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{{about|the chemical element|the nutrient commonly called "sodium"|salt|the PlayStation Home game|Sodium (PlayStation Home)|the racehorse|Sodium (horse)}}
[[File:Stop hand nuvola black.svg|1500px|alt=Stop icon]]{{uw-vandalism4im}}
{{Redirect|Natrium|the town in West Virginia|Natrium, West Virginia}}
{{Infobox sodium}}
'''Sodium''' {{IPAc-en|ˈ|s|oʊ|d|i|ə|m}}<ref>{{citation|last=Wells|first=John C.|date=2008|title=Longman Pronunciation Dictionary|edition=3rd|publisher=Longman|page=755|isbn=9781405881180}}</ref> is a [[chemical element]] with symbol&nbsp;'''Na''' (from Ancient Greek Νάτριο) and [[atomic number]]&nbsp;11. It is a soft, silver-white, highly reactive [[metal]]. In the Periodic table it is in column 1 ([[alkali metals]]), and shares with the other six elements in that column that it has a single electron in its outer shell, which it readily donates, creating a positively charged atom - a [[cation]]. Its only stable [[isotope]] is <sup>23</sup>Na. The free metal does not occur in nature, but instead must be prepared from its compounds. Sodium is the sixth most abundant [[Abundance of elements in Earth's crust|element in the Earth's crust]], and exists in numerous [[minerals]] such as [[feldspar]]s, [[sodalite]] and [[halite|rock salt]] (NaCl). Many salts of sodium are highly water-soluble: sodium ions have been leached by the action of water from the earth's minerals over eons, so that sodium (and chlorine) are the most common dissolved elements by weight in the oceans.

Sodium was first isolated by [[Humphry Davy]] in 1807 by the [[electrolysis]] of [[sodium hydroxide]]. Many sodium compounds are useful, such as [[sodium hydroxide]] ([[lye]]) for [[Soap|soap-making]] and [[sodium chloride]] for use as a de-icing agent and a nutrient ([[edible salt]]).

Sodium is an [[Dietary mineral|essential element]] for all animals and some plants. Sodium ions are the major cation in the [[extracellular fluid]] (ECF) and as such are the major contributor to the ECF [[osmotic pressure]], and thus ECF compartment volume. Loss of only water from the ECF compartment increases the sodium concentration: [[hypernatremia]]. [[Isotonic]] loss of water and sodium from the ECF compartment decreases the size of that compartment: ECF [[hypovolemia]].

Sodium ions are pumped out of cells by [[Na+/K+-ATPase]] in exchange for potassium. It pumps three sodium ions out of the cell for every two potassium ions pumped in, contributing to the fact that about forty times as much potassium is inside cells compared to outside, and about ten times as much sodium is outside cells compared to inside. In nerve cells, a buildup of electrical charge across the cell membrane allows transmission of a nerve impulse - an [[action potential]] - when the charge is dissipated.

==Characteristics==

===Physical===
Sodium at [[standard temperature and pressure]] is a soft silvery metal that oxidizes to grayish white unless immersed in oil or inert gas. Sodium can be readily cut with a knife and is a good conductor of electricity and heat. These properties change dramatically at elevated pressures: at 1.5 [[Bar (unit)|Mbar]], the color changes from silvery metallic to black; at 1.9 Mbar the material becomes transparent with a red color; and at 3 Mbar sodium is a clear and transparent solid. All of these high-pressure [[allotrope]]s are insulators and [[electride]]s.<ref>{{cite journal|last1=Gatti|first1=M.|last2=Tokatly|first2=I.|last3=Rubio|first3=A.|date=2010|title=Sodium: A Charge-Transfer Insulator at High Pressures|journal=[[Physical Review Letters]]|volume=104|issue=21|page=216404|pmid=20867123|doi=10.1103/PhysRevLett.104.216404|bibcode=2010PhRvL.104u6404G|arxiv = 1003.0540 }}</ref>

