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Sodium chloride

For sodium in the diet, see salt.
Sodium chloride
IUPAC name Sodium chloride
Other names Common salt; halite; table salt
CAS number 7647-14-5
RTECS number VZ4725000
Molecular formula NaCl
Molar mass 58.442 g/mol
Appearance white or colorless crystals or powder
Density 2.16 g/cm³, solid
Melting point

801 °C

Boiling point

1465 °C (1738 K)

Solubility in water 35.9 g/100 mL (25 °C)
MSDS External MSDS
Main hazards Irritant and Might Sting
NFPA 704
R-phrases 36
S-phrases none
Flash point Non-flammable
Related Compounds
Other anions NaF, NaBr, NaI
Other cations LiCl, KCl, RbCl,
CsCl, MgCl2, CaCl2
Related salts Sodium acetate
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Sodium chloride, also known as common salt, table salt, or halite, is a chemical compound with the formula NaCl. Sodium chloride is the salt most responsible for the salinity of the ocean and of the extracellular fluid of many multicellular organisms. As the major ingredient in edible salt, it is commonly used as a condiment and food preservative. In one gram of sodium chloride, there are approximately 0.3933 grams of sodium, and 0.6067 g of chlorine.


Production and use

Salt is currently mass produced by evaporation of seawater or brine from other sources, such as brine wells and salt lakes, and by mining rock salt, called halite. In 2002, world production was estimated at 210 million metric tonnes, the top five producers being the United States (40.3 million tonnes), China (32.9), Germany (17.7), India (14.5), and Canada (12.3).[1]

While most people are familiar with the many uses of salt in cooking, they might be unaware that salt is used in a plethora of applications, from manufacturing pulp and paper to setting dyes in textiles and fabric, to producing soaps and detergents. In most of Canada and the northern USA, large quantities of rock salt are used to help clear highways of ice during winter, although "Road Salt" loses its melting ability at temperatures below -15°C to -20°C (5°F to -4°F). Sodium chloride is sometimes used as a cheap and safe desiccant due to its hygroscopic properties, making salting an effective method of food preservation historically. Even though more effective desiccants are available, few are safe for humans to ingest. It is also sometimes used to attract moisture in other applications, for example in a Salt lamp. .

Solubility of NaCl in various solvents
(g NaCl / 100 g of solvent at 25 °C)
H2O 36
Liquid ammonia 3.02
Methanol 1.4
Formic acid 5.2
Sulfolane 0.005
Acetonitrile 0.0003
Acetone 0.000042
Formamide 9.4
Dimethylformamide 0.04
Burgess, J. Metal Ions in Solution
(Ellis Horwood, New York, 1978)
ISBN 0-85312-027-7

Synthetic uses

Salt is also the raw material used to produce chlorine which itself is required for the production of many modern materials including PVC and pesticides. Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the chemical equation

2NaCl + 2H2O → Cl2 + H2 + 2NaOH

Sodium metal is produced commercially through the electrolysis of liquid sodium chloride. This is done in a Down's cell in which sodium chloride is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is more electropositive than sodium, no calcium will be formed at the cathode. This method is less expensive than the previous method of electrolyzing sodium hydroxide.

Sodium chloride is used in other chemical processes for the large-scale production of compounds containing sodium or chlorine. In the Solvay process, sodium chloride is used for producing sodium carbonate and calcium chloride. In the Mannheim process and in the Hargreaves process, it is used for the production of sodium sulfate and hydrochloric acid.

Food Uses

Main article: Salt

Salt is commonly used as a flavour enhancer and preservative for food and has been identified as one of the basic tastes. Excess salt consumption causes elevated levels of blood pressure (hypertension) in some, which in turn is associated with increased risks of heart attack and stroke. Consuming salt in excess can also dehydrate the human body. It is a popular myth that salt also has a practical use in cooking in that it raises the boiling point of water, allowing food to cook quicker, since the temperature of the surrounding water is above 100 degrees Celsius. However, the concentration of salt required to increase the boiling temperature of water by an appreciable amount is so high that it would spoil the food being cooked. Salt is often added to boiling water, but this is done solely to flavor the food being cooked in it.

Biological uses

Many microorganisms cannot live in an overly salty environment: water is drawn out of their cells by osmosis. For this reason salt is used to preserve some foods, such as smoked bacon or fish and can also be used to detach leeches that have attached themselves to feed. It has also been used to disinfect wounds. In medieval times salt would be rubbed into household surfaces as a cleansing agent.

Biological functions

In humans, a high-salt intake was demonstrated to attenuate Nitric Oxide production. Nitric oxide (NO) contributes to vessel homeostasis by inhibiting vascular smooth muscle contraction and growth, platelet aggregation, and leukocyte adhesion to the endothelium [[1]]

Crystal structure

  Sodium chloride forms crystals with cubic symmetry. In these, the larger chloride ions, shown to the right as green spheres, are arranged in a cubic close-packing, while the smaller sodium ions, shown to the right as blue spheres, fill the octahedral gaps between them.

Each ion is surrounded by six ions of the other kind. This same basic structure is found in many other minerals, and is known as the halite structure. This arrangement is known as cubic close packed (ccp).

It is held together with an ionic bond and electrostatic forces.

Salt is also known in the chemical world as a nuclear additive.

Road salt

While salt was once a scarce commodity in history, industrialized production has now made salt plentiful. About 51% of world output is now used by cold countries to de-ice roads in winter, both in grit bins and spread by winter service vehicles. This works because salt and water form an eutectic mixture. For a solution of table salt (sodium chloride, NaCl) in water, the freezing temperature becomes -21 °C (-6 °F) under controlled lab conditions.


Table salt sold for consumption today is not pure sodium chloride. In 1911 magnesium carbonate was first added to salt to make it flow more freely.[2] In 1924 trace amounts of iodine in form of sodium iodide, potassium iodide or potassium iodate were first added, to reduce the incidence of simple goiter.[3]

Salt for de-icing in the UK typically contains sodium hexacyanoferrate (II) at less than 100ppm as an anti-caking agent. In recent years this additive has also been used in table salt.

Common chemicals

Chemicals used in de-icing salts are mostly found to be sodium chloride (NaCl) or calcium chloride (CaCl2). Both are similar and are effective in de-icing roads. When these chemicals are produced, they are mined/made, crushed to fine granules, then treated with an anti-caking agent. Adding salt lowers the freezing point of the water, which allows the liquid to be stable at lower temperatures and allows the ice to melt. Alternative de-icing chemicals have also been used. Chemicals such as calcium magnesium acetate and potassium formate are being produced. These chemicals have few of the negative chemical effects on the environment commonly associated with NaCl and CaCl2.[4][5]

See also

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Sodium chloride


  1. ^ Susan R. Feldman. Sodium chloride. Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons, Inc. Published online 2005. doi:10.1002/0471238961.1915040902051820.a01.pub2
  2. ^ Morton Salt FAQ.
  3. ^ Markel H (1987). ""When it rains it pours": endemic goiter, iodized salt, and David Murray Cowie, MD". American journal of public health 77 (2): 219-29. PMID 3541654.
  4. ^ Finnish Environment Institute (1/9/2007). "Migration of alternative de-icing chemicals in aquifers (MIDAS)". Press release.
  5. ^ Finnish Environment Institute (2/10/2004). "Alternative de-icer found". Press release.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Sodium_chloride". A list of authors is available in Wikipedia.
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