Calcium oxide

Calcium oxide (CaO), commonly known as quicklime or burnt lime, is a widely used chemical compound. It is a white, caustic, alkaline, crystalline solid at room temperature. The broadly used term "lime" connotes calcium-containing inorganic materials, in which carbonates, oxides and hydroxides of calcium, silicon, magnesium, aluminium, and iron predominate. By contrast, quicklime specifically applies to the single chemical compound calcium oxide. Calcium oxide that survives processing without reacting in building products such as cement is called free lime.[5]

Quicklime is relatively inexpensive. Both it and a chemical derivative (calcium hydroxide, of which quicklime is the base anhydride) are important commodity chemicals.

Calcium oxide
Calcium oxide
Calcium oxide powder
IUPAC name
Calcium oxide
Other names
Quicklime, burnt lime, unslaked lime, pebble lime, calcia
3D model (JSmol)
ECHA InfoCard 100.013.763
EC Number
  • 215-138-9
E number E529 (acidity regulators, ...)
RTECS number
  • EW3100000
UN number 1910
Molar mass 56.0774 g/mol
Appearance White to pale yellow/brown powder
Odor Odorless
Density 3.34 g/cm3[1]
Melting point 2,613 °C (4,735 °F; 2,886 K)[1]
Boiling point 2,850 °C (5,160 °F; 3,120 K) (100 hPa)[2]
Reacts to form calcium hydroxide
Solubility in Methanol Insoluble (also in diethyl ether, octanol)
Acidity (pKa) 12.8
−15.0×10−6 cm3/mol
Cubic, cF8
40 J·mol−1·K−1[3]
−635 kJ·mol−1[3]
QP53AX18 (WHO)
Safety data sheet
GHS pictograms GHS05: CorrosiveGHS07: Harmful
GHS signal word Danger
H302, H314, H315, H318, H335
P260, P261, P264, P270, P271, P280, P301+312, P301+330+331, P302+352, P303+361+353, P304+340, P305+351+338, P310, P312, P321, P330, P332+313, P362, P363, P403+233, P405, P501
NFPA 704
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acidNFPA 704 four-colored diamond
Flash point Non-flammable [4]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 5 mg/m3[4]
REL (Recommended)
TWA 2 mg/m3[4]
IDLH (Immediate danger)
25 mg/m3[4]
Related compounds
Other anions
Calcium sulfide
Calcium hydroxide
Other cations
Beryllium oxide
Magnesium oxide
Strontium oxide
Barium oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).


Calcium oxide is usually made by the thermal decomposition of materials, such as limestone or seashells, that contain calcium carbonate (CaCO3; mineral calcite) in a lime kiln. This is accomplished by heating the material to above 825 °C (1,517 °F),[6] a process called calcination or lime-burning, to liberate a molecule of carbon dioxide (CO2), leaving quicklime.

CaCO3(s) → CaO(s) + CO2(g)

The quicklime is not stable and, when cooled, will spontaneously react with CO2 from the air until, after enough time, it will be completely converted back to calcium carbonate unless slaked with water to set as lime plaster or lime mortar.

Annual worldwide production of quicklime is around 283 million tonnes. China is by far the world's largest producer, with a total of around 170 million tonnes per year. The United States is the next largest, with around 20 million tonnes per year.[7]

Approximately 1.8 t of limestone is required per 1.0 t of quicklime. Quicklime has a high affinity for water and is a more efficient desiccant than silica gel. The reaction of quicklime with water is associated with an increase in volume by a factor of at least 2.5.[8]