When sodium or its [[Sodium compounds|compounds]] are introduced into a flame, they turn it yellow,<ref>{{cite book|last=Schumann|first=Walter|title=Minerals of the World|date=5 August 2008|publisher=Sterling|isbn=978-1-4027-5339-8|edition=2nd|page=28|oclc=637302667}}</ref> because the excited [[atomic orbital|3s]] electrons of sodium emit a [[photon]] when they fall from 3p to 3s; the wavelength of this photon corresponds to the [[D-line|D line]] at 589.3&nbsp;nm. [[Spin-orbit interaction]]s involving the electron in the 3p orbital split the D line into two; [[hyperfine structure]]s involving both orbitals cause many more lines.<ref name="Citron-PRL-1977">{{cite journal|last1=Citron|first1=M. L.|last2=Gabel|first2=C.|last3=Stroud|first3=C.|date=1977|title=Experimental Study of Power Broadening in a Two-Level Atom|journal=Physical Review A|volume=16|doi=10.1103/PhysRevA.16.1507|page=1507|issue=4|bibcode=1977PhRvA..16.1507C|last4=Stroud|first4=C.}}</ref>

===Chemical===
[[File:Na-D-sodium D-lines-589nm.jpg|thumb|left|[[Emission spectrum]] for sodium, showing the [[Fraunhofer lines|D line]].]]
[[File:Flametest--Na.swn.jpg|thumb|left|upright|A positive [[flame test]] for sodium has a bright yellow color.]]
When freshly cut, sodium has a bright, silvery luster. If exposed to air, the surface rapidly tarnishes, darkening at first and then forming a white coating of [[sodium hydroxide]] and [[sodium carbonate]].

Sodium is generally less reactive than [[potassium]] and more reactive than [[lithium]].<ref>{{cite web|last=De Leon|first=N|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> Like all the [[alkali metal]]s, it reacts [[Exothermic reaction|exothermically]] with water, to the point that sufficiently large pieces melt to a sphere and may explode; this reaction produces caustic soda ([[sodium hydroxide]]) and flammable [[hydrogen]] gas. When burned in dry air, it mainly forms [[sodium peroxide]] as well as some [[sodium oxide]]. In moist air, sodium hydroxide results.<ref name="Greenwood">{{Greenwood&Earnshaw2nd}}</ref> Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts.<ref>{{cite book|last1=Atkins|first1=Peter W.|last2=de Paula|first2=Julio|title=Physical Chemistry|date=2002|edition=7th|publisher=W. H. Freeman|isbn=978-0-7167-3539-7|oclc=3345182}}</ref> Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy.<ref name="Greenwood"/> However, [[potassium]] and [[lithium]] have even more negative potentials.<ref>{{cite book|last=Davies|first=Julian A.|title=Synthetic Coordination Chemistry: Principles and Practice|date=1996|publisher=World Scientific|isbn=978-981-02-2084-6|oclc=717012347|page=293}}</ref>

===Isotopes===
{{main|Isotopes of sodium}}
Twenty isotopes of sodium are known, but only <sup>23</sup>Na is stable. Two [[radioactive decay|radioactive]], [[cosmogenic]] isotopes are the byproduct of [[cosmic ray spallation]]: <sup>22</sup>Na with a [[half-life]] of 2.6 years and <sup>24</sup>Na with a half-life of 15 hours; all other isotopes have a half-life of less than one minute.<ref>{{cite journal| last=Audi|first=Georges|title=The NUBASE Evaluation of Nuclear and Decay Properties| journal=Nuclear Physics A|volume=729|pages=3–128|date=2003| doi=10.1016/j.nuclphysa.2003.11.001| bibcode=2003NuPhA.729....3A| last2=Bersillon| first2=O.| last3=Blachot| first3=J.| last4=Wapstra| first4=A.H.}}</ref> Two [[nuclear isomer]]s have been discovered, the longer-lived one being <sup>24m</sup>Na with a half-life of around 20.2 microseconds. Acute neutron radiation, such as from a nuclear [[criticality accident]], converts some of the stable <sup>23</sup>Na in human blood to <sup>24</sup>Na; by measuring the concentration of <sup>24</sup>Na in relation to <sup>23</sup>Na, the neutron radiation dosage of the victim can be calculated.<ref>{{cite journal|title=Neutron Activation of Sodium in Anthropomorphous Phantoms|journal=HealthPhysics| volume=8|issue=4| pages=371–379| date=1962| last1=Sanders| first1=F. W.|last2=Auxier|first2=J. A.| doi= 10.1097/00004032-196208000-00005|pmid=14496815}}</ref>

