Calcium chloride
Structure of calcium chloride, (chlorine is green, calcium is gray)
Sample of calcium chloride
IUPAC name
Calcium chloride
Other names
  • Neutral calcium chloride
  • calcium(II) chloride
  • calcium dichloride (1:2)
  • E509
3D model (JSmol)
ECHA InfoCard 100.030.115 Edit this at Wikidata
EC Number
  • 233-140-8
E number E509 (acidity regulators, ...)
RTECS number
  • EV9800000
  • InChI=1S/Ca.2ClH/h;2*1H/q+2;;/p-2 checkY
  • InChI=1/Ca.2ClH/h;2*1H/q+2;;/p-2
  • Cl[Ca]Cl
  • [Ca+2].[Cl-].[Cl-]
  • monohydrate: Cl[Ca]Cl.O
  • dihydrate: Cl[Ca]Cl.O.O
  • hexahydrate: Cl[Ca]Cl.O.O.O.O.O.O
Molar mass 110.98 g·mol−1
Appearance White hygroscopic powder
Odor Odorless
  • 2.15 g/cm3 (anhydrous)
  • 2.24 g/cm3 (monohydrate)
  • 1.85 g/cm3 (dihydrate)
  • 1.83 g/cm3 (tetrahydrate)
  • 1.71 g/cm3 (hexahydrate)[1]
Melting point 772–775 °C (1,422–1,427 °F; 1,045–1,048 K)
260 °C (500 °F; 533 K)
monohydrate, decomposes
175 °C (347 °F; 448 K)
dihydrate, decomposes
45.5 °C (113.9 °F; 318.6 K)
tetrahydrate, decomposes[5]
30 °C (86 °F; 303 K)
hexahydrate, decomposes[1]
Boiling point 1,935 °C (3,515 °F; 2,208 K) anhydrous[1]
74.5 g/100 mL (20 °C)[2]
49.4 g/100 mL (−25 °C)
59.5 g/100 mL (0 °C)
65 g/100 mL (10 °C)
81.1 g/100 mL (25 °C)[1]
102.2 g/100 mL (30.2 °C)
90.8 g/100 mL (20 °C)
114.4 g/100 mL (40 °C)
134.5 g/100 mL (60 °C)
152.4 g/100 mL (100 °C)[3]
Solubility in ethanol
  • 18.3 g/100 g (0 °C)
  • 25.8 g/100 g (20 °C)
  • 35.3 g/100 g (40 °C)
  • 56.2 g/100 g (70 °C)[4]
Solubility in methanol
  • 21.8 g/100 g (0 °C)
  • 29.2 g/100 g (20 °C)
  • 38.5 g/100 g (40 °C)[4]
Solubility in acetone 0.1 g/kg (20 °C)[4]
Solubility in pyridine 16.6 g/kg[4]
Acidity (pKa)
  • 8–9 (anhydrous)
  • 6.5–8.0 (hexahydrate)
−5.47·10−5 cm3/mol[1]
  • 3.34 cP (787 °C)
  • 1.44 cP (967 °C)[4]
  • Pnnm, No. 58 (anhydrous)
  • P42/mnm, No. 136 (anhydrous, >217 °C)[6]
  • 2/m 2/m 2/m (anhydrous)
  • 4/m 2/m 2/m (anhydrous, >217 °C)[6]
a = 6.259 Å, b = 6.444 Å, c = 4.17 Å (anhydrous, 17 °C)[6]
α = 90°, β = 90°, γ = 90°
Octahedral at Ca2+ centres (anhydrous)
  • 72.89 J/(mol·K) (anhydrous)[1]
  • 106.23 J/(mol·K) (monohydrate)
  • 172.92 J/(mol·K) (dihydrate)
  • 251.17 J/(mol·K) (tetrahydrate)
  • 300.7 J/(mol·K) (hexahydrate)[5]
108.4 J/(mol·K)[1][5]
  • −795.42 kJ/mol (anhydrous)[1]
  • −1110.98 kJ/mol (monohydrate)
  • −1403.98 kJ/mol (dihydrate)
  • −2009.99 kJ/mol (tetrahydrate)
  • −2608.01 kJ/mol (hexahydrate)[5]
−748.81 kJ/mol[1][5]
A12AA07 (WHO) B05XA07 (WHO), G04BA03 (WHO)
Occupational safety and health (OHS/OSH):
Main hazards
GHS labelling:
GHS07: Exclamation mark[7]
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
1,000-1,400 mg/kg (rats, oral)[8]
Related compounds
Other anions
Other cations
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Calcium chloride is an inorganic compound, a salt with the chemical formula CaCl2. It is a white crystalline solid at room temperature, and it is highly soluble in water. It can be created by neutralising hydrochloric acid with calcium hydroxide.