A demonstration of slaking of quicklime as a strongly exothermic reaction. Drops of water are added to pieces of quicklime. After a while, a pronounced exothermic reaction occurs ('slaking of lime'). The temperature can reach up to some 300 °C (572 °F).
  • The major use of quicklime is in the basic oxygen steelmaking (BOS) process. Its usage varies from about 30 to 50 kilograms (65–110 lb) per ton of steel. The quicklime neutralizes the acidic oxides, SiO2, Al2O3, and Fe2O3, to produce a basic molten slag.[8]
  • Ground quicklime is used in the production of aerated concrete blocks, with densities of ca. 0.6–1.0 g/cm3 (9.8–16.4 g/cu in).[8]
  • Quicklime and hydrated lime can considerably increase the load carrying capacity of clay-containing soils. They do this by reacting with finely divided silica and alumina to produce calcium silicates and aluminates, which possess cementing properties.[8]
  • Small quantities of quicklime are used in other processes; e.g., the production of glass, calcium aluminate cement, and organic chemicals.[8]
  • Heat: Quicklime releases Thermal energy by the formation of the hydrate, calcium hydroxide, by the following equation:[9]
CaO (s) + H2O (l) ⇌ Ca(OH)2 (aq) (ΔHr = −63.7 kJ/mol of CaO)
As it hydrates, an exothermic reaction results and the solid puffs up. The hydrate can be reconverted to quicklime by removing the water by heating it to redness to reverse the hydration reaction. One litre of water combines with approximately 3.1 kilograms (6.8 lb) of quicklime to give calcium hydroxide plus 3.54 MJ of energy. This process can be used to provide a convenient portable source of heat, as for on-the-spot food warming in a self-heating can, cooking, and heating water without open flames. Several companies sell cooking kits using this heating method.[10]
  • It is known as a food additive to the FAO as an acidity regulator, a flour treatment agent and as a leavener.[11] It has E number E529.
  • Light: When quicklime is heated to 2,400 °C (4,350 °F), it emits an intense glow. This form of illumination is known as a limelight, and was used broadly in theatrical productions before the invention of electric lighting.[12]
  • Cement: Calcium oxide is a key ingredient for the process of making cement.
  • As a cheap and widely available alkali. About 50% of the total quicklime production is converted to calcium hydroxide before use. Both quick- and hydrated lime are used in the treatment of drinking water.[8]
  • Petroleum industry: Water detection pastes contain a mix of calcium oxide and phenolphthalein. Should this paste come into contact with water in a fuel storage tank, the CaO reacts with the water to form calcium hydroxide. Calcium hydroxide has a high enough pH to turn the phenolphthalein a vivid purplish-pink color, thus indicating the presence of water.
  • Paper: Calcium oxide is used to regenerate sodium hydroxide from sodium carbonate in the chemical recovery at Kraft pulp mills.
  • Plaster: There is archeological evidence that Pre-Pottery Neolithic B humans used limestone-based plaster for flooring and other uses.[13][14][15] Such Lime-ash floor remained in use until the late nineteenth century.
  • Chemical or power production: Solid sprays or slurries of calcium oxide can be used to remove sulfur dioxide from exhaust streams in a process called flue-gas desulfurization.
  • Mining: Compressed lime cartridges exploit the exothermic properties of quicklime to break rock. A shot hole is drilled into the rock in the usual way and a sealed cartridge of quicklime is placed within and tamped. A quantity of water is then injected into the cartridge and the resulting release of steam, together with the greater volume of the residual hydrated solid, breaks the rock apart. The method does not work if the rock is particularly hard.[16][17][18]
  • Disposal of corpses: historically, it was believed that quicklime was efficacious in accelerating the decomposition of corpses. This was quite mistaken, and the application of quicklime can even promote preservation; although it can help eradicate the stench of decomposition, which may have led people to suppose it was the actual flesh which had been consumed.[19]


In 80 BC, the Roman general Sertorius deployed choking clouds of caustic lime powder to defeat the Characitani of Hispania, who had taken refuge in inaccessible caves.[20] A similar dust was used in China to quell an armed peasant revolt in 178 AD, when lime chariots equipped with bellows blew limestone powder into the crowds.[21]

Quicklime is also thought to have been a component of Greek fire. Upon contact with water, quicklime would increase its temperature above 150 °C (302 °F) and ignite the fuel.[22]

David Hume, in his History of England, recounts that early in the reign of Henry III, the English Navy destroyed an invading French fleet by blinding the enemy fleet with quicklime.[23] Quicklime may have been used in medieval naval warfare – up to the use of "lime-mortars" to throw it at the enemy ships.[24]