===Occurrence===
<sup>23</sup>Na is created in the [[carbon-burning process]] in stars by fusing two [[carbon]] atoms together; this requires temperatures above 600 megakelvins and a star of at least three solar masses.<ref>{{cite journal|bibcode=1987SvAL...13..214D|title= Sodium Synthesis in Hydrogen Burning Stars|last1=Denisenkov |first=P. A.|last2=Ivanov|first2=V. V.|volume=13| date=1987|page= 214|journal=Soviet Astronomy Letters}}</ref> The Earth's crust contains 2.6% sodium by weight, making it the [[Abundance of the chemical elements|sixth most abundant element]] on Earth.<ref>{{RubberBible86th}}</ref> Sodium's estimated [[Abundance of elements in Earth's crust|crustal abundance]] is 2.36×10<sup>4</sup> milligrams per kilogram.{{citation needed|date=December 2015}} Sodium's estimated [[oceanic abundance]] is 1.08×10<sup>4</sup> milligrams per liter.{{citation needed|date=December 2015}} Because of its high reactivity, it is never found as a pure element. It is found in many different minerals, some very soluble, such as [[halite]] and [[natron]], others much less soluble, such as [[amphibole]] and [[zeolite]]. The insolubility of certain sodium minerals such as [[cryolite]] and [[feldspar]] arises from their polymeric anions, which in the case of feldspar is a polysilicate. In the [[interstellar medium]], sodium is identified by the D [[Fraunhofer lines|spectral line]]; though it has a high vaporization temperature, its abundance allowed it to be detected by [[Mariner 10]] in [[Mercury (planet)|Mercury's]] atmosphere.<ref>{{cite web|url=http://findarticles.com/p/articles/mi_m1200/is_v128/ai_3898126|date= 17 August 1985|first=Dietrick E.|last= Tjrhonsen|accessdate=2008-09-18|title=Sodium found in Mercury's atmosphere|publisher=BNET}}{{dead link|date=January 2016}}</ref> Sodium has been detected in at least one [[comet]]; astronomers watching [[Comet Hale-Bopp]] in 1997 observed a sodium [[Comet tail|tail]], which consisted of neutral atoms (not ions) and extended to some 50 million kilometres in length.<ref name="Cremonese1997">{{cite journal | journal = The Astrophysical Journal Letters | volume = 490 | issue = 2 | pages =L199–L202 | year = 1997 | doi = 10.1086/311040 | title = Neutral Sodium from Comet Hale–Bopp: A Third Type of Tail | author1 = Cremonese, G | author2 = Boehnhardt, H | author3 = Crovisier, J | author4 = Rauer, H | author5 = Fitzsimmons, A | author6 = Fulle, M | author7 = Licandro, J | author8 = Pollacco, D | author9 = Tozzi, G. P | bibcode=1997ApJ...490L.199C|arxiv = astro-ph/9710022 | display-authors = 8 | last10 = West | first10 = R. M.}}</ref>

