Calcium hydroxide

Calcium hydroxide
Names
IUPAC name
Calcium hydroxide
Other names
  • Slaked lime
  • Milk of lime
  • Calcium(II) hydroxide
  • Pickling lime
  • Hydrated lime
  • Portlandite
  • Calcium hydrate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.013.762
EC Number 215-137-3
E number E526 (acidity regulators, ...)
846915
KEGG
RTECS number EW2800000
UNII
Properties
Ca(OH)2
Molar mass 74.093 g/mol
Appearance White powder
Odor Odorless
Density 2.211 g/cm3, solid
Melting point 580 °C (1,076 °F; 853 K) (loses water, decomposes)
  • 1.89 g/L (0 °C)
  • 1.73 g/L (20 °C)
  • 0.66 g/L (100 °C)
5.5×10−6
Solubility
Basicity (pKb) 1.37 (first OH−), 2.43 (second OH−)[1][2]
−22.0·10−6 cm3/mol
1.574
Structure
Hexagonal, hP3[3]
P3m1 No. 164
a = 0.35853 nm, c = 0.4895 nm
Thermochemistry
83 J·mol−1·K−1[4]
−987 kJ·mol−1[4]
Hazards
Safety data sheet See: data page
[5]
R-phrases (outdated) R22, R34
S-phrases (outdated) (S2), S24
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 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
0
3
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
7340 mg/kg (oral, rat)
7300 mg/kg (mouse)
US health exposure limits (NIOSH):
PEL (Permissible)
 TWA 15 mg/m3 (total) 5 mg/m3 (resp.)[6]
REL (Recommended)
 TWA 5 mg/m3[6]
IDLH (Immediate danger)
 N.D.[6]
Related compounds
Other cations
 Magnesium hydroxide
Strontium hydroxide
Barium hydroxide
Related bases
 Calcium oxide
Supplementary data page
Refractive index (n),
Dielectric constant (εr), etc.
Thermodynamic
data
 Phase behaviour
solid–liquid–gas
UV, IR, NMR, MS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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 obtained when calcium oxide (called lime or quicklime) 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. Limewater is the common name for a saturated solution of calcium hydroxide.

Properties

Calcium hydroxide is relatively insoluble in water, with a solubility product Ksp of 5.5 × 10−6. It is large enough that its solutions are basic according to the following reaction:

Ca(OH)2 → Ca2+ + 2 OH−

At ambient temperature, calcium hydroxide (portlandite) dissolves in pure water to produce an alkaline solution with a pH of about 12.4. Calcium hydroxide solutions can cause chemical burns. At high pH value (see common ion effect), its solubility drastically decreases. This behavior is relevant to cement pastes. Its aqueous solutions is called limewater and is a medium strength base that reacts with acids and can attack some metals such as aluminium (amphoteric hydroxide dissolving at high pH) while protecting other metals from corrosion such as iron and steel by passivation of their surface. Limewater turns milky in the presence of carbon dioxide due to formation of calcium carbonate, a process called carbonatation:

Ca(OH)2 + CO2 → CaCO3 + H2O

When heated to 512 °C, the partial pressure of water in equilibrium with calcium hydroxide reaches 101 kPa (normal atmospheric pressure), which decomposes calcium hydroxide into calcium oxide and water.[7]

Ca(OH)2 → CaO + H2O

Structure, preparation, occurrence

SEM image of fractured hardened cement paste, showing plates of calcium hydroxide and needles of ettringite (micron scale)

Calcium hydroxide adopts a polymeric structure, as do all metal hydroxides. The structure is identical to that of Mg(OH)2 (brucite structure); i.e., the cadmium iodide motif. Strong hydrogen bonds exist between the layers.[8]

Calcium hydroxide is produced commercially by treating lime with water:

CaO + H2O → Ca(OH)2

In the laboratory it can be prepared by mixing aqueous solutions of calcium chloride and sodium hydroxide. The mineral form, portlandite, is relatively rare but can be found in some volcanic, plutonic, and metamorphic rocks. It has also been known to arise in burning coal dumps.

The positively charged ionized species CaOH+ has been detected in the atmosphere of S-type stars.[9]

Retrograde solubility

The solubility of calcium hydroxide (portlandite) at 70 °C is about half of its value at 25 °C. The reason of this not so common behaviour is that the dissolution of calcium hydroxide in water is an exothermic process and also obeys to the Le Chatelier's principle. A lowering of temperature will thus favour the elimination of the dissolution heat and increases the equilibrium constant of dissolution of Ca(OH)2, and so increase its solubility at low temperature. This contra-intuitive temperature dependence of the solubility is referred to as "retrograde" or "inverse" solubility. The differently hydrated phases of calcium sulfate (gypsum, bassanite and anhydrite) also exhibit a retrograde solubility for the same reason because their dissolution reactions are exothermic.