Limestone is a substitute for lime in many applications, such as agriculture, fluxing, and sulfur removal. Limestone, which contains less reactive material, is slower to react and may have other disadvantages compared with lime, depending on the application; however, limestone is considerably less expensive than lime. Calcined gypsum is an alternative material in industrial plasters and mortars. Cement, cement kiln dust, fly ash, and lime kiln dust are potential substitutes for some construction uses of lime. Magnesium hydroxide is a substitute for lime in pH control, and magnesium oxide is a substitute for dolomitic lime as a flux in steelmaking.[25]


Because of vigorous reaction of quicklime with water, quicklime causes severe irritation when inhaled or placed in contact with moist skin or eyes. Inhalation may cause coughing, sneezing, labored breathing. It may then evolve into burns with perforation of the nasal septum, abdominal pain, nausea and vomiting. Although quicklime is not considered a fire hazard, its reaction with water can release enough heat to ignite combustible materials.[26]


  1. ^ a b Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.55. ISBN 1439855110.
  2. ^ Calciumoxid Archived 2013-12-30 at the Wayback Machine. GESTIS database
  3. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A21. ISBN 0-618-94690-X.
  4. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0093". National Institute for Occupational Safety and Health (NIOSH).
  5. ^ "free lime" Archived 2017-12-09 at the Wayback Machine.
  6. ^ Merck Index of Chemicals and Drugs, 9th edition monograph 1650
  7. ^ Miller, M. Michael (2007). "Lime". Minerals Yearbook (PDF). U.S. Geological Survey. p. 43.13.
  8. ^ a b c d e f Tony Oates (2007), "Lime and Limestone", Ullmann's Encyclopedia of Industrial Chemistry (7th ed.), Wiley, pp. 1–32, doi:10.1002/14356007.a15_317, ISBN 3527306730
  9. ^ Collie, Robert L. "Solar heating system" U.S. Patent 3,955,554 issued May 11, 1976
  10. ^ Gretton, Lel. "Lime power for cooking - medieval pots to 21st century cans". Old & Interesting. Retrieved 13 February 2018.
  11. ^ "Compound Summary for CID 14778 - Calcium Oxide". PubChem.
  12. ^ Gray, Theodore (September 2007). "Limelight in the Limelight". Popular Science: 84.
  13. ^ Neolithic man: The first lumberjack?. (August 9, 2012). Retrieved on 2013-01-22.
  14. ^ Karkanas, P.; Stratouli, G. (2011). "Neolithic Lime Plastered Floors in Drakaina Cave, Kephalonia Island, Western Greece: Evidence of the Significance of the Site". The Annual of the British School at Athens. 103: 27. doi:10.1017/S006824540000006X.
  15. ^ Connelly, Ashley Nicole (May 2012) Analysis and Interpretation of Neolithic Near Eastern Mortuary Rituals from a Community-Based Perspective. Baylor University Thesis, Texas
  16. ^ Walker, Thomas A (1888). The Severn Tunnel Its Construction and Difficulties. London: Richard Bentley and Son. p. 92.
  17. ^ "Scientific and Industrial Notes". Manchester Times. Manchester, England: 8. 13 May 1882.
  18. ^ US Patent 255042, 14 March 1882
  19. ^ Schotsmans, Eline M.J.; Denton, John; Dekeirsschieter, Jessica; Ivaneanu, Tatiana; Leentjes, Sarah; Janaway, Rob C.; Wilson, Andrew S. (April 2012). "Effects of hydrated lime and quicklime on the decay of buried human remains using pig cadavers as human body analogues". Forensic Science International. 217 (1–3): 50–59. doi:10.1016/j.forsciint.2011.09.025.
  20. ^ Plutarch, "Sertorius 17.1–7", Parallel Lives.
  21. ^ Adrienne Mayor (2005), "Ancient Warfare and Toxicology", in Philip Wexler (ed.), Encyclopedia of Toxicology, 4 (2nd ed.), Elsevier, pp. 117–121, ISBN 0-12-745354-7
  22. ^ Croddy, Eric (2002). Chemical and biological warfare: a comprehensive survey for the concerned citizen. Springer. p. 128. ISBN 0-387-95076-1.
  23. ^ David Hume (1756). History of England. I.
  24. ^ Sayers, W. (2006). "The Use of Quicklime in Medieval Naval Warfare". The Mariner's Mirror. Volume 92. Issue 3. pp. 262–269.
  25. ^
  26. ^ CaO MSDS.

