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{{for|sodium in the diet|Edible salt}}
{{Infobox sodium}}
'''Sodium''' ({{pronEng|ˈsoʊdiəm}}) is a [[chemical element]] which has the symbol '''Na''' (Latin: ''natrium''), atomic number 11, atomic mass 22.9898 g/mol, common [[oxidation number]] +1. Sodium is a soft, [[silver]]y [[white]], highly reactive element and is a member of the [[alkali metals]] within "group 1" (formerly known as ‘group IA’). It has only one stable [[isotope]], <sup>23</sup>Na. Sodium was first isolated by Sir [[Humphry Davy]] in 1807 by passing an electric current through molten [[sodium hydroxide]]. Sodium quickly oxidizes in [[air]] and is violently reactive with [[water]], so it must be stored in an inert medium, such as [[kerosene]]. Sodium is present in great quantities in the earth's oceans as [[sodium chloride]]. It is also a component of many minerals, and it is an [[Dietary mineral|essential element]] for animal life. As such, it is classified as a “dietary inorganic macro-mineral.” 

== Notable characteristics ==
Compared with other alkali metals, sodium is generally less reactive than [[potassium]] and more reactive than [[lithium]],<ref>{{cite web|author=Prof. N. De Leon|title=Reactivity of Alkali Metals|url=http://www.iun.edu/~cpanhd/C101webnotes/modern-atomic-theory/alkali-reac.html|publisher=Indiana University Northwest|accessdate=2007-12-07}}</ref> in accordance with "[[periodic law]]": for example, their reaction in water, chlorine gas, etc.; the reactivity of their [[nitrate]]s, [[chlorate]]s, [[perchlorate]]s, etc. The density of elements generally increase with increasing [[atomic number]], but potassium is less dense than sodium.

Owing to its high reactivity, sodium is found in nature only as a compound and never as the free element. Sodium reacts [[Exothermic reaction|exothermically]] with water: small pea-sized pieces will bounce across the surface of the water until they are consumed by it, whereas large pieces will explode. While sodium reacts with water at room temperature, the sodium piece melts with the heat of the reaction to form a sphere, if the reacting sodium piece is large enough. The reaction with water produces very caustic [[sodium hydroxide]] (lye) and highly flammable hydrogen gas. These are extreme hazards (see ''Precautions'' section below). When burned in air, sodium forms [[sodium peroxide]] Na<sub>2</sub>O<sub>2</sub>, or with limited [[oxygen]], the oxide Na<sub>2</sub>O</sub> (unlike [[lithium]], the [[nitride]] is not formed). If burned in oxygen under pressure, [[sodium superoxide]] NaO<sub>2</sub> will be produced.

When sodium or its compounds are introduced into a flame it will contribute a bright yellow colour.

At room temperature, pure sodium is so soft that it can be easily cut with a butter knife. In air, the bright silvery luster of freshly exposed sodium will rapidly tarnish.

In chemistry, most sodium compounds are considered soluble but nature provides examples of many insoluble sodium compounds such as the feldspars. There are other insoluble sodium salts such as [[sodium bismuthate]] NaBiO<sub>3</sub>, [[sodium octamolybdate]] Na<sub>2</sub>Mo<sub>8</sub>O<sub>25</sub>• 4H<sub>2</sub>O, [[sodium thioplatinate]] Na<sub>4</sub>Pt<sub>3</sub>S<sub>6</sub>, sodium uranate Na<sub>2</sub>UO<sub>4</sub>. [[Sodium meta-antimonate]]'s 2NaSbO<sub>3</sub>•7H<sub>2</sub>O solubility is 0.3g/L as is the [[pyro form]] Na<sub>2</sub>H<sub>2</sub>Sb<sub>2</sub>O<sub>7</sub>• H<sub>2</sub>O of this salt. [[Sodium metaphosphate]] NaPO<sub>3</sub> has a soluble and an insoluble form.<ref>Lange's Handbook of Chemistry</ref> 

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]] the sodium ion as "salty." Receptors for the pure salty taste respond best to sodium, otherwise only to a few other small monovalent cations (Li<sup>+</sup>, NH<sub>4</sub><sup>+</sup>, and somewhat to K<sup>+</sup>). [[Calcium]] ion (Ca<sup>2+</sup>) also tastes salty and sometimes bitter to some people but like potassium, can trigger other tastes.

