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Calcium oxide

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Calcium oxide
Calcium oxide
Ionic crystal structure of calcium oxide
  Ca2+   O2-

Powder sample of white calcium oxide
IUPAC name
Calcium oxide
Other names
  • Lime
  • Quicklime
  • Burnt lime
  • Unslaked lime
  • Free lime (building)
  • Caustic lime
  • Pebble lime
  • Calcia
  • Oxide of calcium
3D model (JSmol)
ECHA InfoCard 100.013.763 Edit this at Wikidata
EC Number
  • 215-138-9
E number E529 (acidity regulators, ...)
RTECS number
  • EW3100000
UN number 1910
  • InChI=1S/Ca.O
  • InChI=1/Ca.O/rCaO/c1-2
  • O=[Ca]
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)
GHS labelling:
GHS05: CorrosiveGHS07: Exclamation mark
H302, H314, H315, H335
P260, P261, P264, P270, P271, P280, P301+P312, P301+P330+P331, P302+P352, P303+P361+P353, P304+P340, P305+P351+P338, P310, P312, P321, P330, P332+P313, P362, P363, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 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 acid
Flash point Non-flammable[4]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 5 mg/m3[4]
REL (Recommended)
TWA 2 mg/m3[4]
IDLH (Immediate danger)
25 mg/m3[4]
Safety data sheet (SDS) ICSC 0409
Related compounds
Other anions
Calcium sulfide
Calcium hydroxide
Calcium selenide
Calcium telluride
Other cations
Beryllium oxide
Magnesium oxide
Strontium oxide
Barium oxide
Radium oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Calcium oxide (formula: 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 compounds, in which carbonates, oxides, and hydroxides of calcium, silicon, magnesium, aluminium, and iron predominate. By contrast, quicklime specifically applies to the single 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 the chemical derivative calcium hydroxide (of which quicklime is the base anhydride) are important commodity chemicals.



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][7] a process called calcination or lime-burning, to liberate a molecule of carbon dioxide (CO2), leaving quicklime behind. This is also one of the few chemical reactions known in prehistoric times.[8]

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.[9]

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.[10]

Hydroxyapatite's free CaO content rises with increased calcination temperatures and longer times. It also pinpoints particular temperature cutoffs and durations that impact the production of CaO, offering information on how calcination parameters impact the composition of the material.


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.[10]
  • Ground quicklime is used in the production of aerated concrete such as blocks with densities of ca. 0.6–1.0 g/cm3 (9.8–16.4 g/cu in).[10]
  • 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.[10]
  • Small quantities of quicklime are used in other processes; e.g., the production of glass, calcium aluminate cement, and organic chemicals.[10]
  • Heat: Quicklime releases thermal energy by the formation of the hydrate, calcium hydroxide, by the following equation:[11]
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.[12]
  • It is known as a food additive to the FAO as an acidity regulator, a flour treatment agent and as a leavener.[13] 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.[14]
  • 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.[10]
  • 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.[15][16][17] 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.
  • Carbon capture and storage: Calcium oxide can be used to capture carbon dioxide from flue gases in a process called calcium looping.
  • 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.[18][19][20]
  • Disposal of corpses: Historically, it was mistakenly believed that quicklime was efficacious in accelerating the decomposition of corpses. The application of quicklime can, in fact, promote preservation. Quicklime can aid in eradicating the stench of decomposition, which may have led people to the erroneous conclusion.[21]
  • It has been determined that the durability of ancient Roman concrete is attributed in part to the use of quicklime as an ingredient. Combined with hot mixing, the quicklime creates macro-sized lime clasts with a characteristically brittle nano-particle architecture. As cracks form in the concrete, they preferentially pass through the structurally weaker lime clasts, fracturing them. When water enters these cracks it creates a calcium-saturated solution which can recrystallize as calcium carbonate, quickly filling the crack. [22]
  • The thermochemical heat storage mechanism is greatly impacted by the sintering of CaO and CaCO3. It demonstrates that the storage materials become less reactive and denser at increasing temperatures. It also pinpoints particular sintering processes and variables influencing the efficiency of these materials in heat storage.



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.[23] 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.[24]

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.[25]

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.[26] Quicklime may have been used in medieval naval warfare – up to the use of "lime-mortars" to throw it at the enemy ships.[27]



Limestone is a substitute for lime in many applications, which include 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.[28]



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, and 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.[29][better source needed]



Calcium oxide is also a separate mineral species (with the unit formula CaO), named 'Lime'.[30][31] It has an isometric crystal system, and can form a solid solution series with monteponite. The crystal is brittle, pyrometamorphic, and is unstable in moist air, quickly turning into portlandite (Ca(OH)2).[32]


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  9. ^ Miller, M. Michael (2007). "Lime". Minerals Yearbook (PDF). U.S. Geological Survey. p. 43.13.
  10. ^ 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 978-3527306732
  11. ^ Collie, Robert L. "Solar heating system" U.S. patent 3,955,554 issued May 11, 1976
  12. ^ Gretton, Lel. "Lime power for cooking - medieval pots to 21st century cans". Old & Interesting. Retrieved 13 February 2018.
  13. ^ "Compound Summary for CID 14778 - Calcium Oxide". PubChem.
  14. ^ Gray, Theodore (September 2007). "Limelight in the Limelight". Popular Science: 84. Archived from the original on 2008-10-13. Retrieved 2009-03-31.
  15. ^ Tel Aviv University (August 9, 2012). "Neolithic man: The first lumberjack?". phys.org. Retrieved 2023-02-02.
  16. ^ 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: 27–41. doi:10.1017/S006824540000006X. S2CID 129562287.
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  20. ^ US Patent 255042, 14 March 1882
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  22. ^ "Riddle solved: Why was Roman concrete so durable?", MIT News, January 6, 2023
  23. ^ Plutarch, "Sertorius 17.1–7", Parallel Lives
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  25. ^ Croddy, Eric (2002). Chemical and biological warfare: a comprehensive survey for the concerned citizen. Springer. p. 128. ISBN 0-387-95076-1.
  26. ^ David Hume (1756). History of England. Vol. I.
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  29. ^ Mallinckrodt Baker Inc. - Strategic Services Division (December 8, 1996). "Hazards". ww25.hazard.com. Archived from the original on May 1, 2012. Retrieved 2023-02-02.{{cite web}}: CS1 maint: unfit URL (link)
  30. ^ "List of Minerals". Ima-mineralogy.org. 21 March 2011.
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