemicals; because it is inexpensive, it is widely used wherever a strong base is needed.
It is also used in producing rayon and other textiles, in making paper, in etching aluminum, in making soaps and detergents, and in a wide variety of other uses.
Sodium carbonate (Na2CO3)
Sodium carbonate exists as anhydrous (Na2CO3) and also as hydrated salt. The decahydrated salt (Na2CO3.10H2O) is known as washing soda while the anhydrous salt is called soda ash.
Occurrence
Large deposits of this salt occur in Owens lake in California and Lake Magadi in British East Africa. It occurs native as Na2CO3.NaHCO3.H2O in Egypt.
During hot weather, soda is also collected from a large number of alkaline lakes.
Manufacture of Sodium Carbonate
Ammonia-soda process (or Solvay process)
This process is the most popularly used method. As Ernest Solvay, the Belgian chemical engineer, devised it in 1864 it is known as Solvay process.
Raw materials
The raw materials for this process are common salt, ammonia and limestone (for supplying CO2 and quicklime).
Principle
When carbon dioxide is passed into a concentrated solution of brine saturated with ammonia, ammonium bicarbonate is produced,
The ammonium bicarbonate then reacts with common salt forming sodium bicarbonate,
Sodium bicarbonate being slightly soluble (in presence of sodium ions) gets precipitated. The precipitated sodium bicarbonate is removed by filtration and changed into sodium carbonate by heating.
The mother liquor remaining after the precipitation of sodium bicarbonate contains ammonium chloride. This is used to regenerate ammonia (one of the raw materials) by steam heating with milk of lime.
Lime is obtained by heating limestone.
Ammonia and carbon dioxide liberated are utilized in making the whole process cyclic and continuous. The only by-product in the process is calcium chloride.
Sodium Bicarbonate NaHCO-3
Sodium Bicarbonate, commonly called baking soda, is a white odourless, crystalline solid, completely soluble in water but slightly soluble in ethanol. It is the mildest of all sodium alkalis.
It is prepared from purified sodium carbonate or sodium hydroxide solution with passing carbon dioxide which is bubbled into the solution of pure carbonate, and the bicarbonate precipitates out to be dried as the bicarbonate is less soluble than the carbonate.
Sodium bicarbonate is also made as an intermediate product in the Solvay process (described above)which is to make sodium carbonate from calcium carbonate by treating sodium chloride with ammonia and carbon dioxide.
The major use of sodium bicarbonate is in baking powders.
Sodium Bicarbonate plays an important role in the products of many diverse industries with functions of releasing CO2 when heated above about 50 C or when reacted with a weak acid makes sodium bicarbonate a key ingredient in food leavening as well as in the manufacture of effervescent salts and beverages.
It can react as an acid or a base in water treatment.
In health and beauty applications, mild abrasivity and ability to reduce odors chemically by neutralizing the acid by-products of bacteria are utilized.
It is also used in treating wool and silk, fire extinguishers, pharmacy, leather, oredressing, metallurgy, in cleaning preparations and industrial & chemical processe.
Uses
food & food processing, beverages , pharmaceuticals , animal foodstuffs , household cleaning products , rubber & plastics foam blowing , fire extinguishers & explosion suppression , effluent & water treatment, flue gas treatment , oil drilling , industrial & chemical processes
Sodium chloride (NaCl) or common salt is an ionic crystal consisting of equal numbers of sodium and chlorine atoms and is an essential component in the human diet, being found in blood sweat and tears.
Occurrence
Sodium chloride is abundant and can be found naturally occurring. It can be found in the mineral halite (pure rock salt) as well as in mixed evaporates in salt lakes.
Sea water also contains 2.7% by weight salt and constitutes 80% of the dissolved minerals in sea water.
Production
Sodium chloride is mined or obtained from brine, when water is added to salt deposits.
Alternatively, it is obtained from sea water. This is commonly known as sea salt and constitutes most table salt. It also contains some impurities.