The most common sodium salt, sodium chloride ([[table salt]]), used for seasoning (for example the English word "salad" refers to salt) and warm-climate food preservation, such as [[pickling]] and making [[Jerky (food)|jerky]] (the high osmotic content of salt inhibits bacterial and fungal growth). As such, salt has been an important commodity in human activities (the English word ''salary'' refers to ''salarium'',  the wafers of salt sometimes given to roman solders along with their other wages). The human requirement for sodium in the diet is about 500&nbsp;mg per day,<ref>{{cite book |title=Implementing recommendations for dietary salt reduction: Where are we? |publisher=DIANE Publishing |isbn=1428929096 }}</ref> 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. For certain people with salt-sensitive blood pressure, this extra intake may cause a negative effect on health.

== Applications ==
[[Image:Na-lamp-3.jpg|thumb|left|150px|A low pressure sodium lamp, glowing with the light of sodium D spectral lines.]]

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<sup>+</sup> ion is vital to animal life. Other uses:
*In certain [[alloy]]s to improve their structure.
*In [[soap]], in combination with [[fatty acid]]s. Sodium soaps are harder (higher melting) soaps than potassium soaps.
*To descale metal (make its surface smooth).
*To purify molten metals.

*In [[sodium vapor lamp]]s, 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 [[Fraunhofer lines|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 [[valve]]s of high-performance  [[internal combustion engine]]s.
*[[NaCl]], a compound of sodium ions and [[chloride]] ions, is an important [[heat transfer]] material.
*In [[organic synthesis]], sodium is used as a [[Redox|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 ==
[[Image:Flametest--Na.swn.jpg|thumb|left|200px|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&nbsp;nanometers.]]

Sodium (sometimes called "soda" in [[English language|English]]) 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. The name sodium probably originates from the Arabic word suda meaning headache as the headache curing properties of sodium carbonate or soda were well known in early times.<ref>Chemical Elements by David E Newton ISBN 0-7876-2847-6</ref>

Sodium's chemical abbreviation ''Na'' was first published by [[Jöns Jakob Berzelius]] in his system of atomic symbols (Thomas Thomson's ''Annals of Philosophy''<ref>{{cite web|url=http://www.vanderkrogt.net/elements/elem/na.html|title=Elementymology &amp; Elements Multidict by Peter van der Krogt|accessdate=2007-06-08}}</ref>) and is a contraction of the element's new [[Latin]] name ''natrium'' which refers to the Egyptian ''[[natron]]''<ref>Chemical Elements by David E Newton ISBN 0-7876-2847-6</ref> word for a natural mineral salt whose primary ingredient is hydrated sodium carbonate. Which historically had several important industrial and household uses later eclipsed by [[soda ash]], [[baking soda]] and other sodium compounds. 

Sodium imparts an intense yellow color to flames. As early as 1860, [[Gustav Kirchhoff|Kirchhoff]] and [[Robert Bunsen|Bunsen]] noted the high sensitivity that a flame test for sodium could give. They state in [[Annalen der Physik|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&nbsp;mg. of sodium chlorate with milk sugar while observing the nonluminous flame before the slit. After a while, it glowed a bright 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 ==
{{seealso|Category:Sodium minerals}}

Sodium is relatively abundant in [[stars]] and the D [[spectral line]]s 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. Though elemental sodium has a rather high vaporization temperature, its relatively high abundance and very intense spectral lines have allowed its presence to be detected by ground telescopes and confirmed by spacecraft ([[Mariner 10]] and [[MESSENGER]]) in the thin atmosphere of the planet [[Mercury (planet)| Mercury]].<ref>[http://findarticles.com/p/articles/mi_m1200/is_v128/ai_3898126] accessed Feb. 5, 2008</ref>

== Commercial production ==

At the end of the 19th century, sodium was chemically prepared by heating [[sodium carbonate]] with [[carbon]] to 1100&nbsp;°C.
:Na<sub>2</sub>CO<sub>3</sub> (liquid) + 2 C (solid) → 2 Na (vapor) + 3 CO (gas).