External links

Agricultural lime

Agricultural lime, also called aglime, agricultural limestone, garden lime or liming, is a soil additive made from pulverized limestone or chalk. The primary active component is calcium carbonate. Additional chemicals vary depending on the mineral source and may include calcium oxide. Unlike the types of lime called quicklime (calcium oxide) and slaked lime (calcium hydroxide), powdered limestone does not require lime burning in a lime kiln; it only requires milling.

The effects of agricultural lime on soil are:

it increases the pH of acidic soil (the lower the pH the more acidic the soil); in other words, soil acidity is reduced and alkalinity increased

it provides a source of calcium and magnesium for plants

it permits improved water penetration for acidic soils

it improves the uptake of major plant nutrients (nitrogen, phosphorus, and potassium) of plants growing on acid soils.Lime may occur naturally in some soils but may require addition of sulfuric acid for its agricultural benefits to be realized. Gypsum is also used to supply calcium for plant nutrition. The concept of "corrected lime potential" to define the degree of base saturation in soils became the basis for procedures now used in soil testing laboratories to determine the "lime requirement" of soils.Other forms of lime have common applications in agriculture and gardening, including dolomitic lime and hydrated lime. Dolomitic lime may be used as a soil input to provide similar effects as agricultural lime, while supplying magnesium in addition to calcium. In livestock farming, hydrated lime can be used as a disinfectant measure, producing a dry and alkaline environment in which bacteria do not readily multiply. In horticultural farming it can be used as an insect repellent, without causing harm to the pest or plant.

Spinner-style lime spreaders are generally used to spread agricultural lime on fields.

Agricultural lime is injected into coal burners at power plants to reduce the pollutants such as NO2 and SO2 from the emissions.

Bone ash

Bone ash is a white material produced by the calcination of bones. Typical bone ash consists of about 55.82% calcium oxide, 42.39% phosphorus pentoxide, and 1.79% water. The exact composition of these compounds varies depending upon the type of bones being used, but generally the formula for bone ash is: Ca5(OH)(PO4)3. Bone ash usually has a density around 3.10 g/mL and a melting point of 1670 °C (3038 °F). Most bones retain their cellular structure through calcination.

Calc-alkaline magma series

The calc-alkaline magma series is one of two main subdivisions of the subalkaline magma series, the other subalkaline magma series being the tholeiitic. A magma series is a series of compositions that describes the evolution of a mafic magma, which is high in magnesium and iron and produces basalt or gabbro, as it fractionally crystallizes to become a felsic magma, which is low in magnesium and iron and produces rhyolite or granite. Calc-alkaline rocks are rich in alkaline earths (magnesia and calcium oxide) and alkali metals and make up a major part of the crust of the continents.

The diverse rock types in the calc-alkaline series include volcanic types such as basalt, andesite, dacite, rhyolite, and also their coarser-grained intrusive equivalents (gabbro, diorite, granodiorite, and granite). They do not include silica-undersaturated, alkalic, or peralkaline rocks.

Calcium hydroxide

Calcium hydroxide (traditionally called slaked lime) is an inorganic compound with the chemical formula Ca(OH)2. It is a colorless crystal or white powder and is produced when quicklime (calcium oxide) is mixed, or slaked with water. It has many names including hydrated lime, caustic lime, builders' lime, slack lime, cal, or pickling lime. Calcium hydroxide is used in many applications, including food preparation, where it has been identified as E number E526. Limewater is the common name for a saturated solution of calcium hydroxide.

Calcium permanganate

Calcium permanganate is an oxidizing agent and chemical compound with the chemical formula Ca(MnO4)2. It consists of the metal calcium and two permanganate ions. It is noncombustible, but, being a strong oxidizing agent, it will accelerate the burning of combustible material. If the combustible material is finely divided, the resulting mixture may be explosive. Contact with liquid combustible materials may result in spontaneous ignition. Contact with sulfuric acid may cause fires or explosions. Mixtures with acetic acid or acetic anhydride can explode if not kept cold. Explosions can occur when mixtures of calcium permanganate and sulfuric acid come into contact with benzene, carbon disulfide, diethyl ether, ethyl alcohol, petroleum, or other organic matter.