Sodium chloride:
• Has a cubic crystalline structure
• Is clear when pure, although may also appear white, grey or brownish, depending upon purity
• Is soluble in water
• Is slightly soluble in other liquids
• Is odourless
• Has a characteristic taste
• Molten sodium chloride is an electrical conductor
Symbol NaCl
Atomic Weight 58.44
Eutectic Composition 23.31% NaCl
Melting Point 801°C
Boiling Point 1465°C
Density 2.17g/cm3
Refractive Index 1.5442
Mohs Hardness 2.5
Co-Efficient of Thermal Expansion @ 0°C 40x10-6
Solubility g/100g H2O at 0°C 35.7
Sodium chloride is used for:
• Windows for analytical instruments
• De-icing
• Food and cooking
• High power lasers
• To produce chlorine and sodium
• Historically it has been used as a form of currency
SODIUM SULPHATE
FORMULA Na2SO4
MOL WT. 142.04
SYNONYMS Disodium monosulfate;
PHYSICAL AND CHEMICAL PROPERTIES
PHYSICAL STATE Hygroscopic white powder, Odorless
MELTING POINT 880 - 888 C
BOILING POINT 1100 C (Decomposes)
SPECIFIC GRAVITY 2.66 - 2.75
SOLUBILITY IN WATER Soluble
pH Aqueous solution is neutral
Sodium sulfate is a white, orthorhombic crystalline solid at room temperatures ( a monoclinic structure at > 100 C, a hexagonal structure at > 250C).
It is reduced to sodium sulfide at high temperature.
But sodium sulfate is a stable compound which does not decompose and does not react with oxidising or reducing agents at normal temperatures.
It is neutral (pH of 7) in water.
Sodium sulfate is most soluble in water at 32.4 C (49.7g/100 g).
Commercial major source of sodium sulfate is salt cake (impure sodium sulfate), a by-product of hydrochloric acid production from sodium chloride by treatment with sulfuric acid.
Sodium sulfate is obtained also as a byproduct of rayon production and sodium dichromate production. The decahydrate is known as Glauber's salt.
About half of the world's production is from the natural mineral form of the decahydrate (mirabilite).
Anhydrous sodium sulfate is found in nature as the mineral thenardite (Na2SO4).
Other sodium sulfate minerals are metasideronatrite Na4Fe2(SO4)4(OH)213H2O, krohnkite Na2Cu(SO4)212H2O, and schairerite Na3(SO4)(F,Cl).
Sodium sulfate is consumed in four major categories; powder detergents as a processing aid and as a filler, wood pulp processing for making kraft paper, textile dyeing processes as a levelling agent to penetrate evenly, and molten glass process to remove small air bubbles.
Sodium sulfate is employed also as a raw material for the production of various chemicals.
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Compounds of Potassium
potassium (I) oxide
• Formula as commonly written: K2O
Synonyms
• potassium (I) oxide
• potassium oxide
• dipotassium oxide
Physical properties
• Colour: yellowish white to grey
• Appearance: crystalline solid
• Melting point: >763°C; 350°C (decomposes
• Density: 2350 kg m-3
Potassium Peroxide
Identifications
• Synonyms/Related:
o Dipotassium peroxide
o K2O2
POTASSIUM HYDROXIDE
FORMULA KOH
MOL WT. 56.1
SYNONYMS Potassium hydrate; Caustic potash; Lye;
Potassium Hydroxide, commonly called caustic potash with formula KOH, is a caustic compound of strong alkaline chemical dissolving readily in water, giving off much heat and forming a caustic solution.
It is a white deliquescent solid in the form of pellets obtained by concentration of purified electrolytic potassium hydroxide solution with very low chloride content.
PHYSICAL AND CHEMICAL PROPERTIES
PHYSICAL STATE odorless white, deliquescent solid
MELTING POINT 360 C
BOLING POINT 1320 C
SPECIFIC GRAVITY 2.044
SOLUBILITY IN WATER Soluble
pH 13.5 (0.1 molar solution)
It reacts violently with acid and is corrosive in moist air toward metals such as zinc, aluminium, tin and lead forming a combustible, explosive gas.
It absorbs rapidly carbon dioxide and water from air.
Contact with moisture or water will generate heat.
Sodium hydroxide (Caustic soda) and potassium hydroxide (Caustic potash) are the two most important caustics. They are closely resembles in chemical properties and applications, e.g., in manufacturing liquid soap, in bleaching, and in manufacturing chemicals.
Potassium hydroxide is the largest-volume potassium chemical for non-fertilizer use.
Potassium Hydroxide is used in chemical manufacturing including potassium carbonate and other potassium chemicals, fertilizers, phosphates, agrochemicals, alkaline batteries and dyes.