It is now produced commercially through the [[electrolysis]] of liquid [[sodium chloride]].<ref name="pauling">Pauling, Linus, ''General Chemistry'', 1970 ed., Dover Publications</ref><ref name="losal">{{cite web|url=http://periodic.lanl.gov/elements/11.html|title=Los Alamos National Laboratory – Sodium|accessdate=2007-06-08}}</ref> This is done in a [[Downs Cell]] in which the NaCl is mixed with [[calcium chloride]] to lower the [[melting point]] below 700&nbsp;°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|Castner process]] of electrolyzing [[sodium hydroxide]].

Very pure sodium can be isolated by the thermal decomposition of [[sodium azide]].<ref>Merck Index, 9th ed., monograph 8325</ref>

Metallic sodium costs about 15 to 20 US cents per [[Pound (mass)|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 looser molecular packing in the liquid, and thus pressure forces equilibrium in the direction of the denser solid phase.

At a pressure of 30 [[gigapascal]]s (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 structure]]s suggesting that the liquid might have unusual properties such as [[Superconductivity|superconduction]] or [[superfluid]]ity.<ref name="gregoryanz">Gregoryanz, E., ''et al.'' (2005). "Melting of dense sodium". ''Physical Review Letters'' '''94''': 185502</ref>

== Compounds == 
{{seealso|Category:Sodium compounds}} 
 
[[Sodium chloride]] or [[halite]], better known as common salt, is the most common compound of sodium, but sodium occurs in many other [[mineral]]s, 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 industries are common salt (Na[[chlorine|Cl]]), [[soda ash]] (Na<sub>2</sub>[[carbon|C]][[oxygen|O]]<sub>3</sub>), [[baking soda]] (Na[[hydrogen|H]]CO<sub>3</sub>), [[sodium hydroxide|caustic soda]] (NaOH), [[sodium nitrate|sodium nitrate]] (Na[[nitrogen|N]]O<sub>3</sub>), di- and tri-sodium phosphates, [[sodium thiosulfate]] (hypo, Na<sub>2</sub>[[sulfur|S]]<sub>2</sub>O<sub>3</sub> · 5H<sub>2</sub>O), and [[borax]] (Na<sub>2</sub>[[boron|B]]<sub>4</sub>O<sub>7</sub> · 10H<sub>2</sub>O).

== Isotopes ==
{{main|Isotopes of sodium}}
There are thirteen [[isotope]]s of sodium that have been recognized. The only stable isotope is <sup>23</sup>Na. Sodium has two [[radioactive]] [[cosmogenic]] isotopes (<sup>22</sup>Na, [[half-life]] = 2.605 years; and <sup>24</sup>Na, half-life ≈ 15 hours).

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

== Precautions ==<!-- This section is linked from [[Potassium]] -->
Extreme care is required in handling elemental/metallic sodium. Sodium is potentially explosive in water (depending on quantity) and is a caustic poison, since it is rapidly converted to sodium hydroxide on contact with moisture. 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 [[Chemical reaction|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 [http://video.google.de/videoplay?docid=-2158222101210607510&q=sodium]) 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. CO<sub>2</sub> and [[Halon 1211|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/(NH<sub>4</sub>)<sub>2</sub>HPO<sub>4</sub> 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)|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 Na<sub>2</sub>CO<sub>3</sub>-based material.

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 soak and wash of any reaction mass 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.