Calcium chloride is commonly encountered as a hydrated solid with generic formula CaCl2·nH2O, where n = 0, 1, 2, 4, and 6. These compounds are mainly used for de-icing and dust control. Because the anhydrous salt is hygroscopic and deliquescent, it is used as a desiccant.[10]


Calcium chloride was apparently discovered in the 15th century but wasn't studied properly until the 18th century.[11] It was historically called "fixed sal ammoniac" (Latin: sal ammoniacum fixum[12]) because it was synthesized during the distillation of ammonium chloride with lime and was nonvolatile (while the former appeared to sublime); in more modern times (18th-19th cc.) it was called "muriate of lime" (Latin: murias calcis, calcaria muriatica[12]).[13]


De-icing and freezing-point depression

Bulk CaCl2 for de-icing in Japan

By depressing the freezing point of water, calcium chloride is used to prevent ice formation and is used to de-ice. This application consumes the greatest amount of calcium chloride. Calcium chloride is relatively harmless to plants and soil. As a de-icing agent, it is much more effective at lower temperatures than sodium chloride. When distributed for this use, it usually takes the form of small, white spheres a few millimeters in diameter, called prills. Solutions of calcium chloride can prevent freezing at temperatures as low as −52 °C (−62 °F), making it ideal for filling agricultural implement tires as a liquid ballast, aiding traction in cold climates.[14]

It is also used in domestic and industrial chemical air dehumidifiers.[15]

Road surfacing

Calcium chloride was sprayed on this road to prevent weathering, giving it a wet appearance even in dry weather.

The second largest application of calcium chloride exploits its hygroscopic nature and the tackiness of its hydrates; calcium chloride is highly hygroscopic and its hydration is an exothermic process. A concentrated solution keeps a liquid layer on the surface of dirt roads, which suppresses the formation of dust. It keeps the finer dust particles on the road, providing a cushioning layer. If these are allowed to blow away, the large aggregate begins to shift around and the road breaks down. Using calcium chloride reduces the need for grading by as much as 50% and the need for fill-in materials as much as 80%.[16]


In the food industry, calcium chloride is frequently employed as a firming agent in canned vegetables, particularly for canned tomatoes and cucumber pickles.[17] It is also used in firming soybean curds into tofu and in producing a caviar substitute from vegetable or fruit juices.[17] It is also used to enhance the texture of various other products, such as whole apples, whole hot peppers, whole and sliced strawberries, diced tomatoes, and whole peaches.[18][19]

The firming effect of calcium chloride can be attributed to several mechanisms:[18]

  1. Complexation, since calcium ions form complexes with pectin, a polysaccharide found in the cell wall and middle lamella of plant tissues.[18]
  2. Membrane stabilization, since calcium ions contribute to the stabilization of the cell membrane.[18]
  3. Turgor pressure regulation, since calcium ions influence cell turgor pressure, which is the pressure exerted by the cell contents against the cell wall.[18]

Calcium chloride's freezing-point depression properties are used to slow the freezing of the caramel in caramel-filled chocolate bars.[17] Also, it is frequently added to sliced apples to maintain texture.[20]

In brewing beer, calcium chloride is sometimes used to correct mineral deficiencies in the brewing water. It affects flavor and chemical reactions during the brewing process, and can also affect yeast function during fermentation.[17]

