Sodium

Sodium ( ) is a metallic element with a symbol Na (from Latin natrium or Arabic ناترون natrun; perhaps ultimately from Egyptian netjerj) and atomic number 11. It is a soft, silvery-white, highly reactive metal and is a member of the alkali metals within "group 1" (formerly known as ‘group IA’). It has only one stable isotope, ²³Na.

Elemental sodium was first isolated by Humphry Davy in 1807 by passing an electric current through molten sodium hydroxide. Elemental sodium does not occur naturally on Earth, because it quickly oxidizes in air

Sodium tends to form water-soluble compounds, such as halides, sulfate, nitrate, carboxylates and carbonates. There are only isolated examples of sodium compounds precipitating from water solution. However, nature provides examples of many insoluble sodium compounds such as the cryolite and the feldspars (aluminum silicates of sodium, potassium and calcium). There are other insoluble sodium salts such as sodium bismuthate NaBiO₃, sodium octamolybdate Na₂Mo₈O₂₅·4H₂O, sodium thioplatinate Na₄Pt₃S₆, sodium uranate Na₂UO₄. Sodium meta-antimonate's 2NaSbO₃·7H₂O solubility is 0.3 g/L as is the pyro form Na₂H₂Sb₂O₇·H₂O of this salt. Sodium metaphosphate NaPO₃ has a soluble and an insoluble form. the word for a natural mineral salt whose primary ingredient is hydrated sodium carbonate. Hydrated sodium carbonate historically had several important industrial and household uses later eclipsed by soda ash, baking soda and other sodium compounds.

Although sodium (sometimes called "soda" in English) has long been recognized in compounds, it was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda.

Sodium is now produced commercially through the electrolysis of liquid sodium chloride, based on a process patented in 1924. This is done in a Downs Cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less electropositive than sodium, no calcium will be formed at the anode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide.

Very pure sodium can be isolated by the thermal decomposition of sodium azide.

Sodium metal in reagent-grade sold for about $1.50/pound ($3.30/kg) in 2009 when purchased in tonne quantities. Lower purity metal sells for considerably less. The market in this metal is volatile due to the difficulty in its storage and shipping. It must be stored under a dry inert gas atmosphere or anhydrous mineral oil to prevent the formation of a surface layer of sodium oxide or sodium superoxide. These oxides can react violently in the presence of organic materials. Sodium will also burn violently when heated in air.

Smaller quantities of sodium, such as a kilogram,cost far more, in the range of $165/kg. This is partially due to the cost of shipping hazardous material.

Applications

Metallic sodium

Sodium in its metallic form can be used to refine some reactive metals, such as zirconium and potassium, from their compounds.

In certain alloys to improve their structure.

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 lines. High-pressure sodium lamps give a more natural peach-colored light, composed of wavelengths 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.

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.

The sodium fusion test uses sodium's high reactivity, low melting point, and the near-universal solubility of its compounds, to qualitatively analyze compounds.

Nuclear reactor cooling

, i. e. melting point of sodium as a function of potassium content in it (in atomic percent).]] Molten sodium is used as a coolant in some types of fast neutron reactors. It has a low neutron absorption cross section, which is required to achieve a high enough neutron flux, and has excellent thermal conductivity. Its high boiling point allows the reactor to operate at ambient pressure. However, using sodium poses certain challenges. The molten metal will readily burn in air and react violently with water, liberating explosive hydrogen. During reactor operation, a small amount of sodium-24 is formed as a result of neutron activation, making the coolant radioactive.

Sodium leaks and fires were a significant operational problem in the first large sodium-cooled fast reactors, causing extended shutdowns at the Monju Nuclear Power Plant and Beloyarsk Nuclear Power Plant.

Where reactors need to be frequently shut down, as is the case with some research reactors, the alloy of sodium and potassium called NaK is used. It melts at −11 °C, so cooling pipes will not freeze at room temperature. Extra precautions against coolant leaks need to be taken in case of NaK, because molten potassium will spontaneously catch fire when exposed to air. The phase diagram with potassium shows that the mixtures with potassium are liquid at room temperature in a wide concentration range. A compound Na₂K melts at 7 °C. The eutectic mixture with a potassium content of 77 % gives a melting point at −12.6 °C.

Sodium compounds

This alkali metal as the Na⁺ ion is vital to animal life.

In soap, as sodium salts of fatty acids. Sodium soaps are harder (higher melting) soaps than potassium soaps.

In some medicine formulations, the salt form of the active ingredient usually with sodium or potassium is a common modification to improve bioavailability.

Sodium chloride (NaCl), a compound of sodium ions and chloride ions, is an important heat transfer material.

Biological role

Maintaining body fluid volume in animals

The serum sodium and urine sodium play important roles in medicine, both in the maintenance of sodium and total body fluid homeostasis, and in the diagnosis of disorders causing homeostatic disruption of salt/sodium and water balance.

In mammals, decreases in blood pressure and decreases in sodium concentration sensed within the kidney result in the production of renin, a hormone which acts in a number of ways, one of them being to act indirectly to cause the generation of aldosterone, a hormone which decreases the excretion of sodium in the urine. As the body of the mammal retains more sodium, other osmoregulation systems which sense osmotic pressure in part from the concentration of sodium and water in the blood, act to generate antidiuretic hormone. This, in turn, causes the body to retain water, thus helping to restore the body's total amount of fluid.

There is also a counterbalancing system, which senses volume. As fluid is retained, receptors in the heart and vessels which sense distension and pressure, cause production of atrial natriuretic peptide, which is named in part for the Latin word for sodium. This hormone acts in various ways to cause the body to lose sodium in the urine. This causes the body's osmotic balance to drop (as low concentration of sodium is sensed directly), which in turn causes the osmoregulation system to excrete the "excess" water. The net effect is to return the body's total fluid levels back toward normal.