==Compounds==
{{Category see also|Sodium compounds}}
[[File:NaCl polyhedra.png|thumb|left|220 px|Structure of [[sodium chloride]], showing octahedral coordination around Na<sup>+</sup> and Cl<sup>−</sup> centres. This framework disintegrates upon dissolution in water and reassembles upon evaporation.]]
Sodium compounds are of immense commercial importance, being particularly central to industries producing [[glass]], [[paper]], [[soap]], and [[textile]]s.<ref name=Ullmann/> The sodium compounds that are the most important include [[table salt]] (Na[[chloride|Cl]]), [[soda ash]] (Na<sub>2</sub>[[carbonate|CO<sub>3</sub>]]), [[baking soda]] (Na[[Bicarbonate|HCO<sub>3</sub>]]), [[sodium hydroxide|caustic soda]] (NaOH), [[sodium nitrate]] (Na[[nitrate|NO<sub>3</sub>]]), di- and tri-sodium phosphates, [[sodium thiosulfate]] (Na<sub>2</sub>[[thiosulfate|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).<ref name="Holl"/> In its compounds, sodium is usually [[ionic bond|ionically bonded]] to water and anions, and is viewed as a [[HSAB|hard]] [[Lewis acid]].<ref>{{cite book|last=Cowan|first=James A.|title=Inorganic Biochemistry: An Introduction|date=1997|publisher=Wiley-VCH|isbn=978-0-471-18895-7|page=7|oclc=34515430}}</ref>

[[File:StericAcidChemStr.png|thumb|right|300px|Two equivalent images of the chemical structure of [[sodium stearate]], a typical soap.]]
Most [[soap]]s are sodium salts of [[fatty acid]]s. Sodium soaps are harder (higher melting) soaps than potassium soaps.<ref name="Holl">{{cite book|publisher=Walter de Gruyter|date=1985|edition=91–100|pages=931–943|isbn=3-11-007511-3|title=Lehrbuch der Anorganischen Chemie|last1=Holleman|first1=Arnold F.|last2=Wiberg|first2=Egon|last3=Wiberg|first3=Nils|language=German}}</ref> [[Sodium chloride]] is extensively used for [[anti-icing]] and [[de-icing]] and as a preservative; [[sodium bicarbonate]] is mainly used for cooking. Along with potassium, many important medicines have sodium added to improve their [[bioavailability]]; although in most cases potassium is the better ion, sodium is selected for its lower price and atomic weight.<ref>{{cite book|last=Remington|first=Joseph P.|title=Remington: The Science and Practice of Pharmacy|date=2006|publisher=Lippincott Williams & Wilkins|isbn=978-0-7817-4673-1|pages=365–366|edition=21st|editor-last=Beringer|editor-first=Paul|oclc=60679584}}</ref> [[Sodium hydride]] is used as a base for various reactions (such as the [[aldol reaction]]) in organic chemistry, and as a reducing agent in inorganic chemistry.<ref>{{cite book|last1=Wiberg|first1=Egon|last2=Wiberg|first2=Nils|last3=Holleman|first3=A. F.|title=Inorganic Chemistry|date=2001|publisher=Academic Press|pages=1103–1104|url=https://books.google.com/books?id=Mtth5g59dEIC&pg=PA1103|isbn=978-0-12-352651-9|oclc=48056955}}</ref>

===Aqueous solutions===
Sodium tends to form water-soluble compounds, such as [[halides]], [[sulfate]]s, [[nitrates]], [[carboxylates]] and [[carbonate]]s. The main aqueous species are the aquo complexes [Na(H<sub>2</sub>O)<sub>''n''</sub>]<sup>+</sup>, where ''n'' = 4–6.<ref name=Lincoln>{{cite book|doi=10.1016/B0-08-043748-6/01055-0|title=Comprehensive Coordination Chemistry II|date=2004|isbn=978-0-08-043748-4|page=515|chapter=Metal Aqua Ions|last1=Lincoln|first1=S.F.|last2=Richens|first2=D.T.|last3=Sykes|first3=A.G.}}</ref> The high affinity of sodium for [[oxygen]]-based [[ligand]]s is the basis of [[crown ether]]s; [[macrolide]] antibiotics, which interfere with Na<sup>+</sup> transport in the infecting organism, are functionally related and more complex.{{citation needed|date=December 2013}}