APPLICATIONS
Potassium Hydroxide is used in chemical manufacturing including potassium carbonate and other potassium chemicals, fertilizers, phosphates, agrochemicals, alkaline batteries and dyes. It is also widely used in soap and bleaching industry.
Potassium carbonate
Potassium carbonate is a white salt, soluble in water (insoluble in alcohol), which forms a strongly alkaline solution. It can be made as the product of potassium hydroxide's absorbent reaction with carbon dioxide. It is deliquescent, often appearing a damp or wet solid. Potassium carbonate is used in the production of soap and glass.
Production
Potassium carbonate is prepared commercially by the electrolysis of potassium chloride. The resulting potassium hydroxide is then carbonated using carbon dioxide to form potassium carbonate, which is often used to produce other potassium compounds.
2KOH + CO2 → K2CO3 + H2O
Uses
Potassium carbonate has been used for soap, glass, and china production.
In the laboratory, it may be used as a mild drying agent where other drying agents such as calcium chloride may be incompatible. However, it is not suitable for acidic compounds.
Mixed with water it causes an exothermic reaction that results in a temperature change, producing heat.
In cuisine, it is used as an ingredient in the production of grass jelly, a food consumed in Chinese and Southeast Asian cuisines.
Potassium carbonate is being used as the electrolyte in many cold fusion experiments.
Potassium bicarbonate (also known as potassium hydrogen carbonate or potassium acid carbonate), is a colorless, odorless, slightly basic, salty substance. The compound is used as a source of carbon dioxide for leavening in baking, extinguishing fire in powder fire extinguishers, acting as a reagent, and a strong buffering agent in medications. The US Food and Drug Administration (FDA) recognizes potassium bicarbonate as "generally recognized as safe". It is used as a base in foods to regulate pH.
Potassium bicarbonate is soluble in water, and is often found added to bottled water to affect taste; however, it is not soluble in alcohol. Decomposition of the substance occurs between 100°C and 120°C into K2CO3 (potassium carbonate), H2O (water), and CO2 (carbon dioxide). In concentrations greater than 0.5%, KHCO3 can have phytotoxic effects on plants (potassium bicarbonate has widespread use in crops, especially for neutralizing acidic soil), although there is no evidence of human carcinogenicity, no adverse effects of overexposure, and no LD50.
Physically, potassium bicarbonate occurs as a crystal or a soft white granular powder. It has a CAS No [298-14-6]. It is manufactured by reacting potassium carbonate with carbon dioxide and water:
K2CO3 + CO2 + H2O → 2 KHCO3
Potassium bicarbonate is used as a fire suppression agent in some dry powder fire extinguishers,. It is about twice as effective in fire suppression as sodium bicarbonate.
potassium chloride (KCl)
Potassium chloride is also commonly known as "Muriate of Potash".
Potash varies in color from pink or red to white depending on the mining and recovery process used. White potash, sometimes referred to as soluble potash, is used primarily for making liquid starter fertilizers.
Manufacture/Extraction
Potassium chloride occurs naturally as sylvite, and in combination with sodium chloride as sylvinite.
It can be extracted from sylvinite.
It is also extracted from salt water .
It is a by-product of the making of nitric acid from potassium nitrate and hydrochloric acid.
Physical properties
In its pure state it is odorless.
It has a white or colorless vitreous crystal, with a crystal structure that cleaves easily in three directions.
Potassium chloride crystals are face-centered cubic.
Potassium chloride has a crystalline structure like many other salts. Structure: face-centered cubic. Lattice Constant: roughly 6.3 angstroms.
In chemistry and physics it is a very commonly used as a standard, for example as a calibration standard solution in measuring electrical conductivity of (ionic) solutions, since carefully prepared KCl solutions have well-reproducible and well-repeatable measurable properties.
Chemical properties
Potassium chloride can react as a source of chloride ion. As with any other soluble ionic chloride, it will precipitate insoluble chloride salts when added to a solution of an appropriate metal ion:
KCl(aq) + AgNO3(aq) → AgCl(s) + KNO3(aq)
Although potassium is more electropositive than sodium, KCl can be reduced to the metal by reaction with metallic sodium at 850 °C because the potassium is removed by distillation
(KCl(l) + Na(l) ⇌ NaCl(l) + K(g)
This method is the main method for producing metallic potassium.
As with other compounds containing potassium, KCl in powdered form gives a lilac flame test result.
KCl is used in medicine, scientific applications, food processing and in judicial execution through lethal injection.
