Sodium

Related Topics:
Sodium :: Sodium_Fluoride :: Sodium_Tetradecyl :: Sodium_Hyaluronate :: Sodium_Polystyrene :: Sodium_Sulfactemide :: Sodium_Ferric_Gluconate :: Sodium_Phosphate_Monobasic_Monohydrate

For sodium in the diet, see Edible salt.

Sodium is the chemical element in the periodic table that has the symbol Na (Natrium in Latin) and atomic number 11. Sodium is a soft, waxy, silvery reactive metal belonging to the alkali metals that is abundant in natural compounds (especially halite). It is highly reactive, burns with a yellow flame, reacts violently with water and oxidizes in air, necessitating storage in an inert environment.

Notable characteristics

Like the other alkali metals, sodium metal is a soft, light-weight, silvery white, reactive metal. Owing to its extreme reactivity, it occurs in nature only combined in compounds, never as a pure elemental metal. Sodium metal floats on water and reacts with it violently, releasing heat and flammable hydrogen gas, and forming a solution of the strong base sodium hydroxide.

Sodium ions are necessary for regulation of blood and body fluids, transmission of nerve impulses, heart activity, and certain metabolic functions. It is widely considered that most people in Western countries consume more than is needed, in the form of sodium chloride, or table salt, and that this can have a negative effect on health. See Edible salt.

Under extreme pressure, sodium departs from standard rules for changing to a liquid state. Most materials need more thermal energy to melt under pressure than they do at normal atmospheric pressure. This is because the molecules are packed closer together and have less room to move. At a pressure of 30 gigapascals (300,000 times sea level atmospheric pressure), the melting temperature of sodium begins to drop. At around 100 gigapascals, sodium will melt near room temperature.

A possible explanation for the aberrant behavior of sodium is that this element has one free electron that is pushed closer to the other 10 electrons when placed under pressure, forcing interactions that are not normally present. While under pressure, solid sodium assumes several odd crystal structures suggesting that the liquid might have unusual properties such as superconduction or superfluidity. (Gregoryanz, et al., 2005)

Applications

Sodium in its metallic form is an essential component in the making of esters and in the manufacture of organic compounds. This alkali metal is also a component of sodium chloride (NaCl) which is vital to life. Other uses:

In certain alloys to improve their structure.
In soap, in combination with fatty acids.
To descale metal (make its surface smooth).
To purify molten metals.
In sodium vapor lamps, an efficient means of producing light from electricity.
As a heat transfer fluid in some types of nuclear reactors and inside the hollow valves of high-performance internal combustion engines.
Sodium sulphate consumption is heavily influenced by the detergent market, which accounts for over 45% of world demand. ^*

History

Sodium (English, soda) has long been recognized in compounds, but was not isolated until 1807 by Sir Humphry Davy through the electrolysis of caustic soda. In medieval Europe a compound of sodium with the Latin name of sodanum was used as a headache remedy. Sodium's symbol, Na, comes from the neo-Latin name for a common sodium compound named natrium, which comes from the Greek nítron, a kind of natural salt. As early as 1860 Kirchhoff and Bunsen noted the sensitivity that a flame test for sodium could have. Stating in Annalen der Physik und der Chemie in the paper "Chemical Analysis by Observation of Spectra": "In a corner of our 60 cu.m. room farthest away from the apparatus, we exploded 3 mg. of sodium chlorate with milk sugar while observing the nonluminous flame before the slit. After a few minutes, the flame gradually turned yellow and showed a strong sodium line that disappeared only after 10 minutes. From the weight of the sodium salt and the volume of air in the room, we easily calculate that one part by weight of air could not contain more than 1/20 millionth weight of sodium."

Occurrence and production

Sodium is relatively abundant in stars and the D spectral lines of this element are among the most prominent in star light. Sodium makes up about 2.6% by weight of the Earth's crust making it the fourth most abundant element overall and the most abundant alkali metal.

At the end of the 19th century, sodium was chemically prepared by heating sodium carbonate with carbon to 1100 °C.

Na₂CO₃ (liquid) + 2 C (solid, coke) → 2 Na (vapor) + 3 CO (gas).

It is now produced commercially through the electrolysis of liquid sodium chloride. This is done in a Down's cell in which the NaCl 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.

The production is done by a few companies in the world : - Metaux Speciaux in France in Moutiers in the French Alps, - DuPont in the USA at Niagara Falls, - a few Chinese producers appeared recently in Inner Mongolia and Henan.

Metallic sodium costs about 15 to 20 US cents per pound (US$0.30/kg to US$0.45/kg) in 1997 but reagent grade (ACS) sodium cost about US$35 per pound (US$75/kg) in 1990.

See also Sodium minerals.

Compounds

Sodium chloride or halite, better known as common salt, is the most common compound of sodium, but sodium occurs in many other minerals, such as amphibole, cryolite, soda niter and zeolite. Sodium compounds are important to the chemical, glass, metal, paper, petroleum, soap, and textile industries. Soap is generally a sodium salt of certain fatty acids.

The sodium compounds that are the most important to industry are common salt (NaCl), soda ash (Na₂C O₃), baking soda (NaHCO₃), caustic soda (NaOH), Chile saltpeter (NaNO₃), di- and tri-sodium phosphates, sodium thiosulfate (hypo, Na₂S₂O₃ · 5H₂O), and borax (Na₂B₄O₇ · 10H₂O).

See also Sodium compounds.

Isotopes

There are thirteen isotopes of sodium that have been recognized. The only stable isotope is ²³Na. Sodium has two radioactive cosmogenic isotopes (²²Na, half-life = 2.605 years; and ²⁴Na, half-life ≈ 15 hours).

Acute neutron radiation exposure (e.g., from a nuclear criticality accident) converts some of the stable ²³Na in human blood plasma to ²⁴Na. By measuring the concentration of this isotope, the neutron radiation dosage to the victim can be computed.

Sodium has been proposed as a "salting" material for nuclear weapons (cobalt is another, better-known salting material). A jacket of ²³Na, irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, would transmute into the radioactive isotope ²⁴Na with a half-life of 14.959 hours and produce approximately 2.7 MeV of gamma radiation, significantly increasing the radioactivity of the weapon's fallout for several hours. Such a weapon is not known to have ever been built, tested, or used.

Precautions

Sodium's powdered form is highly explosive in water and is a poison when uncombined or combined with many other elements. This metal should be handled carefully at all times. Sodium must be stored either in an inert atmosphere, or under a liquid hydrocarbon such as mineral oil or kerosene.

Physiology and sodium ions

Sodium ions play a diverse and important role in many physiological processes. Excitable cells, for example, rely on the entry of Na⁺ to cause a depolarization. An example of this is signal transduction in the human central nervous system.

Some potent neurotoxins, such as batrachotoxin, increase the sodium ion permeability of the cell membranes in nerves and muscles, causing a massive and irreversible depolarization of the membranes, with potentially fatal consequences.

References

Los Alamos National Laboratory – Sodium
Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.

External links