Direct precipitation of sodium salts from aqueous solutions is rare because sodium salts typically have a high affinity for water; an exception is [[sodium bismuthate]] (NaBiO<sub>3</sub>).<ref>{{cite book|title=Lange's Handbook of Chemistry|publisher=McGraw-Hill|date=1998|isbn=0-07-016384-7|last1=Dean|first1=John Aurie|last2=Lange|first2=Norbert Adolph}}</ref> Because of this, sodium salts are usually isolated as solids by evaporation or by precipitation with an organic solvent, such as [[ethanol]]; for example, only 0.35 g/L of sodium chloride will dissolve in ethanol.<ref>{{cite book|last=Burgess|first= J.|title=Metal Ions in Solution|publisher=Ellis Horwood|location=New York|date=1978|isbn=0-85312-027-7}}</ref> [[Crown ether]]s, like [[15-crown-5]], may be used as a [[phase-transfer catalyst]].<ref>{{cite book|last1=Starks|first1=Charles M.|last2=Liotta|first2=Charles L.|last3=Halpern|first3=Marc|title=Phase-Transfer Catalysis: Fundamentals, Applications, and Industrial Perspectives|date=1994|publisher=Chapman & Hall|page=162|isbn=978-0-412-04071-9|oclc=28027599}}</ref>

Sodium content in bulk may be determined by treating with a large excess of [[uranyl zinc acetate]]; the hexahydrate (UO<sub>2</sub>)<sub>2</sub>ZnNa(CH<sub>3</sub>CO<sub>2</sub>)·6H<sub>2</sub>O precipitates and [[gravimetry|can be weighed]]. Caesium and rubidium do not interfere with this reaction, but potassium and lithium do.<ref>{{cite journal|journal=J. Am. Chem. Soc.|doi=10.1021/ja01386a008|date=1929|last1=Barber|first1=H. H.|last2=Kolthoff|first2=I. M.|volume=51|issue=11|page=3233|title=Gravimetric Determination of Sodium by the Uranyl Zinc Acetate Method. Ii. Application in the Presence of Rubidium, Cesium, Potassium, Lithium, Phosphate or Arsenate}}</ref> Lower concentrations of sodium may be determined by [[atomic absorption spectrophotometry]]<ref>{{cite journal|url=http://www.jbc.org/content/206/2/807|journal=J. Biol. Chem.|volume=206|issue=2|pages=807–15|date=1954|pmid=13143043|last1=Kingsley|first1=G. R.|last2=Schaffert|first2=R. R.|title=Micro-flame Photometric Determination of Sodium, Potassium and Calcium in Serum with Solvents}}</ref> or by [[potentiometry]] using ion-selective electrodes.<ref>{{cite journal|last=Levy|first=G. B.|title=Determination of Sodium with Ion-Selective Electrodes|journal=Clinical Chemistry|url=http://www.clinchem.org/content/27/8/1435|volume=27|issue=8|pages=1435–1438|date=1981|pmid=7273405}}</ref>

===Electrides and sodides===
Like the other alkali metals, sodium dissolves in ammonia and some amines to give deeply colored solutions; evaporation of these solutions leaves a shiny film of metallic sodium. The solutions contain the [[coordination complex]] (Na(NH<sub>3</sub>)<sub>6</sub>)<sup>+</sup>, whose positive charge is counterbalanced by [[electride|electrons as anions]]; [[cryptand]]s permit the isolation of these complexes as crystalline solids. Cryptands, like crown ethers and other [[ionophore]]s, have a high affinity for the sodium ion; derivatives of the [[alkalide]] Na<sup>−</sup> are obtainable<ref>{{cite journal|journal=[[J. Am. Chem. Soc.]]|last1=Dye|first1=J. L.|last2=Ceraso|first2=J. M.|author3=Mei Lok Tak|last4=Barnett|first4=B. L.|last5=Tehan|first5=F. J.|title=Crystalline Salt of the Sodium Anion (Na<sup>−</sup>)|date=1974|volume=96|issue=2|pages=608–609|doi=10.1021/ja00809a060}}</ref> by the addition of cryptands to solutions of sodium in ammonia via [[disproportionation]].<ref>{{cite book|last1=Holleman|first1=A. F.|last2=Wiberg|first2=E.|last3=Wiberg|first3=N.|title=Inorganic Chemistry|publisher=Academic Press|date=2001|isbn=978-0-12-352651-9|oclc=48056955}}</ref>