Potassium sulfate (K2SO4)
Potassium sulfate (K2SO4) (also known as potash of sulfur) is a non-flammable white crystalline salt which is soluble in water. The chemical is commonly used in fertilizers, providing both potassium and sulfur.
Manufacture
Potassium sulfate can be synthesised by the decomposition of potassium chloride with sulfuric acid. Hhydrogen chloride evaporates and can be used to produce hydrochloric acid.
The Hargreaves method is basically the same process with different starting materials. Sulfur dioxide, oxygen and water (the starting materials for sulfuric acid) are reacted with potassium chloride. Hydrochloric acid evaporates off.
It is obtained as a by-product in many chemical reactions including the production of nitric acid. This can be done by mixing the following: 2 Parts Potassium Nitrate to 1 Part Sulfuric Acid (molar ratio).
2KNO3 + H2SO4 ---> 2HNO3 + K2SO4
To purify the crude product, it can be dissolved in hot water and then filtered and cooled, causing the bulk of the dissolved salt to crystallize with characteristic promptitude.
Properties
The anhydrous crystals form a double six-sided pyramid, but are in fact classified as rhombic.
They are transparent, very hard and have a bitter, salty taste.
The salt is soluble in water, but insoluble in solutions of potassium hydroxide (sp. gr. 1.35), or in absolute ethanol.
It melts at 1078 °C.
Uses
The principal use of potassium sulfate is as a fertilizer. The crude salt is also used occasionally in the manufacture of glass.
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Compounds of Magnesium
Magnesium oxide
Roasting either magnesium carbonate or magnesium hydroxide produces the oxygen compound magnesium oxide, commonly called magnesia, MgO, a white solid used in the manufacture of high-temperature refractory bricks, electrical and thermal insulators, cements, fertilizer, rubber, and plastics. It is used medically as a laxative.
Magnesium peroxide
Magnesium peroxide is a fine powder peroxide with a white to white-off color. It releases oxygen by breaking down at a controlled rate with a hydrous fluid.
In contact with water it decomposes by the reactions:
MgO2+ 2 H2O → Mg(OH)2 + H2O2
2 H2O2 → 2 H2O + O2
Applications
Magnesium peroxide being environmentally benign and its stable oxygen release are used widely in the cosmetic, agricultural, pharmaceutical, and environmental industries. It is used to reduce contaminant levels in groundwater. Magnesium peroxide is used in the bioremediation of contaminated soil and can improve the soil quality for plant growth and metabolism. It also used in the aquaculture industry for bioremediation.
Commercially, magnesium peroxide exists as a form of compound of magnesium peroxide and magnesium hydroxide
http://en.wikipedia.org/wiki/Magnesium_peroxide
Magnesium hydroxide
Magnesium hydroxide, Mg(OH)2, is a white powder produced in large quantities from seawater by the addition of milk of lime (calcium hydroxide). It is the primary raw material in the production of magnesium metal. In water it forms a suspension known as milk of magnesia, which has long been used as an antacid and a laxative.
Magnesium carbonate
Magnesium carbonate, MgCO3, occurs in nature as the mineral magnesite and is an important source of elemental magnesium. It can be produced artificially by the action of carbon dioxide on a variety of magnesium compounds. The odourless white powder has many industrial uses—e.g., as a heat insulator for boilers and pipes and as an additive in food, pharmaceuticals, cosmetics, rubbers, inks, and glass.
Magnesium chloride,
The action of hydrochloric acid on magnesium hydroxide produces magnesium chloride, MgCl2, a colourless, deliquescent (water-absorbing) substance employed in magnesium metal production, in the manufacture of a cement for heavy-duty flooring, and as an additive in textile manufacture.
Magnesium sulfate
Magnesium sulfate, MgSO-4, is a colourless, crystalline substance formed by the reaction of magnesium hydroxide with sulfur dioxide and air. A hydrate form of magnesium sulfate called kieserite, MgSO-4.H2O, occurs as a mineral deposit.
Synthetically prepared magnesium sulfate is sold as Epsom salt, MgSO-4.7H2O. In industry magnesium sulfate is used in the manufacture of cements and fertilizers and in tanning and dyeing; in medicine it serves as a purgative.
http://www.britannica.com/eb/article-4446/magnesium
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Compounds of Calcium
A number of calcium compounds such as calcium oxide, calcium hydroxide, calcium carbide, calcium chloride, calcium carbonate, calcium sulphate, etc. occur either in nature or are easily prepared. All these compounds find extensive use in our day-to-day living.
