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calcium oxide

calcium oxide, chemical compound, CaO, a colorless, cubic crystalline or white amorphous substance. It is also called lime, quicklime, or caustic lime, but commercial lime often contains impurities, e.g., silica, iron, alumina, and magnesia. It is prepared by heating calcium carbonate (e.g., limestone) in a special lime kiln to about 500°C to 600°C, decomposing it into the oxide and carbon dioxide. Calcium oxide is widely used in industry, e.g., in making porcelain and glass; in purifying sugar; in preparing bleaching powder, calcium carbide, and calcium cyanamide; in water softeners; and in mortars and cements. In agriculture it is used for treating acidic soils (liming). It is incandescent when heated to high temperatures; the Drummond light, or limelight, provides a brilliant white light by heating a cylinder of lime with the flame of an oxyhydrogen torch. Calcium oxide is a basic anhydride, reacting with water to form calcium hydroxide; during the reaction (slaking) much heat is given off and the solid nearly doubles its volume.

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calcium oxide

calcium oxide (quicklime) White solid (CaO) made by heating calcium carbonate (CaCO3) at high temperatures. It is used industrially to treat acidic soil and to make porcelain and glass, bleaching powder, caustic soda, mortar and cement. Calcium oxide reacts with water to form calcium hydroxide (Ca(OH)2).

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Calcium oxide

Calcium oxide

Calcium oxide (CaO), more commonly known as lime or quick lime, has been studied by scholars as far back as the pre-Christian era. In his book Historia Naturalis, for example, Pliny the Elder discussed the preparation, properties, and uses of lime. Probably the first scientific paper on the substance was Dr. Joseph Blacks Experiments Upon Magnesia, Alba, Quick-lime, and Some Other Alkaline Substances, written in 1755.

Lime does not occur naturally since it reacts so readily with water (to form hydrated lime) and carbon dioxide (to form limestone). It is produced in very large quantities synthetically, however, by the heating of limestone. For many years, calcium oxide has ranked among the top ten chemicals in the United States in terms of production. Other common names by which the compound is known include burnt lime, unslaked lime, fluxing lime, and calx.

In its pure form, calcium oxide occurs as white crystals, white or gray lumps, or a white granular powder. It has a very high melting point of 4,662°F(2, 572°C) and a boiling point of 5,162°F (2, 850°C). It dissolves in and reacts with water to form calcium hydroxide and is soluble in acids and some organic solvents.

Like other calcium compounds, calcium oxide is used for many construction purposes, as in the manufacture of bricks, mortar, plaster, and stucco. Its high melting point makes it attractive as a refractory material, as in the lining of furnaces. The compound is also used in the manufacture of various types of glass. Common soda-lime glass, for example, contains about 12% calcium oxide, while high-melting aluminosilicate glass contains about 20% calcium oxide. One of the new forms of glass used to coat surgical implants contains an even higher ratio of calcium oxide, about 24% of the compound.

Among the many other applications of calcium oxide are its uses in the production of pulp and paper, in the removal of hair from animal hides, in clarifying cane and beet sugar, in poultry feeds, and as a drilling fluid.

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Calcium Oxide

Calcium oxide

Calcium oxide (CaO), more commonly known as lime or quick lime, has been studied by scholars as far back as the pre-Christian era. In his book Historia Naturalis, for example, Pliny the Elder discussed the preparation, properties, and uses of lime. Probably the first scientific paper on the substance was Dr. Joseph Black's "Experiments Upon Magnesia, Alba, Quick-lime, and Some Other Alkaline Substances," written in 1755.

Lime does not occur naturally since it reacts so readily with water (to form hydrated lime) and carbon dioxide (to form limestone). It is produced in very large quantities synthetically, however, by the heating of limestone. For many years, calcium oxide has ranked among the top ten chemicals in the United States in terms of production. Other common names by which the compound is known include burnt lime, unslaked lime, fluxing lime, and calx.

In its pure form, calcium oxide occurs as white crystals, white or gray lumps, or a white granular powder. It has a very high melting point of 4,662°F (2,572°C) and a boiling point of 5,162°F (2,850°C). It dissolves in and reacts with water to form calcium hydroxide and is soluble in acids and some organic solvents.

Like other calcium compounds, calcium oxide is used for many construction purposes, as in the manufacture of bricks, mortar, plaster, and stucco. Its high melting point makes it attractive as a refractory material, as in the lining of furnaces. The compound is also used in the manufacture of various types of glass . Common soda-lime glass, for example, contains about 12% calcium oxide, while high-melting aluminosilicate glass contains about 20% calcium oxide. One of the new forms of glass used to coat surgical implants contains an even higher ratio of calcium oxide, about 24% of the compound.

Among the many other applications of calcium oxide are its uses in the production of pulp and paper, in the removal of hair from animal hides, in clarifying cane and beet sugar, in poultry feeds, and as a drilling fluid.

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Calcium Oxide

Calcium Oxide

OVERVIEW

Calcium oxide (KAL-see-um OK-side) is an odorless crystalline or powdery solid that, in a pure form, is white to off-gray. It often appears with a yellowish or brownish tint to the presence of impurities, especially iron. Calcium oxide reacts with water to form calcium hydroxide (Ca(OH)2) with the evolution of significant amounts of heat. The compound is strongly caustic.