===Organosodium compounds===
[[File:Monensin2.png|thumb|The structure of the complex of sodium (Na<sup>+</sup>, shown in yellow) and the antibiotic [[monensin]]-A.]]
Many organosodium compounds have been prepared. Because of the high polarity of the C-Na bonds, they behave like sources of [[carbanion]]s (salts with organic [[anion]]s). Some well known derivatives include [[sodium cyclopentadienide]] (NaC<sub>5</sub>H<sub>5</sub>) and trityl sodium ((C<sub>6</sub>H<sub>5</sub>)<sub>3</sub>CNa).<ref>{{OrgSynth |first1=W. B. | last1=Renfrow, Jr. |first2= C. R. | last2= Hauser |year=1943 |title=Triphenylmethylsodium |volume= |pages= |collvol=2 |collvolpages=607 |prep=CV2P0607}}</ref>

==History==
Salt has been an important commodity in human activities, as shown by the English word ''salary'', which derives from ''salarium'', the wafers of salt sometimes given to Roman soldiers along with their other wages. In medieval Europe, a compound of sodium with the Latin name of ''sodanum'' was used as a [[headache]] remedy. The name sodium is thought to originate from the Arabic ''suda'' , meaning headache, as the headache-alleviating properties of sodium carbonate or soda were well known in early times.<ref name=newton>{{cite book|last=Newton|first=David E.|editor-last=Baker|editor-first=Lawrence W.|title=Chemical Elements|date=1999|isbn=978-0-7876-2847-5|oclc=39778687}}</ref> The chemical abbreviation for sodium was first published by [[Jöns Jakob Berzelius]] in his system of atomic symbols,<ref>{{cite web|url=http://elements.vanderkrogt.net/element.php?sym=Na|title=Elementymology &amp; Elements Multidict|first = Peter|last = 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 name=newton/> a natural mineral salt primarily made of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds. Although sodium, sometimes called ''soda'', had long been recognised in compounds, the metal itself was not isolated until 1807 by [[Humphry Davy|Sir Humphry Davy]] through the [[electrolysis]] of [[sodium hydroxide]].<ref name=Davy1807>{{cite journal|first=Humphry|last=Davy|title=On some new phenomena of chemical changes produced by electricity, particularly the decomposition of the fixed alkalies, and the exhibition of the new substances which constitute their bases; and on the general nature of alkaline bodies|date=1808|volume=98|journal=Philosophical Transactions of the Royal Society of London|pages=1–44|url=https://books.google.com/?id=gpwEAAAAYAAJ&pg=PA57|doi=10.1098/rstl.1808.0001}}</ref><ref name="weeks">{{cite journal|doi=10.1021/ed009p1035|title=The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium|date=1932|last1=Weeks|first1=Mary Elvira|authorlink1=Mary Elvira Weeks|journal=Journal of Chemical Education|volume=9|issue=6|page=1035|bibcode=1932JChEd...9.1035W}}</ref>

Sodium imparts an intense yellow color to flames. As early as 1860, [[Gustav Kirchhoff|Kirchhoff]] and [[Robert Bunsen|Bunsen]] noted the high sensitivity of a sodium flame test, and stated in [[Annalen der Physik|Annalen der Physik und Chemie]]:<ref name="bunsen1">{{cite journal|doi =10.1002/andp.18601860602| title=Chemische Analyse durch Spectralbeobachtungen|date=1860 |last1=Kirchhoff |first1=G.|last2=Bunsen|first2=R. |journal=Annalen der Physik und Chemie |volume=186|issue=6 |pages=161–189|bibcode=1860AnP...186..161K}}</ref>

<blockquote>In a corner of our 60&nbsp;m<sup>3</sup> 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.</blockquote>