In cheesemaking, calcium chloride is sometimes added to processed (pasteurized/homogenized) milk to restore the natural balance between calcium and protein in casein. It is added before the coagulant.[17]

Calcium chloride is also commonly used as an "electrolyte" in sports drinks and other beverages, including bottled water.[21][17]

The average intake of calcium chloride as food additives has been estimated to be 160–345 mg/day.[22] Calcium chloride is permitted as a food additive in the European Union for use as a sequestrant and firming agent with the E number E509.[17] It is considered as generally recognized as safe (GRAS) by the U.S. Food and Drug Administration.[23] Its use in organic crop production is generally prohibited under the US National Organic Program.[24]

Calcium chloride contains approximately 27.2% or 272 mg of elemental calcium per gram. This means that for every gram of calcium chloride, there are 272 mg of actual, absorbable calcium. Calcium chloride has a very salty taste and can cause mouth and throat irritation at high concentrations, so it is typically not the first choice for long-term oral supplementation (as a calcium supplement).[25][26] Calcium chloride, characterized by its low molecular weight and high water solubility, readily breaks down into calcium and chloride ions when exposed to water. These ions are efficiently absorbed from the intestine.[27] However, caution should be exercised when handling calcium chloride, for it has the potential to release heat energy upon dissolution in water. This release of heat can lead to trauma and burns in the mouth, throat, esophagus, and stomach. In fact, there have been reported cases of stomach necrosis resulting from burns caused by accidental ingestions of big amounts of dry calcium chloride.[28][29]

The extremely salty taste of calcium chloride is used to flavor pickles without increasing the food's sodium content.[17]

Calcium chloride is used to prevent cork spot and bitter pit on apples by spraying on the tree during the late growing season.[30]

Laboratory and related drying operations

Drying tubes are frequently packed with calcium chloride. Kelp is dried with calcium chloride for use in producing sodium carbonate. Anhydrous calcium chloride has been approved by the FDA as a packaging aid to ensure dryness (CPG 7117.02).[31]

The hydrated salt can be dried for re-use but will dissolve in its own water of hydration if heated quickly and form a hard amalgamated solid when cooled.

Other applications

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Calcium chloride is used in concrete mixes to accelerate the initial setting, but chloride ions lead to corrosion of steel rebar, so it should not be used in reinforced concrete.[32] The anhydrous form of calcium chloride may also be used for this purpose and can provide a measure of the moisture in concrete.[33]

Calcium chloride is included as an additive in plastics and in fire extinguishers, in blast furnaces as an additive to control scaffolding (clumping and adhesion of materials that prevent the furnace charge from descending), and in fabric softener as a thinner.

The exothermic dissolution of calcium chloride is used in self-heating cans and heating pads.

Calcium Chloride is used as a water hardener in the maintenance of hot tub water, as insufficiently hard water can lead to corrosion and foaming.

In the oil industry, calcium chloride is used to increase the density of solids-free brines. It is also used to provide inhibition of swelling clays in the water phase of invert emulsion drilling fluids.

CaCl2 acts as flux material, decreasing the melting point, in the Davy process for the industrial production of sodium metal through the electrolysis of molten NaCl.

Calcium chloride is also used in the production of activated charcoal.

Calcium chloride can be used to precipitate fluoride ions from water as insoluble CaF2.

Calcium chloride is also an ingredient used in ceramic slipware. It suspends clay particles so that they float within the solution, making it easier to use in a variety of slipcasting techniques.

Calcium chloride dihydrate (20 percent by weight) dissolved in ethanol (95 percent ABV) has been used as a sterilant for male animals. The solution is injected into the testes of the animal. Within one month, necrosis of testicular tissue results in sterilization.[34][35]

Cocaine producers in Colombia import tons of calcium chloride to recover solvents that are on the INCB Red List and are more tightly controlled.[36]

Metal reduction flux

Similarly, CaCl2 is used as a flux and electrolyte in the FFC Cambridge electrolysis process for titanium production, where it ensures the proper exchange of calcium and oxygen ions between the electrodes.