It is prepared from the reaction of potassium permanganate with calcium chloride or from the reaction of aluminium permanganate with calcium oxide. It can be also prepared by reacting manganese dioxide with a solution of calcium hypochlorite and a little bit of calcium hydroxide to increase the pH level. If manganese dioxide is heated with calcium hydroxide with an oxidier such as Ca(NO3)2, Ca(ClO3)2, or Ca(ClO4)2, it will produce calcium manganate or mangamite ('hypomanganate').


Calx is a substance formed from an ore or mineral that has been heated.Calx, especially of a metal, is now known as an oxide. According to the obsolete phlogiston theory, the calx was the true elemental substance, having lost its phlogiston in the process of combustion."Calx" is also sometimes used in older texts on artist's techniques to mean calcium oxide.

Corrosive substance

A corrosive substance is one that will damage or destroy other substances with which it comes into contact by means of a chemical reaction.


Degrees of general hardness (dGH or °GH) is a unit of water hardness, specifically of general hardness. General hardness is a measure of the concentration of divalent metal ions such as calcium (Ca2+) and magnesium (Mg2+) per volume of water. Specifically, 1 dGH is defined as 10 milligrams (mg) of calcium oxide (CaO) per litre of water. Since CaO has a molar mass of 56.08 g/mol, 1 dGH is equivalent to 0.17832 mmol per litre of elemental calcium and/or magnesium ions.

In water testing, paper strips often measure hardness in parts per million (ppm), where one part per million is defined as one milligram of calcium carbonate (CaCO3) per litre of water. Consequently, 1 dGH corresponds to 10 ppm CaO but 17.848 ppm CaCO3 which has a molar mass of 100.09 g/mol.

Heart Mountain (Wyoming)

Heart Mountain is an 8,123-foot (2,476 m) klippe just north of Cody in the U.S. state of Wyoming, rising from the floor of the Bighorn Basin. The mountain is composed of limestone and dolomite of Ordovician through Mississippian age (about 500 to 350 million years old), but it rests on the Willwood Formation, rocks that are about 55 million years old—rock on the summit of Heart Mountain is thus almost 300 million years older than the rocks at the base. For over one hundred years, geologists have tried to understand how these older rocks came to rest on much younger strata.

The carbonate rocks that form Heart Mountain were deposited on a basement of ancient (more than 2.5 billion years old) granite when the area was covered by a large shallow tropical sea. Up until 50 million years ago, these rocks lay about 25 miles (40 kilometers) to the northwest, where the eastern Absaroka Range now stands.

Between 75 and 50 million years ago, a period of mountain-building called the Laramide Orogeny caused uplift of the Beartooth Range and subsidence of the Bighorn and Absaroka Basins. Just south of the Beartooth Range, this orogeny uplifted an elongate, somewhat lower plateau which sloped gently to the southeast toward the Bighorn Basin and to the south toward the Absaroka Basin. Immediately following this period of mountain-building, volcanic eruptions began to form the now extinct volcanoes of the Absaroka Range that lie to the south of the Beartooths and extend into Yellowstone National Park. Between 50 and 48 million years ago a sheet of rock about 500 square miles (1,300 square kilometers) in area detached from the plateau south of the Beartooths and slid tens of kilometers to the southeast and south into the Bighorn and Absaroka Basins. This sheet, consisting of Ordovician through Mississippian carbonate rocks and overlying Absaroka volcanic rocks, was probably originally about 4–5 kilometers thick. Although the slope was less than 2 degrees, the front of the landslide traveled at least 25 miles (40 km) and the slide mass ended up covering over 1,300 square miles (>3,400 km2). This is by far the largest rockslide known on land on the surface of the earth and is comparable in scale to some of the largest known submarine landslides.Many models have been proposed to explain what caused this huge slab of rocks to start sliding and what allowed it to slide so far on such a low slope, fragmenting, thinning and extending as it went. Most geologists who have worked in the area agree that Absaroka volcanism played a role in the sliding and many suggest that a major volcanic or steam explosion initiated movement. Another model involves injection of numerous igneous dikes with the resulting heating of water within pores in rocks causing an increase in pressure which initiated sliding. Some geologists have suggested that hot pressurized water (hydrothermal fluids), derived from a volcano which sat north of Cooke City, Montana, effectively lubricated the sliding surface. Another possibility is that once the slide was moving, friction heated the limestone along the sliding surface, creating pseudotachylite, which then further broke down to calcium oxide and carbon dioxide gas (or supercritical fluid). The gas supported the slide in the way that air pressure supports a hovercraft, allowing the slide to move easily down the very low slope. When the rockslide stopped, the carbon dioxide cooled and recombined with calcium oxide to form the cement-like carbonate rock now found in the fault zone. The consensus favors catastrophic sliding and calculations suggest that the front of the sliding mass may have advanced at a speed of over 100 miles/hour (160 km/h), meaning that the mountain traveled to its present location in approximately 30 minutes.In the 48 million years since the slide occurred, erosion has removed most of the portion of the slide sheet which moved out into the Bighorn Basin, leaving just one big block of carbonate rocks—Heart Mountain. Farther south, a large block of carbonate rock forms Sheep Mountain, which lies just south of the road that goes from Cody into Yellowstone Park. Some of the best views of the sliding surface, called the Heart Mountain fault, can be found along the Chief Joseph Highway (Wyoming Highway 296). The fault is particularly well exposed in Cathedral Cliffs, where it appears as a remarkably straight and nearly horizontal line just above a 2–3-meter-high cliff.