Maintaining electric potential in animal tissues

Sodium cations are important in neuron (brain and nerve) function, and in influencing osmotic balance between cells and the interstitial fluid, with their distribution mediated in all animals (but not in all plants) by the so-called Na⁺/K⁺-ATPase pump. Sodium is the chief cation in fluid residing outside cells in the mammalian body (the so-called extracellular compartment), with relatively little sodium residing inside cells. The volume of extracellular fluid is typically 15 liters in a 70 kg human, and the 50 grams of sodium it contains is about 90% of the body's total sodium content.

Role of sodium in plant biology

Although sodium is not considered an essential micronutrient in most plants, it is necessary in the metabolism of some C4 plants, e.g. Rhodes grass, amaranth, Joseph's coat, and pearl millet. Within these C4 plants, sodium is used in the regeneration of phosphoenolpyruvate (PEP) and the synthesis of chlorophyll. In addition, the presence of sodium can offset potassium requirements in many plants by substituting in several roles, such as: maintaining turgor pressure, serving as an accompanying cation in long distance transport, and aiding in stomatal opening and closing.

Due to increasing soil salinity, osmotic stress and sodium toxicity in plants, especially in agricultural crops, have become worldwide phenomena. High levels of sodium in the soil solution limit the plants' ability to uptake water due to decreased soil water potential and, therefore, may result in wilting of the plant. In addition, excess sodium within the cytoplasm of plant cells can lead to enzyme inhibition, which may result in symptoms such as necrosis, chlorosis, and possible plant death. To avoid such symptoms, plants have developed methods to combat high sodium levels, such as: mechanisms limiting sodium uptake by roots, compartmentalization of sodium in cell vacuoles, and control of sodium in long distance transport. Many plants store excess sodium in old plant tissue, limiting damage to new growth.

Dietary uses

The most common sodium salt, sodium chloride ('table salt' or 'common salt'), is used for seasoning and warm-climate food preservation, such as pickling and making jerky (the high osmotic content of salt inhibits bacterial and fungal growth). The human requirement for sodium in the diet is about 1.5 grams 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 physiologically needed. Low sodium intake may lead to sodium deficiency (hyponatremia).

Persons suffering from severe dehydration caused by diarrhea, such as that by cholera, can be treated with oral rehydration therapy, in which they drink a solution of sodium chloride, potassium chloride and glucose. This simple, effective therapy saves the lives of millions of children annually in the developing world.

==Precautions== Extreme care is required in handling elemental/metallic sodium. Sodium is potentially explosive in water (depending on quantity), and it is rapidly converted to sodium hydroxide on contact with moisture and sodium hydroxide is a corrosive substance. The powdered form may combust spontaneously in air or oxygen. Sodium must be stored either in an inert (oxygen and moisture free) atmosphere (such as nitrogen or argon), 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 by people wearing eye protection. However, the sodium-water reaction does not scale up 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 may appear to be stably reacting with water, until splashing covers more of the reaction mass, causing thermal runaway and an explosion which scatters molten sodium, lye solution, and sometimes flame. (18.5 g explosion ) This behavior is unpredictable, and among the alkali metals it is usually sodium which invites this surprise phenomenon, 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 larger potassium pieces.

Sodium is much more reactive than magnesium; a reactivity which can be further enhanced due to sodium's much lower melting point. When sodium catches fire in air (as opposed to just the hydrogen gas generated from water by means of its reaction with sodium) it more easily produces temperatures high enough to melt the sodium, exposing more of its surface to the air and spreading the fire.

Few common fire extinguishers work on sodium fires. Water, of course, exacerbates sodium fires, as do water-based foams. CO₂ and Halon are often ineffective on sodium fires, which reignite when the extinguisher dissipates. Among the very few materials effective on a sodium fire are Pyromet and Met-L-X. Pyromet is a NaCl/(NH₄)₂HPO₄ mix, with flow/anti-clump agents. It smothers the fire, drains away heat, and melts to form an impermeable crust. This is the standard dry-powder canister fire extinguisher 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+, a graphite based dry powder with an organophosphate flame retardant; and Na+, a Na₂CO₃-based material. Alternatively, plain dry sand can effectively slow down the oxygen and humidity flow to the sodium.

Because of the reaction scale problems discussed above, disposing of large quantities of sodium (more than 10 to 100 grams) must be done through a licensed hazardous materials disposer. Smaller quantities may be broken up and neutralized carefully with ethanol (which has a much slower reaction than water), or even methanol (where the reaction is more rapid than ethanol's but still less than in water), but care should nevertheless 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 the alcohol reaction appears complete, and all pieces of reaction debris have been broken up or dissolved, a mixture of alcohol and water, then pure water, may then be carefully used for a final cleaning. This should be allowed to stand a few minutes until the reaction products are diluted more thoroughly and flushed down the drain. The purpose of the final water soaking and washing of any reaction mass or container which may contain sodium, is to ensure that alcohol does not carry unreacted sodium into the sink trap, where a water reaction may generate hydrogen in the trap space which can then be potentially ignited, causing a confined sink trap explosion.

See also

Alkali metals

Sodium compounds

References

External links

Etymology of "natrium" – source of symbol Na

The Wooden Periodic Table Table's Entry on Sodium

Dietary Sodium

Sodium isotopes data from The Berkeley Laboratory Isotopes Project's

Category:Desiccants Category:Dietary minerals Category:Reducing agents Category:Chemical elements Category:Alkali metals Category:Biology and pharmacology of chemical elements Category:Nuclear reactor coolants