sodium chloride

Sodium chloride
IUPAC name	Sodium chloride
Other names	Common salt; halite; table salt; rock salt
Identifiers
CAS number	[7647-14-5]
PubChem	5234
RTECS number	VZ4725000
ChemSpider ID	5044
Properties
Molecular formula	NaCl
Molar mass	58.443 g/mol
Appearance	Colorless/white crystalline solid
Odor	Odorless
Density	2.165 g/cm3
Melting point	801 °C (1074 K)
Boiling point	1465 °C (1738 K)
Solubility in water	35.6 g/100 mL (0 °C) 35.9 g/100 mL (25 °C) 39.1 g/100 mL (100 °C)
Solubility	soluble in glycerol, ethylene glycol, formic acid insoluble in HCl
Solubility in methanol	1.49 g/100 mL
Solubility in ammonia	2.15 g/100 mL
Refractive index (nD)	1.5442 (589 nm)
Structure
Crystal structure	Cubic (see text), cF8
Space group	Fm3m, No. 225
Coordination geometry	Octahedral (Na+) Octahedral (Cl−)
Hazards
MSDS	External MSDS
EU Index	Not listed
NFPA 704	0 1 0
Flash point	Non-flammable
LD50	3000–8000 mg/kg (oral in rats, mice, rabbits)[1]
Related compounds
Other anions	Sodium fluoride Sodium bromide Sodium iodide
Other cations	Lithium chloride Potassium chloride Rubidium chloride Caesium chloride
Supplementary data page
Structure and properties	n, εr, etc.
Thermodynamic data	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 references

Sodium chloride, also known as common salt, table salt, or halite, is an ionic 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.

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 (in million tones) being the United States (40.3), China (32.9), Germany (17.7), India (14.5) and Canada (12.3).^[2]

As well as the familiar uses of salt in cooking, salt is used in many applications, from manufacturing pulp and paper, to setting dyes in textiles and fabric, to producing soaps, detergents, and other bath products. It is the major source of industrial chlorine and sodium hydroxide, and used in almost every industry.

Sodium chloride is sometimes used as a cheap and safe desiccant because it appears to have 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.

Israeli and Jordanian salt evaporation ponds at the south end of the Dead Sea.

Mounds of salt, Salar de Uyuni, Bolivia.

Modern rock salt mine near Mount Morris, New York, United States.

Evaporation lagoons, Aigues-Mortes, France.

Synthetic uses

Sodium chloride 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 now 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.

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. It can also be used to detach leeches that have attached themselves to feed. It is also used to disinfect wounds.

Optical uses

Pure NaCl crystal is an optical compound with a wide transmission range from 200 nm to 20 um. It was often used in the infrared spectrum range and it is still used sometimes.

NaCl crystal is soft, hygroscopic and inexpensive. This limits its application to protected environment or for short term uses (prototyping). Exposed to free air NaCl optics will "rot".

Today tougher crystals like ZnSe are used instead of NaCl (for the IR spectral range).

Optical data

• Transmitivity: 92% (from 400 nm to 13μm)
• Refractive Index: 1.494 @ 10μm
• Reflection Loss: 7.5% @ 10μm (2 surfaces)
• dN/dT: -36.2 x 10-6/°C @ 0.7μm

Household uses

Since at least medieval times, people have used salt as a cleansing agent rubbed on household surfaces. It is also used in many brands of shampoo, and popularly to de-ice driveways and patches of ice.

Biological functions

In humans, a high-salt intake has long been known to generally raise blood pressure, especially in certain individuals. More recently, it 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.^[3][4]

Crystal structure

The crystal structure of sodium chloride. Each ion has six nearest neighbors, with octahedral geometry.

Sodium chloride forms crystals with face-centered 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 silver spheres, fill all the cubic gaps between them. Each ion is surrounded by six ions of the other kind; the surrounding ions are located at the vertices of a perfect octahedron.

This same basic structure is found in many other minerals and is commonly known as the halite or rock-salt crystal structure. It can be represented as a face-centered cubic (fcc) lattice with a two atom basis. The first atom is located at each lattice point, and the second atom is located half way between lattice points along the fcc unit cell edge.

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

Road salt

See Also: Magnesium chloride

While salt was once a scarce commodity in history, industrialized production has now made salt plentiful. Approximately 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. Calcium chloride is preferred over sodium chloride, since CaCl₂ releases energy upon forming a solution with water, heating any ice or snow it is in contact with. It also lowers the freezing point, depending on the concentration. NaCl does not release heat upon solution; however, it does lower the freezing point. It is also more readily available and does not have any special handling or storage requirements, unlike calcium chloride. The salinity (S) of water is measured as grams salt per kilogram (1000g) water, and the freezing temperatures are as follows.

Additives

Most 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.^[5] 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.^[6]

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.^[7][8]

See also

Wikibooks Cookbook has a recipe/module on Salt

Wikimedia Commons has media related to: Sodium chloride

• Biosalinity
• Black salt
• Halite, the mineral form of sodium chloride
• Salinity
• Salt
• Soap
• Salting the earth

References

Further reading

1. Kaufmann, Dale W., SODIUM CHLORIDE, The Production and Properties of Salt and Brine, ACS Monograph 145, American Chemical Society, Washington, D.C., 1960, 743 pages. Library of Congress Catalog Card Number 79-50778.
2. Kurlansky, Mark, SALT, A World History, Walker and Company, New York, 2002, 484 pages. ISBN 0-8027-1373-4

External links

• The Salt Manufacturers Association website
• Salt Institute website
• Salt Archive website
• Video of rotating rock salt unit cell (divx, 378kb)
• Salt United States Geological Survey Statistics and Information
• US Road Management website
• Salt Intake in Cold Weather
• Oxford MSDS
• JtBaker MSDS