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Calcium oxide

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Calcium oxide
Calcium oxide
Names
IUPAC name
Calcium oxide
Other names
Quicklime, burnt lime, unslaked lime, pebble lime
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.013.763 Edit this at Wikidata
E number E529 (acidity regulators, ...)
485425
RTECS number
  • EW3100000
UNII
UN number 1910
  • InChI=1/Ca.O/rCaO/c1-2
    Key: ODINCKMPIJJUCX-BFMVISLHAU
  • [Ca]=O
Properties
CaO
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 3,850 °C (6,960 °F; 4,120 K) (100 hPa)[2]
Reacts to form calcium hydroxide
Solubility in Methanol Insoluble (also in diethyl ether, n-octanol)
Acidity (pKa) 12.8
Structure
cubic, cF8
Thermochemistry
40 J·mol−1·K−1[3]
−635 kJ·mol−1[3]
Pharmacology
QP53AX18 (WHO)
Hazards
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 hazards (white): no code
3
0
2
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) Hazard.com
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 (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 which 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.

Preparation

A demonstration of slaking of quicklime as a strongly exoenergetic reaction. Drops of water are added to pieces of quicklime. After a while, a pronounced exoenergetic reaction occurs ('slaking of lime'). The temperature can reach up to some 300 ºС.
Planned and performed by Marina Stojanovska, Miha Bukleski and Vladimir Petruševski, Department of Chemistry, FNSM, Ss. Cyril and Methodius University, Skopje, Macedonia.

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]

Usage

  • The major use of quicklime is in the Basic oxygen steelmaking (BOS) process. Its usage varies from about 30–50 kg/t 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.[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.
  • 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 prior to the invention of electric lighting.[10]
  • 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.[11][12][13] 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.

As a weapon

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. A similar dust was used in China to quell an armed peasant revolt in 178 CE, when "lime chariots" equipped with bellows blew limestone powder into the crowds.[14]

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.[15]

Quicklime is also thought to have been a component of Greek fire. Upon contact with water, quicklime would increase its temperature above 150 °C and ignite the fuel.[16]

Safety

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.[17]

References

  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 1-4398-5511-0.
  2. ^ Calciumoxid. 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". DictionaryOfConstruction.com.
  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. ^ Gray, Theodore (September 2007). "Limelight in the Limelight". Popular Science: 84.
  11. ^ Neolithic man: The first lumberjack?. Phys.org (August 9, 2012). Retrieved on 2013-01-22.
  12. ^ 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.
  13. ^ Connelly, Ashley Nicole (May 2012) Analysis and Interpretation of Neolithic Near Eastern Mortuary Rituals from a Community-Based Perspective. Baylor University Thesis, Texas
  14. ^ Adrienne Mayor (2005), "Ancient Warfare and Toxicology", in Philip Wexler (ed.), Encyclopedia of Toxicology, vol. 4 (2nd ed.), Elsevier, pp. 117–121, ISBN 0-12-745354-7
  15. ^ David Hume (1756). History of England. Vol. I.
  16. ^ Croddy, Eric (2002). Chemical and biological warfare: a comprehensive survey for the concerned citizen. Springer. p. 128. ISBN 0-387-95076-1.
  17. ^ CaO MSDS. hazard.com
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