| IUPAC name
| Other names
|Appearance||White hygroscopic solid|
|Density|| 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)
|Melting point||775 °C (1,427 °F; 1,048 K) (anhydrous)|
|Boiling point||1,935 °C (3,515 °F; 2,208 K)|
74.5 g/100 ml (20 °C)
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)
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)
|Solubility|| Soluble in glacial acetic acid, alcohols|
Insoluble in liq. ammonia, DMSO, ethyl acetate
|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)
|Solubility in methanol|| 21.8 g/100 g (0 °C)|
29.2 g/100 g (20 °C)
38.5 g/100 g (40 °C)
|Solubility in acetone||0.01 g/100 g (20 °C)|
|Solubility in pyridine||1.66 g/100 g|
|Acidity (pKa)|| 8–9 (anhydrous)|
Std enthalpy of
| −795.42 kJ/mol (anhydrous)|
−1110.98 kJ/mol (monohydrate)
−1403.98 kJ/mol (dihydrate)
−2009.99 kJ/mol (tetrahydrate)
−2608.01 kJ/mol (hexahydrate)
|Safety data sheet||Sigma-Aldrich|
|Lethal dose or concentration (LD, LC):|
LD50 (Median dose)
|1,000 mg/kg (rats, oral)|
| Magnesium chloride|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Calcium chloride serves as a source of calcium ions for reactions. Because calcium carbonate and sulfate are more or less insoluble, calcium chloride can be used to produce any metal chloride by introduction of a soluble sulfate or carbonate to its solutions.
Molten calcium chloride can be electrolyzed to yield calcium metal and chlorine gas:
- CaCl2(l) → Ca(s) + Cl2(g)
Calcium chloride is white, solid at room temperature. The anhydrous salt is very hygroscopic, it will quickly absorb water from the air to form a solution, a property known as deliquescence. Anhydrous calcium chloride must be stored in a very dry environment.
Calcium chloride can be cheaply purchased at hardware stores, stored in desiccator bags, that contain the anhydrous form. Some road salts are another source, and can be obtained in large quantities, though these products may also contain other chlorides or impurities. The anhydrous form can be produced by strongly heating the hydrate. Often, the dried calcium chloride will harden into a solid mass, that takes time to break, during which it will rapidly absorb water from air as it cools.
Calcium chloride can be obtained by the reaction of calcium hydroxide or calcium carbonate and hydrochloric acid. If these compounds are hard to come by, clean white limestone, seashells, or well-cleaned egg shells can be dissolved in hydrochloric acid, and the solute filtered. Calcium carbonate is then reprecipitated by addition of sodium carbonate or sodium bicarbonate, and finally dissolved in hydrochloric acid and filtered again. It is highly recommended that it is prepared strictly from a pure calcium salt though, as this process is time-consuming, costly, and fails to remove magnesium or iron impurities.
- Electrolysis of molten, anhydrous calcium chloride to produce calcium metal.
- Make ammonium chloride
- Building a homemade desiccator
- Make a dry box
- Dry solvents
Anhydrous calcium chloride can be an irritant due to its strong desiccating properties. Otherwise, it is only minimally toxic.
Calcium chloride should be stored in sealed containers or bags, as it is extremely hygroscopic. Zipper bags are also good to store anhydrous CaCl2.
No special disposal is required. Calcium chloride is non-toxic to the environment, and it's even used as anti-icing during the winter.