== Physiology and sodium ions ==
{{main|action potential}}

Sodium ions play a diverse and important role in many physiological processes. Excitable animal cells, for example, rely on the entry of Na<sup>+</sup> to cause a [[depolarization]]. An example of this is [[signal transduction]] in the human [[central nervous system]], which depends on sodium ion motion across the nerve cell membrane, in all nerves.

Some potent [[neurotoxin]]s, such as [[batrachotoxin]], increase the sodium ion permeability of the [[cell membrane]]s in nerves and muscles, causing a massive and irreversible [[depolarization]] of the membranes, with potentially fatal consequences. However, drugs with smaller effects on sodium ion motion in nerves may have diverse pharmacological effects which range from anti-depressant to anti-seizure actions.

{{main|hyponatremia|hypernatremia|diuretic|vasopressin}}

Sodium is the primary [[cation]] (positive ion) in extracellular fluids in animals and humans. These fluids, such as blood plasma and extracellular fluids in other tissues, bathe cells and carry out transport functions for nutrients and wastes. Sodium is also the principal cation in seawater, although the concentration there is about 3.8 times what it is normally in extracellular body fluids.

Although the system for maintaining optimal salt and water balance in the body is a complex one, one of the primary ways in which the human body keeps track of loss of body water is that [[osmoreceptor]]s in the [[hypothalamus]] sense a balance of sodium and water concentration in extracellular fluids. Relative loss of body water will cause sodium concentration to rise higher than normal, a condition known as [[hypernatremia]]. This ordinarily results in thirst. Conversely, an excess of body water caused by drinking will result in too little sodium in the blood ([[hyponatremia]]), a condition which is again sensed by the [[hypothalamus]], causing a decrease in [[vasopressin]] hormone secretion from the [[posterior pituitary]], and a consequent loss of water in the urine, which acts to restore blood sodium concentrations to normal.

Severely dehydrated persons, such as people rescued from ocean or desert survival situations, usually have very high blood sodium concentrations. These must be very carefully and slowly returned to normal, since too-rapid correction of hypernatremia may result in brain damage from cellular swelling, as water moves suddenly into cells with high [[osmolar]] content.

Because the [[hypothalamus]]/[[osmoreceptor]] system ordinarily works well to cause drinking or urination to restore the body's sodium concentrations to normal, this system can be used in medical treatment to regulate the body's total fluid content, by first controlling the body's sodium content. Thus, when a powerful [[diuretic]] drug is given which causes the kidneys to excrete sodium, the effect is accompanied by an excretion of body water (water loss accompanies sodium loss). This happens because the kidney is unable to efficiently retain water while excreting large amounts of sodium. In addition, after sodium excretion, the [[osmoreceptor]] system may sense lowered sodium concentration in the blood and then direct compensatory urinary water loss in order to correct the [[hyponatremia|hyponatremic]] (low blood sodium) state.

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 <ref>Relationship between Salt Intake, Nitric Oxide
and Asymmetric Dimethylarginine and Its Relevance to Patients with End-Stage Renal Disease, Tomohiro Osanai, Naoto Fujiwara, Masayuki Saitoh, Satoko Sasaki, Hirofumi Tomita, Masayuki Nakamura, Hiroshi Osawa, Hideaki Yamabe, Ken Okumura, 2002, http://content.karger.com/ProdukteDB/produkte.asp?Aktion=ShowPDF&ProduktNr=223997&Ausgabe=228460&ArtikelNr=63555</ref>

== Atomic spectral lines ==
[[Image: Na-lamp-2.jpg|left|thumb|150px|A Low Pressure Sodium / Sodium Oxide (LPS/SOX) [[streetlamp]] at full power (detail)]]
[[Image:Starfield Optical Range - sodium laser.jpg|thumb|right|200px|A [[FASOR (laser physics)|FASOR]] tuned to the D2A component of the sodium D line, used at the [[Starfire Optical Range]] to excite sodium atoms [[sodium layer|in the upper atmosphere]].]]