Calcium oxide (CaO)
Calcium oxide is commonly known as quicklime, and is a material of primary importance in the building industry.
Preparation
Quicklime (CaO) is prepared by heating very strongly, limestone (CaCO3) in a lime kiln.
Smaller pieces of limestone are introduced from the top and heating is done from the lower end.
Limestone decomposes at about 1000°C to give calcium oxide.
Properties
Calcium oxide (quicklime) is a white amorphous solid. On heating, quicklime (CaO) glows at high temperatures. This glow of white dazzling light is called lime light. Quicklime melts at 2870 K (2597°C).
On exposure to the atmosphere, it absorbs moisture and carbon dioxide to finally give calcium carbonate.
When water is poured over quicklime, a lot of heat is produced giving out steam with a hissing sound. This is called slaking of lime.
Uses
For the manufacture of bleaching powder, glass, calcium carbide, soda ash etc.
For white washing of buildings.
For tanning of leather.
As a fertilizer for acidic soil.
For softening hard water.
In building and construction industry as an important raw material.
Calcium peroxide (CaO2)
Calcium peroxide (CaO2) is a solid peroxide with a white or yellowish color.
When in contact with water it will immediately begin to decompose releasing oxygen.
It will dissolve in acid to form hydrogen peroxide.
Calcium hydroxide (Ca(OH)-2)
Calcium hydroxide in solid powdered form is called slaked lime. A suspension of slaked lime in water is called milk of lime.
A clear saturated solution of Ca(OH)2 is called lime water.
Preparation
From quicklime
Calcium hydroxide is obtained by treating quicklime with water
From calcium chloride
Calcium hydroxide is obtained as white precipitate when sodium hydroxide is added to a concentrated solution of calcium chloride.
Properties
Calcium hydroxide is a white amorphous solid (density: 2.08 g/mL).
Calcium hydroxide is sparingly soluble in water. The solubility decreases with temperature.
With carbon dioxide
When carbon dioxide is passed through lime water, it becomes milky due to formation of calcium carbonate.
The milkiness disappears on passing CO2 gas in excess because calcium carbonate changes to the soluble calcium bicarbonate.
When slaked lime is exposed to air, it absorbs carbon dioxide to form calcium carbonate.
This reaction forms the basis of the white washing of buildings.
With acids
Slaked lime (Ca(OH)2) dissolves in dilute hydrochloric acid.
However, it is not soluble in dilute sulphuric acid because the calcium sulphate formed is a water-insoluble salt.
With chlorine
Slaked lime reacts with chlorine to form bleaching powder.
Uses
As a building material.
For white washing buildings.
For the softening of hard water.
As lime water, it is used for the detection of carbon dioxide.
In tanning industry.
Calcium carbonate (CaCO3)
Occurrence
The most abundant mineral of calcium is calcium carbonate. It occurs in nature in different forms, such as limestone, marble, chalk etc.
Calcium carbonate occurs abundantly as dolomite, MgCO3.CaCO3, a mixture of calcium and magnesium carbonates. It is the chief constituent of shells of sea animals and also of bones along with tricalcium phosphate.
Preparation
Laboratory preparation
Calcium carbonate is prepared in the laboratory by passing carbon dioxide gas into lime water.
Calcium carbonate is also obtained by adding the solution of a soluble carbonate to soluble calcium salt e.g.,
The resulting precipitate is filtered, washed and dried. The product obtained is known as precipitated chalk.
Excess of carbon dioxide should be avoided since this leads to the formation of calcium hydrogen carbonate.
Properties
Calcium carbonate is a white fluffy powder. It is almost insoluble in water.
Action of heat
When heated to 1200 K, it decomposes to give quicklime and carbondioxide.
Action of heat
When heated to 1200 K, it decomposes to give quicklime and carbondioxide.
With acids
Calcium carbonate reacts with dilute acids to liberate carbon dioxide.
Uses
As a building material in the form of marble.
In the manufacture of quick lime.
As a raw material for the manufacture of sodium carbonate in Solvay process.
In the extraction of metals such as iron (as flux).
As a constituent of toothpaste.
Calcium sulphate or Gypsum (CaSO4)
Occurrence
In nature calcium sulphate occurs as
Anhydrite, CaSO4
Gypsum, CaSO4.2H2O
Large quantities of gypsum are available in India in the states of Punjab and Rajasthan.