The nearby Heart Mountain War Relocation Center, where a number of Japanese Americans were interned during World War II, was named after the peak.


Keilhauite (also known as yttrotitanite) is a variety of the mineral titanite of a brownish black color, related to titanite in form. It consists chiefly of silicon dioxide, titanium dioxide, calcium oxide, and yttrium oxide. The variety was described in 1841 and named for Baltazar Mathias Keilhau (1797-1858) a Norwegian geologist.Keilhauite has a chemical formula of (CaTi,Al2,Fe23+,Y23+)SiO5. It differs from titanite only in that calcium is substituted by up to 10 percent (Y,Ce)2O3.

Lime (material)

Lime is a calcium-containing inorganic mineral composed primarily of oxides, and hydroxide, usually calcium oxide and/ or calcium hydroxide. It is also the name for calcium oxide which occurs as a product of coal seam fires and in altered limestone xenoliths in volcanic ejecta. The word lime originates with its earliest use as building mortar and has the sense of sticking or adhering.These materials are still used in large quantities as building and engineering materials (including limestone products, cement, concrete, and mortar), as chemical feedstocks, and for sugar refining, among other uses. Lime industries and the use of many of the resulting products date from prehistoric times in both the Old World and the New World. Lime is used extensively for wastewater treatment with ferrous sulfate.

The rocks and minerals from which these materials are derived, typically limestone or chalk, are composed primarily of calcium carbonate. They may be cut, crushed, or pulverized and chemically altered. Burning (calcination) of these minerals in a lime kiln converts them into the highly caustic material burnt lime, unslaked lime or quicklime (calcium oxide) and, through subsequent addition of water, into the less caustic (but still strongly alkaline) slaked lime or hydrated lime (calcium hydroxide, Ca(OH)2), the process of which is called slaking of lime.

When the term is encountered in an agricultural context, it usually refers to agricultural lime, which is crushed limestone, not a product of a lime kiln. Otherwise it most commonly means slaked lime, as the more dangerous form is usually described more specifically as quicklime or burnt lime.


Limelight (also known as Drummond light or calcium light) is a type of stage lighting once used in theatres and music halls. An intense illumination is created when an oxyhydrogen flame is directed at a cylinder of quicklime (calcium oxide), which can be heated to 2,572 °C (4,662 °F) before melting. The light is produced by a combination of incandescence and candoluminescence. Although it has long since been replaced by electric lighting, the term has nonetheless survived, as someone in the public eye is still said to be "in the limelight". The actual lights are called "limes", a term which has been transferred to electrical equivalents.

List of desiccants

A desiccant is a substance that absorbs water. It is most commonly used to remove humidity that would normally degrade or even destroy products sensitive to moisture.