Medical use

Calcium chloride infusions may be used as an intravenous therapy to prevent hypocalcemia.


Although the salt is non-toxic in small quantities when wet, the strongly hygroscopic properties of non-hydrated calcium chloride present some hazards. It can act as an irritant by desiccating moist skin. Solid calcium chloride dissolves exothermically, and burns can result in the mouth and esophagus if it is ingested. Ingestion of concentrated solutions or solid products may cause gastrointestinal irritation or ulceration.[37]

Consumption of calcium chloride can lead to hypercalcemia.[38]


Flame test of CaCl2

Calcium chloride dissolves in water, producing chloride and the aquo complex [Ca(H2O)6]2+. In this way, these solutions are sources of "free" calcium and free chloride ions. This description is illustrated by the fact that these solutions react with phosphate sources to give a solid precipitate of calcium phosphate:

3 CaCl2 + 2 PO3−4 → Ca3(PO4)2 + 6 Cl

Calcium chloride has a very high enthalpy change of solution, indicated by considerable temperature rise accompanying dissolution of the anhydrous salt in water. This property is the basis for its largest-scale application.

Molten calcium chloride can be electrolysed to give calcium metal and chlorine gas:

CaCl2 → Ca + Cl2


Structure of the polymeric [Ca(H2O)6]2+ center in crystalline calcium chloride hexahydrate, illustrating the high coordination number typical for calcium complexes.

In much of the world, calcium chloride is derived from limestone as a by-product of the Solvay process, which follows the net reaction below:[10]

2 NaCl + CaCO3 → Na2CO3 + CaCl2

North American consumption in 2002 was 1,529,000 tonnes (3.37 billion pounds).[39] In the US, most of calcium chloride is obtained by purification from brine. As with most bulk commodity salt products, trace amounts of other cations from the alkali metals and alkaline earth metals (groups 1 and 2) and other anions from the halogens (group 17) typically occur.[10]


Calcium chloride occurs as the rare evaporite minerals sinjarite (dihydrate) and antarcticite (hexahydrate).[40][41][42] Another natural hydrate known is ghiaraite – a tetrahydrate.[43][42] The related minerals chlorocalcite (potassium calcium chloride, KCaCl3) and tachyhydrite (calcium magnesium chloride, CaMg2Cl6·12H2O) are also very rare.[44][45][42] The same is true for rorisite, CaClF (calcium chloride fluoride).[46][42]