Preparation
In laboratory, calcium sulphate is prepared by the action of sulphuric acid on calcium chloride, calcium oxide or calcium carbonate.
Calcium sulphate is obtained as dihydrate (CaSO4.2H2O) from the solution.
From gypsum
Calcium sulphate can be obtained by heating gypsum above 200°C.
Properties
Calcium sulphate exists as a white solid in two forms:
Dihydrate (CaSO4.2H2O)
Anhydrous (CaSO4)
It is sparingly soluble in water. The solubility increases upto 40°C, beyond which the solubility decreases.
Action of heat
During initial heating of calcium sulphate dihydrate, the crystal structure
changes. On further heating at 390 K, it loses water and forms CaSO4.5H2O or (CaSO4)2.H2O. This is known at Plaster of Paris. On heating above 437 K, it becomes anhydrous CaSO4 and does not set when mixed with water. It is known as Dead Burnt Plaster.
Above 1475 K (< 1200°C) it decomposes to give CaO.
With calcium sulphide
When heated with calcium sulphide at 1475 K (1200°C), it gives calcium oxide, and sulphur dioxide.
Uses
For making plaster of paris and ammonium sulphate.
Gypsum (CaSO)4.2H2O) is used as a fertilizer as well as in the manufacture of cement.
For manufacturing ammonium sulphate fertilizer.
As a drying agent.
Preparation of black board chalks.
Plaster of Paris [CaSO4.1/2H2O]
Calcium sulphate with half a molecule of water per molecule of the salt (hemi-hydrate) is called plaster of paris.
Preparation
It is prepared by heating gypsum (CaSO4.2H2O) at 120°C in rotary kilns, where it gets partially dehydrated.
The temperature should be kept below 140°C otherwise further dehydration will take place and the setting property of the plaster will be partially reduced.
Properties
It is a white powder. When mixed with water (1/3 of its mass), it evolves heat and quickly sets to a hard porous mass within 5 to 15 minutes. During setting, a slight expansion (about 1%) in volume occurs so that it fills the mould completely and takes a sharp impression. The process of setting occurs as follows:
The first step is called the setting stage, and the second, the hardening stage. The setting of Plaster of Paris is catalyzed by sodium chloride, while it is reduced by borax, or alum.
Uses
In surgery for setting broken or fractured bones.
For making casts for statues, in dentistry, for surgical instruments, and toys etc.
In making black board chalks, and statues.
In construction industry.
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Compounds of Aluminium
Alumina
Production
Aluminium oxide, also known as alumina, is the main component of bauxite, the principal ore of aluminium. The largest manufacturers in the world of alumina are Alcoa, Alcan and Rusal.[citation needed] Companies which specialise in the production of speciality aluminium oxides and aluminium hydroxides include Alcan and Almatis. The bauxite ore is made up of impure Al2O3, Fe2O3, and SiO2. Bauxite is purified by the Bayer process:
Al2O3 + 3 H2O + 2 NaOH → 2NaAl(OH)4
The Fe2O3 does not dissolve in the base. The SiO2 dissolves as silicate Si(OH)62-. Upon filtering, Fe2O3 is removed. When the Bayer liquor is cooled, Al(OH)3 precipitates, leaving the silicates in solution. The mixture is then calcined (heated strongly) to give aluminium oxide:[3]
2Al(OH)3 + heat → Al2O3 + 3H2O
The formed Al-2O-3 is alumina.
Uses
Alumina output in 2005Annual world production of alumina is approximately 45 million tonnes, over 90% of which is used in the manufacture of aluminium metal.[3]. The major uses of speciality aluminium oxides are in refractories, ceramics, polishing and abrasive applications. Large tonnages are also used in the manufacture of zeolites, coating titania pigments and as a fire retardant/smoke suppressant.
In lighting and photography, alumina is a medium for chromatography, available in basic (pH 9.5), acidic (pH 4.5 when in water) and neutral formulations. In 1961, GE developed "Lucalox", a transparent alumina used in sodium vapor lamps.[citation needed] Aluminium oxide is also used in preparation of coating suspensions in compact fluorescent lamps.
Health and medical applications include it as a material in hip replacements,[3] in water filters (derived water treatment chemicals such as aluminium sulfate, aluminium chlorohydrate and sodium aluminate, are one of the few methods available to filter water-soluble fluorides out of water), and even in toothpaste formulations.
