calcium chloride

calcium chloride

calcium chloride, CaCl2, chemical compound that is crystalline, lumpy, or flaky, is usually white, and is very soluble in water. The anhydrous compound is hygroscopic; it rapidly absorbs water and is used to dry gases by passing them through it. Calcium chloride is commercially available usually as the dihydrate, CaCl2·2H2O; it is used to melt ice on roads, to control dust, in brines for refrigeration, and as a preservative in foods. It is also used in the monohydrate and hexahydrate forms. Calcium chloride is a byproduct of the Solvay process (a major source of the compound) and is present in natural brines.
| Section8 = }} Calcium chloride (CaCl2), is an ionic compound of calcium and chlorine. It is highly soluble in water and it is deliquescent. It is a salt that is solid at room temperature, and it behaves as a typical ionic halide. It has several common applications such as brine for refrigeration plants, ice and dust control on roads, and in concrete. It can be produced directly from limestone, but large amounts are also produced as a by-product of the Solvay process. Because of its hygroscopic nature, it must be kept in tightly-sealed containers.

Chemical properties

Calcium Chloride can serve as a source of calcium ions in solution, for instance for precipitation because many calcium compounds are insoluble:

3 CaCl2(aq) + 2 K3PO4(aq) → Ca3(PO4)2(s) + 6 KCl(aq)

Molten CaCl2 can be electrolysed to give calcium metal:

CaCl2(l) → Ca(s) + Cl2(g)

Uses in industry

Millions of tonnes of calcium chloride are made each year in the US alone, and it has a wide variety of industrial applications:

Because it is strongly hygroscopic, air or other gases may be channeled through a column of calcium chloride to remove moisture. In particular, calcium chloride is usually used to pack drying tubes to exclude atmospheric moisture from a reaction set-up while allowing gases to escape. It can also be added to liquids to remove suspended or dissolved water. The dissolving process is highly exothermic and rapidly produces temperatures of around 60 °C (140 °F). In this capacity, it is known as a drying agent or desiccant. It is converted to a brine as it absorbs the water or water vapor from the substance to be dried:

CaCl2 + 2 H2O → CaCl2·2H2O

Aided by the intense heat evolved during its dissolution, calcium chloride is also used as an ice-melting compound. Unlike the more-common sodium chloride (rock salt or halite), it is relatively harmless to plants and soil; however, recent observations in Washington state suggest it may be particularly harsh on roadside evergreen trees. It is also more effective at lower temperatures than sodium chloride. When distributed for this use, it usually takes the form of small white balls a few millimetres in diameter, called prills (see picture at top of page).

Used for its hygroscopic property, it can be applied to keep a liquid layer on the surface of the roadway, which holds dust down. It is used in concrete mixes to help speed up the initial setting, however chloride ion leads to corrosion of steel rebar, so it should not be used in reinforced concrete.

Aqueous calcium chloride (in solution with water) lowers the freezing point as low as -52°C (-62°F), making it ideal for filling agricultural implement tires as a liquid ballast, aiding traction in cold climates.

Other industrial applications include use as an additive in plastics, as a drainage aid for wastewater treatment, as an additive in fire extinguishers, as an additive in control scaffolding in blast furnaces, and as a thinner in fabric softener.

North American consumption in 2002 was 1,687,000 tons (3.7 billion pounds). A Dow Chemical Company manufacturing facility in Michigan, houses about 35% of the total U.S. production capacity for calcium chloride.

Uses in food

As an ingredient, it is listed as a permitted food additive in the European Union for use as a sequestrant and firming agent with the E number E509, and considered as generally recognized as safe (GRAS) by the U.S. Food and Drug Administration. (21 CFR § 184.1193) The average intake of calcium chloride as food additives has been estimated to be 160-345 mg/day for individuals. Ingestion of concentrated or pure calcium chloride products may cause gastrointestinal irritation or ulceration. The anhydrous form has been approved by the FDA as a packaging aid to ensure dryness (CPG 7117.02).

Calcium chloride is commonly used as an electrolyte and has an extremely salty taste, as found in sports drinks and other beverages such as Smartwater and Nestle bottled water. It can also be used as a preservative to maintain firmness in canned vegetables or in higher concentrations in pickles to give a salty taste while not increasing the food's sodium content. It is even found in snack foods, including Cadbury Caramilk chocolate bars (purpose unknown).

It can be used to make a caviar substitute from vegetable or fruit juices or added to processed milk to restore the natural balance between calcium and protein for the purposes of making cheese such as brie and stilton. Calcium chloride's exothermic properties are exploited in many 'self heating' food products where it is activated (mixed) with water to start the heating process, providing a non-explosive, dry fuel that is easily activated.

In brewing beers (esp. ales and bitters), calcium chloride is sometimes used to correct mineral deficiencies in the brewing water (calcium is important for enzyme function during the mash, for kettle protein coagulation (the "hot break") and yeast metabolism) and adds permanent hardness to the water. The chloride ions enhance flavour and give a perception of sweetness and fuller flavour, whereas the sulfate ions in Gypsum, which is also used to add calcium ions to brewing water, tend to impart a drier, crisper flavour with more bitterness.

Uses in drugs

Calcium chloride can be injected as intravenous therapy for the treatment of hypocalcaemia (low serum calcium). It can be used for insect bites or stings (such as Black Widow Spider bites); sensitivity reactions, particularly when characterized by urticaria (hives); magnesium intoxication; as an aid in management of the acute symptoms in lead colic; in cardiac resuscitation, particularly after open heart surgery. Parenteral calcium can be used when epinephrine has failed to improve weak or ineffective myocardial contractions. Calcium chloride injection may antagonize cardiac toxicity as measured by electrocardiogram (ECG/EKG).

It can help to protect the myocardium from dangerously-high levels of serum potassium in hyperkalemia. Calcium chloride can be used to quickly treat Calcium Channel Blocker toxicity, from the side effects of drugs such as Diltiazem (Cardizem)—helping avoid potential heart attacks.

The aqueous form of calcium chloride is used in genetic transformation of cells by increasing the cell membrane permeability, inducing competence for DNA uptake (allowing DNA fragments to enter the cell more readily).

It can also be used in the reef aquarium hobby for adding bio-available calcium in solution for calcium-using animals such as algae, snails, hard tube worms, and corals although the use of calcium hydroxide (kalkwasser mix) or a calcium reactor is the preferred method of adding calcium. However, calcium chloride is the quickest method to increase calcium levels as it dissolves readily in water.


Calcium chloride is an irritant, particularly on moist skin. Wear gloves and goggles or a full face shield to protect hands and eyes; avoid inhalation.

Dry calcium chloride reacts exothermically when exposed to water. Burns can result in the mouth and esophagus if humans or other animals ingest dry calcium chloride pellets. Small children are more susceptible than adults (who generally have had experience trying to eat hot food, and can react accordingly) so calcium chloride pellets should be kept out of their reach.

Natural occurence

Natural occurrence of a dihydrate (mineral sinjarite) and hexahydrate (antarcticite) is very rare and connected mainly with dry lakes and brines. Chlorocalcite KCaCl3 is a related mineral (also very rare).


General references

  • Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.

External links

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