One notable [[atomic spectral line]] of [[sodium vapor]] is the so-called D-line, which may be observed directly as the sodium flame-test line (see above) and also the major light output of low-pressure [[sodium lamp]]s (these produce an unnatural yellow, rather than the peach-colored glow of high pressure lamps). The D-line is one of the classified [[Fraunhofer lines]] observed in the visible spectrum of the sun's [[electromagnetic radiation]]. Sodium vapor in the upper layers of the sun creates a dark line in the emitted spectrum of electromagnetic radiation by absorbing visible light in a band of wavelengths around 589.5&nbsp;nm. This wavelength corresponds to transitions in atomic sodium in which the valence-electron transitions from a 3p to 3s [[electronic state]]. Closer examination of the visible spectrum of atomic sodium reveals that the D-line actually consists of two lines called the D<sub>1</sub> and D<sub>2</sub> lines at 589.8&nbsp;nm and 589.2&nbsp;nm, respectively. This [[fine structure]] results from a [[spin-orbit interaction]] of the valence electron in the 3p [[electronic state]]. The spin-orbit interaction couples the [[spin angular momentum]] and [[orbital angular momentum]] of a 3p electron to form two states that are respectively notated as <math>3p(^2P^o_{1/2})</math> and <math>3p(^2P^o_{3/2})</math> in the [[LS coupling]] scheme. The 3s state of the electron gives rise to a single state which is notated as <math>3s(^2S_{1/2})</math> in the LS coupling scheme. The D<sub>1</sub>-line results from an electronic transition between <math>3s(^2S_{1/2})</math> lower state and <math>3p(^2P^o_{1/2})</math> upper state. The D<sub>2</sub>-line results from an electronic transition between <math>3s(^2S_{1/2})</math> lower state and <math>3p(^2P^o_{3/2})</math> upper state. Even closer examination of the visible spectrum of atomic sodium would reveal that the D-line actually consists of a lot more than two lines. These lines are associated with [[hyperfine structure]] of the 3p upper states and 3s lower states. Many different transitions involving visible light near 589.5&nbsp;nm may occur between the different upper and lower hyperfine levels.<ref name="Citron-PRL-1977">{{cite journal| author= Citron, M. L., et al.| date=1977| title=Experimental study of power broadening in a two level atom | journal=Physical Review A | volume=16| page=1507 | doi=10.1103/PhysRevA.16.1507  | url=http://prola.aps.org/abstract/PRA/v16/i4/p1507_1}}</ref><ref>{{cite web| author=Daniel A. Steck| title=Sodium D. Line Data | url=http://george.ph.utexas.edu/~dsteck/alkalidata/sodiumnumbers.pdf| publisher=Los Alamos National Laboratory (technical report)}}</ref>


== Sodium Atmospheres==
One surprising and little known fact is that the planet Mercury has a Sodium atmosphere. This was determined by ground based observations. Apparently, the solar wind scours the surface of this planet and kicks up lots of ions of sodium. Mercury's sodium atmosphere is dynamic and constantly being replenished. Europa also seems to have an extended sodium atmosphere. 
== See Also ==
*[[:Category:Sodium compounds|Sodium compounds]]
*[[:Category:Alkali metals|Alkali metals]]
{{alkalimetals}}

== References ==
{{reflist}}
*Rebecca J. Donatelle. Health, The Basics. 6th ed. San Francisco: Pearson Education, Inc. 2005.


== External links ==
{{Commons|Sodium}}
{{wiktionary|sodium}}
*[http://www.webelements.com/webelements/elements/text/Na/index.html WebElements.com – Sodium]
*[http://www.theodoregray.com/PeriodicTable/Elements/011/index.html The Wooden Periodic Table Table's Entry on Sodium]
*[http://www.americanheart.org/presenter.jhtml?identifier=4708 Dietary Sodium]
*[http://ie.lbl.gov/education/parent/Na_iso.htm Sodium isotopes data from ''The Berkeley Laboratory Isotopes Project's'']

[[Category:Desiccants]]
[[Category:Dietary minerals]]
[[Category:Sodium|Sodium]]
[[Category:Reducing agents]]

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