Calcium hydroxide

Source: Wikipedia, the free encyclopedia.

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
  • Calcium dihydroxide
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
100.013.762 Edit this at Wikidata
EC Number
  • 215-13
E number E526 (acidity regulators, ...)
846915
KEGG
RTECS number
  • EW2800000
UNII
  • InChI=1S/Ca.2H2O/h;2*1H2/q+2;;/p-2 checkY
    Key: AXCZMVOFGPJBDE-UHFFFAOYSA-L checkY
  • InChI=1/Ca.2H2O/h;2*1H2/q+2;;/p-2
    Key: AXCZMVOFGPJBDE-NUQVWONBAD
  • [Ca+2].[OH-].[OH-]
  • [OH-].[OH-].[Ca+2]
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)
  • (retrograde solubility, i.e., unusually decreasing with T)
5.02×10−6 [1]
Solubility
Acidity (pKa) 12.63 (first OH), 11.57 (second OH)[2][3] [clarification needed]
−22.0·10−6 cm3/mol
1.574
Structure
Hexagonal, hP3[4]
P3m1 No. 164
a = 0.35853 nm, c = 0.4895 nm
Thermochemistry
83 J·mol−1·K−1[5]
Std enthalpy of
formation
fH298)
−987 kJ·mol−1[5]
Hazards
GHS labelling:
GHS05: Corrosive GHS07: Exclamation mark
Danger
H314, H335, H402
P261, P280, P305+P351+P338
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 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
7340 mg/kg (oral, rat)
7300 mg/kg (mouse)
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 15 mg/m3 (total) 5 mg/m3 (resp.)[7]
REL (Recommended)
TWA 5 mg/m3[7]
IDLH
(Immediate danger)
N.D.[7]
Safety data sheet (SDS) [6]
Related compounds
Other cations
Magnesium hydroxide
Strontium hydroxide
Barium hydroxide
Related bases
Calcium oxide
Supplementary data page
Calcium hydroxide (data page)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Calcium hydroxide (traditionally called slaked lime) is an

saturated solution
of calcium hydroxide.

Properties

Calcium hydroxide is poorly soluble in water, with a retrograde solubility increasing from 0.66 g/L at 100 °C to 1.89 g/L at 0 °C. With a

solubility product Ksp of 5.02×10−6 at 25 °C,[1][clarification needed
] its dissociation in water is large enough that its solutions are basic according to the following dissolution reaction:

Ca(OH)2 → Ca2+ + 2 OH

At ambient temperature, calcium hydroxide (

bases, which react with acids and can attack some metals such as aluminium[citation needed] (amphoteric hydroxide dissolving at high pH), while protecting other metals, such as iron and steel, from corrosion by passivation of their surface. Limewater turns milky in the presence of carbon dioxide due to the formation of insoluble 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:[8]

Ca(OH)2 → CaO + H2O

Calcium hydroxide reacts with hydrogen chloride to first give calcium hydroxychloride and then calcium chloride.

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.[9]

Calcium hydroxide is produced commercially by treating (slaking) lime with water:

CaO + H2O → Ca(OH)2

In the laboratory it can be prepared by mixing

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.[10]

Retrograde solubility

According to Hopkins and Wulff (1965),

Julius Thomsen (1883)[13] (see Thomsen–Berthelot principle), when the presence of ions in aqueous solutions was still questioned. Since, it has been studied in detail by many authors, a.o., Miller and Witt (1929)[14] or Johnston and Grove (1931)[15] and refined many times (e.g., Greenberg and Copeland (1960);[16] Hopkins and Wulff (1965);[11] Seewald and Seyfried (1991);[17] Duchesne and Reardon (1995)[18]
).

The reason for this rather uncommon behavior is that the dissolution of calcium hydroxide in water involves an entropy decrease, due to the ordering of water molecules around the doubly charged calcium ion, This counter-intuitive temperature dependence of the solubility is referred to as "retrograde" or "inverse" solubility. The variably hydrated phases of calcium sulfate (gypsum, bassanite and anhydrite) also exhibit a slight retrograde solubility due to the presence of the calcium ion. However, other calcium salts like calcium chloride show an increase in solubility since the enthalpy change is larger and dominates the free energy change during dissolution. In all cases, the dissolution is exothermic.[citation needed]

Uses

Calcium hydroxide is commonly used to prepare lime mortar.

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.[citation needed]

It is also used in the preparation of ammonia gas (NH3), using the following reaction:

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

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.[9] 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.[citation needed]

In

Neonectria galligena. The trees are sprayed when they are dormant in winter to prevent toxic burns from the highly reactive calcium hydroxide. This use is authorised in the European Union and the United Kingdom under Basic Substance regulations.[19]

Calcium hydroxide is used in dentistry, primarily in the specialty of endodontics.

Food industry

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

Native American uses

niacin (vitamin B3), and is also considered tastier and easier to digest. Nixtamal is often ground into a flour, known as masa, which is used to make tortillas and tamales.[citation needed
]

In chewing

yopo, a psychedelic snuff prepared from the beans of some Anadenanthera species.[23]

Asian uses

Calcium hydroxide is typically added to a bundle of

sublingual
absorption.

It is used in making

Pathan diaspora, Afghanistan, Pakistan, India and Bangladesh. Villagers also use calcium hydroxide to paint their mud houses
in Afghanistan, Pakistan and India.

Health risks

Unprotected exposure to Ca(OH)2, as with any strong base, can cause severe skin irritation, chemical burns, blindness, lung damage or rashes.[6]

See also

References

  1. ^ .
  2. ^ "Sortierte Liste: pKb-Werte, nach Ordnungszahl sortiert. – Das Periodensystem online".
  3. ^ ChemBuddy dissociation constants pKa and pKb
  4. .
  5. ^ .
  6. ^ a b "MSDS Calcium hydroxide" (PDF). Archived from the original (PDF) on 25 March 2012. Retrieved 21 June 2011.
  7. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0092". National Institute for Occupational Safety and Health (NIOSH).
  8. .
  9. ^ .
  10. .
  11. ^ .
  12. ^ Berthelot, M. (1875). Dissolution des acides et des alcalis. [Dissolution of acids and alkalis]. In: Annales de Chimie et de Physique. Vol. 4, pp. 445–536.
  13. ^ Thomsen J. (1883). Thermochemische untersuchungen [Thermochemical studies]. Vol. III, Johann Ambrosius Barth Verlag, Leipzig.
  14. ISSN 0092-7325
    .
  15. .
  16. .
  17. .
  18. .
  19. ^ European Union (13 May 2015). "COMMISSION IMPLEMENTING REGULATION (EU) 2015/762 of 12 May 2015 approving the basic substance calcium hydroxide in accordance with Regulation (EC) No 1107/2009 of the European Parliament and of the Council concerning the placing of plant protection products on the market, and amending the Annex to Commission Implementing Regulation (EU) No 540/2011". Retrieved 12 May 2022.
  20. ^ Pesticide Research Institute for the USDA National Organic Program (23 March 2015). "Hydrated Lime: Technical Evaluation Report" (PDF). Agriculture Marketing Services. Retrieved 17 July 2019.
  21. ^ Borsook, Alec (6 August 2015). "Cooking with Alkali". Nordic Food Lab.
  22. ^ "Preparation of Mushroom Growing Substrates". North American Mycological Association. Retrieved 8 July 2021.
  23. PMID 3887041
    .

External links