Aluminium oxide is also used for its strength. Most pre-finished wood flooring now uses aluminium oxide as a hard protective coating. In 2004, 3M developed a technique for making a ceramic composed of aluminium oxide and rare earth elements to produce a strong glass called transparent alumina. Alumina can be grown as a coating on aluminium by anodising or by plasma electrolytic oxidation (see the "Properties" section, above). Both its strength and abrasive characteristics are due to aluminium oxide's great hardness (position 9 on the Mohs scale of mineral hardness).
It is widely used as a coarse or fine abrasive, including as a much less expensive substitute for industrial diamond. Many types of sandpaper use aluminium oxide crystals. In addition, its low heat retention and low specific heat make it widely used in grinding operations, particularly cutoff tools. As the powdery abrasive mineral aloxite, it is a major component, along with silica, of the cue tip "chalk" used in billiards. (See William A. Spinks, cue chalk co-inventor, for more information.) Aluminium oxide powder is used in some CD/DVD polishing and scratch-repair kits. Its polishing qualities are also behind its use in toothpaste.
Aluminium oxide is widely used in the fabrication of superconducting devices, particularly single electron transistors and superconducting quantum interference devices (SQUID), where it is used to form highly resistive quantum tunnelling barriers.
http://en.wikipedia.org/wiki/Aluminium_oxide
Aluminium chloride
Preparation
Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride.[1]
2 Al + 3 Cl2 → 2 AlCl3
2 Al + 6 HCl → 2 AlCl3 + 3 H2
Hydrated forms are prepared by dissolving aluminium oxides with hydrochloric acid.
The solid has a low melting and boiling point, and is covalently bonded. It sublimes at 178 °C. Molten AlCl3 conducts electricity poorly, unlike more ionic halides such as sodium chloride. It exists in the solid state as a six-coordinate layer lattice.
Aluminium chloride is a powerful Lewis acid, capable of forming stable Lewis acid-base adducts with even weak Lewis bases such as benzophenone or mesitylene.[3] Not surprisingly it forms AlCl4− in the presence of chloride ion.
In water, partial hydrolysis forms HCl gas or H3O+, as described in the overview above. Aqueous solutions behave similarly to other aluminium salts containing hydrated Al3+ ions, giving a gelatinous precipitate of aluminium hydroxide upon reaction with the correct quantity of aqueous sodium hydroxide:
AlCl3(aq) + 3 NaOH(aq) → Al(OH)3(s) + 3NaCl(aq)
http://en.wikipedia.org/wiki/Aluminium_chloride
Alum
Alum is a salt that in chemistry is a combination of an alkali metal, such as sodium, potassium, or ammonium and a trivalent metal, such as aluminum, iron, or chromium.
The most common form, potassium aluminum sulfate, or potash alum, is one form that has been used in food processing. Another, sodium aluminum sulfate, is an ingredient in commercially produced baking powder.
The potassium-based alum has been used to produce crisp cucumber and watermelon-rind pickles.
Alum, refers to a specific chemical compound and a class of chemical compounds.
The specific compound is the hydrated aluminum potassium sulfate with the formula KAl(SO4)2.12H2O.
The class of compounds known as alums have the related stoichiometry, AB(SO4)2.12H2O.
Production
Alum from alunite
In order to obtain alum from alunite, it is calcined and then exposed to the action of air for a considerable time. During this exposure it is kept continually moistened with water, so that it ultimately falls to a very fine powder. This powder is then lixiviated with hot water, the liquor decanted, and the alum allowed to crystallize. The alum schists employed in the manufacture of alum are mixtures of iron pyrite, aluminium silicate and various bituminous substances, and are found in upper Bavaria, Bohemia, Belgium, and Scotland. These are either roasted or exposed to the weathering action of the air. In the roasting process, sulfuric acid is formed and acts on the clay to form aluminium sulfate, a similar condition of affairs being produced during weathering. The mass is now systematically extracted with water, and a solution of aluminium sulfate of specific gravity 1.16 is prepared. This solution is allowed to stand for some time (in order that any calcium sulfate and basic ferric sulfate may separate), and is then evaporated until ferrous sulfate crystallizes on cooling; it is then drawn off and evaporated until it attains a specific gravity of 1.40. It is now allowed to stand for some time, decanted from any sediment, and finally mixed with the calculated quantity of potassium sulfa te (or if ammonium alum is required, with ammonium sulfate), well agitated, and the alum is thrown down as a finely-divided precipitate of alum meal. If much iron should be present in the shale then it is preferable to use potassium chloride in place of potassium sulfate.