See also


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  3. ^ Seidell A, Linke WF (1919). Solubilities of Inorganic and Organic Compounds (second ed.). New York: D. Van Nostrand Company. p. 196.
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  5. ^ a b c d e f Pradyot P (2019). Handbook of Inorganic Chemicals. The McGraw-Hill Companies, Inc. p. 162. ISBN 978-0-07-049439-8.
  6. ^ a b c d Müller U (2006). Inorganic Structural Chemistry (second ed.). England: John Wiley & Sons Ltd. p. 33. ISBN 978-0-470-01864-4.
  7. ^ a b c Sigma-Aldrich Co., Calcium chloride.
  8. ^ Garrett DE (2004). Handbook of Lithium and Natural Calcium Chloride. Elsevier. p. 379. ISBN 978-0-08-047290-4. Archived from the original on 31 October 2023. Retrieved 29 August 2018. Its toxicity upon ingestion, is indicated by the test on rats: oral LD50 (rat) is 1.0–1.4 g/kg (the lethal dose for half of the test animals, in this case rats...)
  9. ^ "MSDS of Calcium chloride". Fisher Scientific. Archived from the original on 25 September 2015. Retrieved 7 July 2014.
  10. ^ a b c Robert Kemp, Suzanne E. Keegan "Calcium Chloride" in Ullmann's Encyclopedia of Industrial Chemistry 2000, Wiley-VCH, Weinheim. doi:10.1002/14356007.a04_547
  11. ^ Peck EL, Hamilton JH, Lewis JR, Hogan MB, Kusian RN, Cope WJ (1954). Proceedings of the First Annual Heating and Air Conditioning Conference: 1953-1955. University of Utah, Department of Metallurgy. Archived from the original on 15 March 2024. Retrieved 4 February 2024.
  12. ^ a b Hartmann PK (1816). Pharmacologia Dynamica: Usui Academico Adcommodata (in Latin). Kupffer et Wimmer. Archived from the original on 29 December 2023. Retrieved 29 December 2023.
  13. ^ Ottley WC (1826). A dictionary of chemistry and of mineralogy as connected with it. Murray. Archived from the original on 29 December 2023. Retrieved 29 December 2023.
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  19. ^ "Apple Caviar Technique". StarChefs Studio. April 2004. Archived from the original on 29 June 2022. Retrieved 9 August 2006.
  20. ^ Sitbon C, Paliyath G (1 January 2011). "4.28 - Pre- and Postharvest Treatments Affecting Nutritional Quality". In Moo-Young M (ed.). Comprehensive Biotechnology (Second ed.). Academic Press. pp. 349–357. doi:10.1016/B978-0-08-088504-9.00275-0. ISBN 978-0-08-088504-9. Archived from the original on 19 March 2012. Retrieved 17 March 2024 – via ScienceDirect.
  21. ^ "Why Your Bottled Water Contains Four Different Ingredients". 24 July 2014. Archived from the original on 8 February 2019. Retrieved 17 March 2024.
  22. ^ Calcium Chloride SIDS Initial Assessment Profile, UNEP Publications, SIAM 15, Boston, 22–25 October 2002, pp. 13–14.
  23. ^ 21 CFR § 184.1193
  24. ^ 7 CFR § 205.602 Archived 29 April 2021 at the Wayback Machine
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  27. ^ "Archived copy" (PDF). Archived (PDF) from the original on 16 March 2024. Retrieved 16 March 2024.((cite web)): CS1 maint: archived copy as title (link)
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  30. ^ "Cork Spot and Bitter Pit of Apples", Richard C. Funt and Michael A. Ellis,
  31. ^ "CPG 7117.02". FDA Compliance Articles. US Food and Drug Administration. March 1995. Archived from the original on 13 December 2007. Retrieved 3 December 2007.
  32. ^ "Accelerating Concrete Set Time". Federal Highway Administration. 1 June 1999. Archived from the original on 17 January 2007. Retrieved 16 January 2007.
  33. ^ National Research Council (U.S.). Building Research Institute (1962). Adhesives in Building: Selection and Field Application; Pressure-sensitive Tapes. National Academy of Science-National Research Council. pp. 24–5.
  34. ^ Koger, Nov 1977, "Calcium Chloride, Practical Necrotizing Agent", Journal of the American Association of Bovine Practitioners (USA), (Nov 1977), v. 12, p. 118–119
  35. ^ Jana K, Samanta P (2011). "Clinical evaluation of non-surgical sterilization of male cats with single intra-testicular injection of calcium chloride". BMC Vet. Res. 7: 39. doi:10.1186/1746-6148-7-39. PMC 3152893. PMID 21774835.
  36. ^ Smith M, Simpson C (26 October 2020). "Narcos Are Waging a New Drug War Over a Texas Company's Basic Chemical". Bloomberg. Archived from the original on 26 October 2020. Retrieved 26 October 2020.((cite web)): CS1 maint: bot: original URL status unknown (link)
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  40. ^ "Sinjarite". Archived from the original on 3 March 2023. Retrieved 6 November 2020.
  41. ^ "Antarcticite". Archived from the original on 1 May 2023. Retrieved 6 November 2020.
  42. ^ a b c d "List of Minerals". 21 March 2011. Archived from the original on 15 March 2013. Retrieved 6 November 2020.
  43. ^ "Ghiaraite". Archived from the original on 3 March 2023. Retrieved 6 November 2020.
  44. ^ "Chlorocalcite". Archived from the original on 30 May 2023. Retrieved 6 November 2020.
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  46. ^ "Rorisite". Archived from the original on 3 March 2023. Retrieved 6 November 2020.