Uses

One significant application of calcium hydroxide is as a flocculant, in water and sewage treatment. It forms a fluffy charged solid that aids in the removal of smaller particles from water, resulting in a clearer product. This application is enabled by the low cost and low toxicity of calcium hydroxide. It is also used in fresh water treatment for raising the pH of the water so that pipes will not corrode where the base water is acidic, because it is self-regulating and does not raise the pH too much.

It is also used in the preparation of ammonia gas, using the following reaction:

Ca(OH)2 + 2NH4Cl → 2NH3 + CaCl2 + 2H2O

Another large application is in the paper industry, where it is an intermediate in the reaction in the production of sodium hydroxide. This conversion is part of the causticizing step in the Kraft process for making pulp.[8] In the causticizing operation burned lime is added to green liquor which is a solution primarily of sodium carbonate and sodium sulfate produced by dissolving smelt, which is the molten form of these chemicals from the recovery furnace.

Food industry

Because of its low toxicity and the mildness of its basic properties, slaked lime is widely used in the food industry to:

Native American uses

Dry treated corn (left), and untreated corn (right) after boiling in water with calcium hydroxide (15 ml, or 1 Tbsp, lime for 500 g of corn) for 15 minutes.

In Spanish, calcium hydroxide is called cal. Corn cooked with cal (nixtamalization) becomes hominy (nixtamal), which significantly increases the bioavailability of niacin, and it is also considered tastier and easier to digest.

In chewing coca leaves, calcium hydroxide is usually chewed alongside to keep the alkaloid stimulants chemically available for absorption by the body. Similarly, Native Americans traditionally chewed tobacco leaves with calcium hydroxide derived from burnt mollusc shells to enhance the effects. It has also been used by some indigenous American tribes as an ingredient in yopo, a psychedelic snuff prepared from the beans of some Anadenanthera species.[10]

Asian uses

Calcium hydroxide is typically added to a bundle of areca nut and betel leaf to keep the alkaloid stimulants chemically available to enter the bloodstream via sublingual absorption.

Afghan uses

It is used in making naswar (also known as nass or niswar), a type of dipping tobacco made from fresh tobacco leaves, calcium hydroxide (chuna), and wood ash. It is consumed most in the Pathan diaspora, Afghanistan, Pakistan, India, Bangladesh and also in Sweden and Norway. Villagers also use calcium hydroxide to paint their mud houses in Afghanistan, Pakistan and India.

Health risks

Unprotected exposure to Ca(OH)2 can cause severe skin irritation, chemical burns, blindness or lung damage.[5]

See also

References

  1. ↑ "Sortierte Liste: pKb-Werte, nach Ordnungszahl sortiert. – Das Periodensystem online".
  2. ↑ ChemBuddy dissociation constants pKa and pKb
  3. ↑ Petch, H. E. (1961). "The hydrogen positions in portlandite, Ca(OH)2, as indicated by the electron distribution". Acta Crystallographica. 14 (9): 950. doi:10.1107/S0365110X61002771.
  4. 1 2 Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A21. ISBN 0-618-94690-X.
  5. 1 2 "MSDS Calcium hydroxide" (PDF). Archived from the original (PDF) on 25 March 2012. Retrieved 21 June 2011.
  6. 1 2 3 "NIOSH Pocket Guide to Chemical Hazards #0092". National Institute for Occupational Safety and Health (NIOSH).
  7. ↑ Halstead, P.E.; Moore, A.E. (1957). "The Thermal Dissociation Of Calcium Hydroxide". Journal of the Chemical Society. 769: 3873. doi:10.1039/JR9570003873.
  8. 1 2 Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  9. ↑ Jørgensen, Uffe G. (1997), "Cool Star Models", in van Dishoeck, Ewine F., Molecules in Astrophysics: Probes and Processes, International Astronomical Union Symposia. Molecules in Astrophysics: Probes and Processes, 178, Springer Science & Business Media, p. 446, ISBN 079234538X.
  10. ↑ de Smet, Peter A. G. M. (1985). "A multidisciplinary overview of intoxicating snuff rituals in the Western Hemisphere". Journal of Ethnopharmacology. 3 (1): 3–49. doi:10.1016/0378-8741(85)90060-1.

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