Sodium

For sodium in the diet, see Edible salt.

Notable characteristics

Like the other alkali metals, sodium metal is a soft, light-weight, silvery white, reactive metal. Owing to its extreme reactivity, in nature it occurs only combined into compounds, and never as a pure elemental metal. Sodium metal floats on water, and reacts violently with it releasing heat, flammable hydrogen gas and caustic sodium hydroxide solution. Depending on the mass of sodium used and the amount of aggitation, this reaction may be explosive.

Sodium ions are necessary for regulation of blood and body fluids, transmission of nerve impulses, heart activity, and certain metabolic functions. Interestingly, sodium is needed by animals, which maintain high concentrations in their blood and extracellular fluids, but the ion is not needed by plants. A completely plant-based diet, therefore, will be very low in sodium. This requires some herbivores to obtain their sodium from salt licks and other mineral sources. The animal need for sodium is probably the reason for the highly conserved ability to taste sodium ion, as "salty." Receptors for the salty taste respond best to sodium, and otherwise only to a few other small monovalent cations (Li⁺, NH₄⁺, and to some extent also K⁺). Sodium salts without exception are soluble in water, and all of them are salty to the taste.

The most common sodium salt, sodium chloride (table salt), used for seasoning and food preservation, has been an important commodity in human activities — so much so that historically salt has been used as a medium of monetary exchange (the English word salary refers to salt).

The human requirement for sodium in the diet is less than 500 mg per day, which is typically less than a tenth as much as many diets "seasoned to taste." Most people consume far more sodium than is physiolgically needed. For certain people with salt-sensitive blood pressure, this extra intake may cause a negative effect on health. See edible salt.

Applications

File:Na-lamp-3.jpg

Sodium in its metallic form can be used to refine some reactive metals, such as zirconium and potassium, from their compounds. This alkali metal as the Na⁺ is vital to animal life. Other uses:

In certain alloys to improve their structure.
In soap, in combination with fatty acids. Sodium soaps are harder soaps than potassium soaps.
To descale metal (make its surface smooth).
To purify molten metals.
In sodium vapor lamps, an efficient means of producing light from electricity (see the picture), often used for street lighting in cities. Low-pressure sodium lamps give a distinctive yellow-orange light which consists primarily of the twin sodium D spectral lines. High-pressure sodium lamps give a more natural peach-colored light, which is spread much more widely across the spectrum.
As a heat transfer fluid in some types of nuclear reactors and inside the hollow valves of high-performance internal combustion engines.
NaCl, a compound of sodium ions and chloride ions, is an important heat transfer material.
In organic synthesis, sodium is used as a reducing agent, for example in the Birch reduction.
In chemistry, sodium is often used either alone or with potassium in an alloy, NaK as a desiccant for drying solvents. Used with benzophenone, it forms an intense blue coloration when the solvent is dry and oxygen-free.

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. The difference between the English name, Soda, and the abbreviation, Na stems from Berzelius' publication of his system of atomic symbols in Thomas Thomson's Annals of Philosophy^[1].

Sodium imparts an intense yellow color to flames. As early as 1860 Kirchhoff and Bunsen noted the high sensitivity that a flame test for sodium could give. They state 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

See also sodium minerals.

The flame test for sodium displays a brilliantly bright yellow emission due to the so called "sodium D-lines" at 588.9950 and 589.5924 nanometers.

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.

Metallic sodium cost 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.

Phase behavior under pressure

Under extreme pressure, sodium departs from common melting behavior. Most materials require higher temperatures to melt under pressure than they do at normal atmospheric pressure. This is because they expand on melting due to loser molecular packing in the liquid, and thus pressure forces equilibrium in the direction of the denser solid phase.

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 at 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)

Compounds

See also sodium 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. Hard soaps are generally sodium salt of certain fatty acids (potassium produces softer or liquid soaps).

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).

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.

Precautions

Sodium's metallic form is highly explosive in water and is a poison when uncombined with other elements. The powdered form may combust spontaneously in air or oxygen. 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.

The reaction of sodium and water is a familiar one in chemistry labs, and is reasonably safe if amounts of sodium smaller than a pencil eraser are used, and the reaction is done behind a plastic shield glass by people wearing eye protection. However, the sodium-water reaction does not scale well, and is treacherous when larger amounts of sodium are used. Larger pieces of sodium melt under the heat of the reaction, and the molten ball of metal is buoyed up by hydrogen and appears to be stably reacting with water until splashing covers more of the reaction mass, causing thermal runaway and an explosion which scatters molten sodium and lye. This behavior is unpredictable, and usually happens with sodium, because lithium is not reactive enough to do it, and potassium is so reactive that chemistry students are not tempted to try the reaction with large potassium pieces.

Sodium is much more reactive than magnesium. It burns at high temperatures and also melts, which spreads the flames and exposes even more surface area to the air.

Few common fire extinguishers work on sodium fires. Water of course exacerbates sodium fires. CO₂, foam, and Halon are all ineffective on sodium fires, which reignite when the extinquisher disappates. Among the very few materials effective on a sodium metal fire are Pyromet and Met-L-X. Pyromet is a NaCl/(NH₄)₂HPO₄ mix, with flow/anti-clump agents. It smoothers the fire, drains away heat, and melts to form an impermeable crust. This is the standard dry-powder canister fire extingusher for all classes of fires. Met-L-X is mostly sodium chloride, NaCl, with approximately 5% Saran plastic as a crust-former, and flow/anti-clumping agents. It is most commonly hand-applied, with a scoop. Other extreme fire extinguishing materials include Lith-X, a graphite based dry powder with an organophosphate flame retardant; and Na-X, a Na₂CO₃-based material.

Disposing of large quantities of sodium must be done through a licensed hazardous materials disposer. Small quantities may be neutralized carefully with ethanol (which has a much slower reaction than water), but care should be taken as the caustic products from the ethanol or methanol reaction are just as hazardous to eyes and skin as those from water. After alcohol reaction appears complete, water may then be carefully used for a final cleaning, before the reaction products are diluted and flushed down the drain.

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
Gregoryanz, E.; et al. (2005). "Melting of dense sodium". Physical Review Letters: in press. {{cite journal}}: Explicit use of et al. in: |author= (help)
Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.

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External links

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Theodore Gray's "Sodium Party"

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