List of desiccants:

Activated alumina



Bentonite clay

Calcium chloride

Calcium oxide

Calcium sulfate (Drierite)

Cobalt(II) chloride

Copper(II) sulfate

Lithium chloride

Lithium bromide

Magnesium sulfate

Magnesium perchlorate

Molecular sieve

Phosphorus pentoxide

Potassium carbonate

Potassium hydroxide

Silica gel


Sodium chlorate

Sodium chloride

Sodium hydroxide

Sodium sulfate


Sulfuric acid

Madeira Island

Madeira is a Portuguese island, and is the largest and most populous of the Madeira Archipelago. It has an area of 740.7 km2, including Ilhéu de Agostinho, Ilhéu de São Lourenço, Ilhéu Mole (northwest). As of 2011, Madeira had a total population of 262,456.

The island is the top of a massive submerged shield volcano that rises about 6 km (3.7 mi) from the floor of the Atlantic Ocean. The volcano formed atop an east-west rift in the oceanic crust along the African Plate, beginning during the Miocene epoch over 5 million years ago, continuing into the Pleistocene until about 700,000 years ago. This was followed by extensive erosion, producing two large amphitheatres open to south in the central part of the island. Volcanic activity later resumed, producing scoria cones and lava flows atop the older eroded shield. The most recent volcanic eruptions were on the west-central part of the island only 6,500 years ago, creating more cinder cones and lava flows.Madeira is the largest island of the group with an area of 741 km2 (286 sq mi), a length of 57 km (35 mi) (from Ponte de São Lourenço to Ponte do Pargo), while approximately 22 km (14 mi) at its widest point (from Ponte da Cruz to Ponte São Jorge), with a coastline of 150 km (90 mi). It has a mountain ridge that extends along the centre of the island, reaching 1,862 metres (6,109 feet) at its highest point (Pico Ruivo), while much lower (below 200 metres) along its eastern extent. The primitive volcanic foci responsible for the central mountainous area, consisted of the peaks: Ruivo (1,862 m), Torres (1,851 m), Arieiro (1,818 m), Cidrão (1,802 m), Cedro (1,759 m), Casado (1,725 m), Grande (1,657 m), Ferreiro (1,582 m). At the end of this eruptive phase, an island circled by reefs was formed, its marine vestiges are evident in a calcareous layer in the area of Lameiros, in São Vicente (which was later explored for calcium oxide production). Sea cliffs, such as Cabo Girão, valleys and ravines extend from this central spine, making the interior generally inaccessible. Daily life is concentrated in the many villages at the mouths of the ravines, through which the heavy rains of autumn and winter usually travel to the sea.

Peraluminous rock

Peraluminous rocks are igneous rocks that have a molecular proportion of aluminium oxide higher than the combination of sodium oxide, potassium oxide and calcium oxide. This contrasts with peralkaline in which the alkalis are higher, metaluminous where aluminium oxide concentration is lower than the combination, but above the alkalis, and subaluminous in which aluminia concentration is lower than the combination. Examples of peraluminous minerals include biotite, muscovite, cordierite, andalusite and garnet.

Peraluminous corresponds to the aluminum saturation index values greater than 1.Peralumneous magmas can form S-type granitoids and have been linked to collisional orogenies and to the formation of tin, tungsten and silver deposits such as those in the Bolivian tin belt.

Plate glass

Plate glass, flat glass or sheet glass is a type of glass, initially produced in plane form, commonly used for windows, glass doors, transparent walls, and windscreens. For modern architectural and automotive applications, the flat glass is sometimes bent after production of the plane sheet. Flat glass stands in contrast to container glass (used for bottles, jars, cups) and glass fibre (used for thermal insulation, in fibreglass composites, and optical communication).

Flat glass has a higher magnesium oxide and sodium oxide content than container glass, and a lower silica, calcium oxide, and aluminium oxide content. (From the lower soluble oxide content comes the better chemical durability of container glass against water, which is required especially for storage of beverages and food).