KEY FACTS

OTHER NAMES:

Lime; quicklime; burnt lime; calx; unslaked lime; fluxing lime

FORMULA:

CaO

ELEMENTS:

Calcium, oxygen

COMPOUND TYPE:

Metallic oxide

STATE:

Solid

MOLECULAR WEIGHT:

56.08 g/mol

MELTING POINT:

2,898°C (5,248°F)

BOILING POINT:

Not available

SOLUBILITY:

Reacts with water to form calcium hydroxide; soluble in acids; insoluble in alcohol and most organic solvents

Calcium oxide has been known since ancient times. The Roman writer Cato the Elder (234–149 bce) described one method of producing the compound in 184 bce. Another early Roman scholar, Pliny the Elder (23–79 ce) discussed the compound at length in his book Historia Naturalis (Natural History), published in 70 ce. By the early fifteenth century, most of Europe was using calcium oxide (widely referred to as lime) in the construction of buildings. The Scottish chemist Joseph Black (1728–1799) carried out some of the earliest scientific studies of calcium oxide. He found that when the compound is exposed to air, it combines with carbon dioxide to form calcium carbonate.

HOW IT IS MADE

The process for making calcium oxide is believed to be one of the first chemical reactions known to humans, dating back to prehistoric times. When limestone (calcium carbonate; CaCO3) is heated, carbon dioxide (CO2) is driven off, leaving calcium oxide behind. The reaction was probably discovered very early in human history because limestone is a common, readily available material in the form of chalk and sea shells, and the amount of heat needed to produce the reaction can easily be produced in a simple wood fire. A more efficient method for carrying out the reaction is to heat the limestone in a kiln (oven) at temperatures of 500°C to 900°C (900°F to 1,600°F), resulting in a more complete conversion of calcium carbonate to calcium oxide. This method is still used today for the commercial preparation of calcium oxide.

COMMON USES AND POTENTIAL HAZARDS

The most important single use of calcium oxide is in metallurgy, particularly in the production of steel. When calcium oxide is added to the furnace in which steel is made, it reacts with sulfur, phosphorus, silica, and other impurities present in the mixture from which steel is produced. The complex mixture that results can be poured off the top of the molten steel in the form of a slag, a nonmetallic waste formed during the production of metals. Calcium oxide plays a comparable role in the manufacture of other metals, such as aluminum and magnesium. About 40 percent of all the calcium oxide produced in the United States goes to metallurgical applications.

The next most important use of calcium oxide is in pollution control devices. Smoke that leaves a factory's smokestack, for example, contains oxides of sulfur and nitrogen that, in the atmosphere, combine with water to form sulfuric acid and nitric acid. To prevent the formation of these pollutants, "scrubbers" can be installed in smokestacks. The scrubbers contain some chemical that reacts with and neutralizes the oxides of sulfur and nitrogen. One of the most common compounds used for this purpose is calcium oxide. About 15 percent of all calcium oxide used in the United States goes to this application.

Interesting Facts

  • Calcium oxide is often used to "lime" lake waters that have been acidified by acid rain. It reacts with and neutralizes acids in the lake formed when nitric and sulfuric acid in acid rain are carried to earth by rain, snow, sleet, and other forms of precipitation.
  • When calcium oxide is heated near its melting point, it gives off a brilliant white light. In the years before electricity was available for lighting, particularly during the second half of the nineteenth century, heated lime was used to produce the bright lights needed to illuminate stage productions. From this practice came the expression "being in the limelight" to refer to anyone who was in public view of large groups of people.
  • Because it was thought to accelerate the decomposition of soft tissue, quicklime has historically been used in the burial of diseased animals and humans. For example, bodies of plague victims in London in 1666 were directed to be buried in quicklime.

Some other uses of calcium oxide include:

  • In water treatment plants, to control acidity of the water being treated and to remove impurities present in the water;
  • In the construction industry, where it is used to make plaster, mortar, stucco, bricks, and other building materials;
  • As a filler to strengthen paper products;
  • As a refractory, a heat-resistant material used to line the insides of furnaces;
  • In the production of other chemical materials;
  • As an additive for poultry feed;
  • In insecticides and fungicides;
  • In the removal of hair from hides before tanning; and
  • As a food additive to maintain proper acidity and give bulk to a food product.

Exposure to calcium oxide can cause damage to the skin, eyes, nose, and respiratory system. People who use the product in their line of work or at home (for garden purposes, for example) must exercise extreme caution to avoid swallowing, breathing in, or otherwise coming into contact with the chemical. If such contact occurs, it should be washed off completely and medical assistance requested.

Words to Know

CAUSTIC
Strongly basic or alkaline material that irritates or corrodes living tissue.
METALLURGY
The science of working with metals.

FOR FURTHER INFORMATION

"Calcium Oxide or Quicklime." Peters Chemical Company. http://www.peterschemical.com/calcium_oxide_or_quicklime.htm (accessed on September 27, 2005).

"Chemical of the Week—Lime." http://scifun.chem.wisc.edu/chemweek/Lime/lime.html (accessed on September 27, 2005).

"Cheminfo." Canadian Centre for Occupational Health and Safety. http://intox.org/databank/documents/chemical/calcoxid/cie11.htm (accessed on September 27, 2005).

"Metallurgy Information Area." National Lime Association. http://www.lime.org/ENV02/Metal802.htm#IS (accessed on September 27, 2005).

See AlsoCalcium Carbonate; Calcium Hydroxide

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