Alum from clays or bauxite
In the preparation of alum from clays or from bauxite, the material is gently calcined, then mixed with sulfuric acid and heated gradually to boiling; it is allowed to stand for some time, the clear solution drawn off and mixed with acid potassium sulfate and allowed to crystallize. When cryolite is used for the preparation of alum, it is mixed with calcium carbonate and heated. By this means, sodium aluminate is formed; it is then extracted with water and precipitated either by sodium bicarbonate or by passing a current of carbon dioxide through the solution. The precipitate is then dissolved in sulfuric acid, the requisite amount of potassium sulfate added and the solution allowed to crystallize.
Types of alum
Soda alum
Sodium alum, Na2SO4·Al2(SO4)3·24H2O, occurs in nature as the mineral mendozite. It is very soluble in water, and is extremely difficult to purify. In the preparation of this salt, it is preferable to mix the component solutions in the cold, and to evaporate them at a temperature not exceeding 60 °C. 100 parts of water dissolve 110 parts of sodium alum at 0 °C, and 51 parts at 16 °C. Soda alum is used in the acidulent of food as well as in the manufacture of baking powder.
Ammonium alum
Ammonia alum, NH4Al(SO4)2·12H2O, a white crystalline double sulfate of aluminium, is used in water purification, in vegetable glues, in porcelain cements, in natural deodorants (though potassium alum is more commonly used), in tanning, dyeing and in fireproofing textiles.
Uses
Alum in Makeup: Alum was often used as a base in skin whiteners and treatments during the late 16th Century in the Elizabethan fashion.
Shaving alum: is a powdered form of alum used as an astringent to prevent bleeding from small shaving cuts. The styptic pencils sold for this purpose contain aluminium sulfate or potassium aluminium sulfate. Similar products are also used on animals to prevent bleeding after nail-clipping. Alum in block form (usually potassium alum) is used as an aftershave, rubbed over the wet freshly shaved face.
Hair Stiffener: Alum was used in rock form in the 1950's to rub on the front short hair of a "crewcut". When the hair dried, it would stay up all day.
Crystal deodorant: Alum was used in the past as a natural underarm deodorant in Europe, Mexico, Thailand, the Far East and in the Philippines where it is called Tawas. It is now commercially sold for this purpose in many countries, often in a plastic case that protects the crystal and makes it resemble other non-liquid deodorants. Typically potassium alum is used.
Alum powder, found amongst spices at most grocery stores, is used in pickling recipes as a preservative, to maintain crispness, and as an ingredient in some play dough recipes. It is also commonly cited as a home remedy or pain relief for canker sores.
Fire retardant: By soaking and then drying cloth and paper materials they can be made fireproof.
Wax: Alum is used in the Middle East as a component in wax, compounded with other ingredients to create a hair-removal substance.
Foamite: Alum is used to make foamite which is used in many fire extinguishers for chemical and oil fires.
Adjuvant: Alum is used regularly as an adjuvant (enhances immune response to a given immunogen when given with it) in human immunizations.
Antibacterial agent: Alum works as a deodorant because Alum inhibits bacterial growth. This fits the definition of an antibacterial agent. Styptic pencils or Alum powder/crystals can be applied to cuts that have a mild infection.
http://en.wikipedia.org/wiki/Alum
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web sites
thermal decomposition of carbonates, hydroxides and nitrates
http://www.docbrown.info/page01/ExIndChem/ExIndChem.htm-----------------
JEE Question 2007 paper II
Statement - 1
Alkali metals dissolve in liquid ammonia to give blue solutions.
Because
Statement - 2
Alkali metals in liquid ammonia give solvated species of the type [M(NH-3)-n]^+ (M= alkali metals.
(A) Statement – 1 is True, Statement – 2 is True; Statement – 2 is a correct explanation for
statement – 1
(B) Statement – 1 is True, Statement – 2 is True; Statement – 2 is Not a correct explanation for
Statement – 1.
(C) Statement – 1 is True, Statement – 2 is False
(D) Statement – 1 is False, Statement – 2 is True
Solution: B
Blue colour appears due to solvated electrons in liquid ammonia.
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