Most flat glass is soda–lime glass, produced by the float glass process. Other processes for making flat glass include:

Rolling (rolled plate glass, figure rolled glass)

Overflow downdraw method

Blown plate method

Broad sheet method

Window crown glass technique

Cylinder blown sheet method

Fourcault process

Machine drawn cylinder sheet method

Plate polishing


Rhodonite is a manganese inosilicate, (Mn, Fe, Mg, Ca)SiO3 and member of the pyroxenoid group of minerals, crystallizing in the triclinic system. It commonly occurs as cleavable to compact masses with a rose-red color (the name comes from the Greek ῥόδος rhodos, rosy), often tending to brown because of surface oxidation.

Rhodonite crystals often have a thick tabular habit, but are rare. It has a perfect, prismatic cleavage, almost at right angles. The hardness is 5.5–6.5, and the specific gravity is 3.4–3.7; luster is vitreous, being less frequently pearly on cleavage surfaces. The manganese is often partly replaced by iron, magnesium, calcium, and sometimes zinc, which may sometimes be present in considerable amounts; a greyish-brown variety containing as much as 20% of calcium oxide is called bustamite; fowlerite is a zinciferous variety containing 7% of zinc oxide.

The inosilicate (chain silicate) structure of rhodonite has a repeat unit of five silica tetrahedra. The rare polymorph pyroxmangite, formed at different conditions of pressure and temperature, has the same chemical composition but a repeat unit of seven tetrahedra.

Rhodonite has also been worked as an ornamental stone. In the iron and manganese mines at Pajsberg near Filipstad and Långban in Värmland, Sweden, small brilliant and translucent crystals (pajsbergite) and cleavage masses occur. Fowlerite occurs as large, rough crystals, somewhat resembling pink feldspar, with franklinite and zinc ores in granular limestone at Franklin Furnace in New Jersey.

Rhodonite is the official gemstone of the Commonwealth of Massachusetts.

Soda–lime glass

Soda–lime glass, also called soda–lime–silica glass, is the most prevalent type of glass, used for windowpanes and glass containers (bottles and jars) for beverages, food, and some commodity items. Glass bakeware is often made of borosilicate glass. Soda–lime glass accounts for about 90% of manufactured glass.Soda–lime glass is relatively inexpensive, chemically stable, reasonably hard, and extremely workable. Because it can be resoftened and remelted numerous times, it is ideal for glass recycling. It is used in preference to chemically-pure silica, which is silicon dioxide (SiO2), otherwise known as fused quartz. Whereas pure silica has excellent resistance to thermal shock, being able to survive immersion in water while red hot, its high melting temperature (1723 °C) and viscosity make it difficult to work with. Other substances are therefore added to simplify processing. One is the "soda", or sodium carbonate (Na2CO3), which lowers the glass-transition temperature. However, the soda makes the glass water-soluble, which is usually undesirable. To provide for better chemical durability, the "lime" is also added. This is calcium oxide (CaO), generally obtained from limestone. In addition, magnesium oxide (MgO) and alumina, which is aluminium oxide (Al2O3), contribute to the durability. The resulting glass contains about 70 to 74% silica by weight.

The manufacturing process for soda–lime glass consists in melting the raw materials, which are the silica, soda, lime (in the form of (Ca(OH)2), dolomite (CaMg(CO3)2, which provides the magnesium oxide), and aluminium oxide; along with small quantities of fining agents (e.g., sodium sulfate (Na2SO4), sodium chloride (NaCl), etc.) in a glass furnace at temperatures locally up to 1675 °C. The temperature is only limited by the quality of the furnace structure material and by the glass composition. Relatively inexpensive minerals such as trona, sand, and feldspar are usually used instead of pure chemicals. Green and brown bottles are obtained from raw materials containing iron oxide. The mix of raw materials is termed batch.

Soda–lime glass is divided technically into glass used for windows, called flat glass, and glass for containers, called container glass. The two types differ in the application, production method (float process for windows, blowing and pressing for containers), and chemical composition. Flat glass has a higher magnesium oxide and sodium oxide content than container glass, and a lower silica, calcium oxide, and aluminium oxide content. From the lower content of highly water-soluble ions (sodium and magnesium) in container glass comes its slightly higher chemical durability against water, which is required especially for storage of beverages and food.

Calcium compounds
Mixed oxidation states
+1 oxidation state
+2 oxidation state
+3 oxidation state
+4 oxidation state
+5 oxidation state
+6 oxidation state
+7 oxidation state
+8 oxidation state


This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.