Calcium oxide

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Calcium oxide
Calcium oxide
Calcium oxide powder.JPG
Identifiers
CAS number 1305-78-8 YesY
PubChem 14778
ChemSpider 14095
UNII C7X2M0VVNH
UN number 1910
RTECS number EW3100000
ATCvet code QP53AX18
Jmol-3D images Image 1
Properties
Molecular formula 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 2,850 °C (5,160 °F; 3,120 K) , 3123 K (100 hPa)[3]
Solubility in water 1.19 g/L (25 °C), 0.57 g/L (100 °C), exothermic reaction[2]
Solubility in acids soluble (also in glycerol, sugar solution)
Solubility in methanol insoluble (also in diethyl ether, n-octanol)
Acidity (pKa) 12.8
Thermochemistry
Std molar
entropy
So298
40 J·mol−1·K−1[4]
Std enthalpy of
formation
ΔfHo298
−635 kJ·mol−1[4]
Hazards
MSDS [1]
EU Index Not listed
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Related compounds
Other anions Calcium sulfide
Calcium hydroxide
Other cations Beryllium oxide
Magnesium oxide
Strontium oxide
Barium oxide
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references

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[edit]

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 metric tons. 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]

Usage[edit]

CaO (s) + H2O (l) is in equilibrium with 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.[9]
  • Cement: Calcium oxide is a key ingredient for the process of making cement.
  • As an alkali in biodiesel production[10][11]
  • 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.[12][13][14] 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.

Use as a weapon[edit]

Historian and philosopher 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:

D’Albiney employed a stratagem against them, which is said to have contributed to the victory: Having gained the wind of the French, he came down upon them with violence; and throwing in their faces a great quantity of quicklime, which he purposely carried on board, he so blinded them, that they were disabled from defending themselves.[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]

Health issues[edit]

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[edit]

  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. ^ Committee on Water Treatment Chemicals, Food and Nutrition Board, Assembly of Life Sciences, National Research Council (1982). Water Chemicals Codex. p. 20. ISBN 0-309-07368-5. 
  3. ^ Calciumoxid. GESTIS database
  4. ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A21. ISBN 0-618-94690-X. 
  5. ^ "free lime" DictionaryOfConstruction.com. WebFinance, Inc. May 29, 2014 <http://www.dictionaryofconstruction.com/definition/free-lime.html>.
  6. ^ Merck Index of chemicals and Drugs , 9th edition monograph 1650
  7. ^ Miller, M. Michael (2007). "Lime". Minerals Yearbook. U.S. Geological Survey. p. 43.13. 
  8. ^ Collie, Robert L. "Solar heating system" U.S. Patent 3,955,554 issued May 11, 1976
  9. ^ Gray, Theodore (September 2007). "Limelight in the Limelight". Popular Science: 84. 
  10. ^ Kozu, Masato; et al (2008). "Calcium oxide as a solid base catalyst for transesterification of soybean oil and its application to biodiesel production". Fuel (Elsevier) 87 (12). doi:10.1016/j.fuel.2007.10.019. Retrieved 19 March 2014. 
  11. ^ Zhu, Huaping; et al (2006). "Preparation of Biodiesel Catalyzed by Solid Super Base of Calcium Oxide and Its Refining Process". Chinese Journal of Catalysis (Elsevier) 27 (5): 391–396. doi:10.1016/S1872-2067(06)60024-7. Retrieved 19 March 2014. 
  12. ^ Neolithic man: The first lumberjack?. Phys.org (August 9, 2012). Retrieved on 2013-01-22.
  13. ^ Karkanas, Panagiotis; Stratouli, Georgia (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. 
  14. ^ Connelly, Ashley Nicole (May 2012) Analysis and Interpretation of Neolithic Near Eastern Mortuary Rituals from a Community-Based Perspective. Baylor University Thesis, Texas
  15. ^ David Hume (1756). History of England 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

External links[edit]