==Commercial production==
Enjoying rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.<ref name=Ullmann/> Metallic sodium was first produced commercially in 1855 by [[carbothermal reduction]] of [[sodium carbonate]] at 1100&nbsp;°C,{{citation needed|date=August 2012}} in what is known as the [[Deville process]]:<ref name=kirk>{{cite book|last1=Eggeman |first1=Tim|title=Kirk-Othmer Encyclopedia of Chemical Technology|publisher=John Wiley & Sons|date=2007 |doi=10.1002/0471238961.1915040912051311.a01.pub3|last2=Updated By Staff|isbn=0-471-23896-1|chapter=Sodium and Sodium Alloys}}</ref><ref>{{cite journal|jstor=27757153|pages=205–221 |last1=Oesper|first1=R. E.|last2=Lemay| first2=P.| title=Henri Sainte-Claire Deville, 1818–1881| volume=3| journal= Chymia|date=1950|doi=10.2307/27757153}}</ref><ref>{{cite journal|doi =10.1021/ed067p1046|title=Sodium|date=1990| last=Banks |first=Alton|journal=Journal of Chemical Education|volume=67 |issue=12|page=1046 |bibcode=1990JChEd..67.1046B}}</ref><!-- Eggeman's source is actually part of the Kirk-Othmer Encyclopedia of Chemical Technology, but the Wiley Online Library text for that has restricted access. -->
:Na<sub>2</sub>CO<sub>3</sub> + 2 C → 2 Na + 3 CO
A related process based on the reduction of sodium hydroxide was developed in 1886.<ref name=kirk/>

Sodium is now produced commercially through the [[electrolysis]] of molten [[sodium chloride]], based on a process patented in 1924.<ref name="pauling">Pauling, Linus, ''General Chemistry'', 1970 ed., Dover Publications</ref><ref name="losal">{{cite web|url=http://periodic.lanl.gov/11.shtml|title=Los Alamos National Laboratory – Sodium|accessdate=2007-06-08}}</ref> This is done in a [[Downs Cell|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 deposited at the cathode. This method is less expensive than the previous [[Castner process]] of electrolyzing [[sodium hydroxide]].

Reagent-grade sodium in tonne quantities sold for about US$3.30/kg in 2009; lower purity metal sells for considerably less.{{citation needed|date=August 2012}} The market for sodium 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 compound|organic]] materials. Smaller quantities of sodium cost far more, in the range of US$165/kg; the high cost is partially due to the expense of shipping [[hazardous material]].<ref>{{cite web| url=http://www.mcssl.com/store/gallium-source/sodium-metal |title=007-Sodium Metal| publisher=Mcssl.com |accessdate=2010-11-27}}</ref>

==Applications==
Though metallic sodium has some important uses, the major applications of sodium use is in its many compounds; millions of tons of the [[sodium chloride|chloride]], [[sodium hydroxide|hydroxide]], and [[sodium carbonate|carbonate]] are produced annually.

===Free element===
Metallic sodium is mainly used for the production of [[sodium borohydride]], [[sodium azide]], [[indigo dye|indigo]], and [[triphenylphosphine]]. Previous uses were for the making of [[tetraethyllead]] and titanium metal; because applications for these chemicals were discontinued, the production of sodium declined after 1970.<ref name=Ullmann>Alfred Klemm, Gabriele Hartmann, Ludwig Lange, "Sodium and Sodium Alloys" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. {{DOI|10.1002/14356007.a24_277}}</ref> Sodium is also used as an alloying metal, an [[anti-scaling agent]],<ref>{{cite book|url=https://books.google.com/books?id=LI4KmKqca78C&pg=PA76|page=76|title=Metal cleaning: bibliographical abstracts, 1842–1951|last=Harris|first=Jay C.|date=1949|publisher=[[American Society for Testing and Materials]]|oclc=1848092}}</ref> and as a reducing agent for metals when other materials are ineffective. Note the free element is not used as a scaling agent, ions in the water are exchanged for sodium ions. [[Sodium vapor lamp]]s are often used for street lighting in cities and give colours ranging from yellow-orange to peach as the pressure increases.<ref>{{cite book|url=https://books.google.com/books?id=0d7u9Nr33zIC&pg=PA112|pages=112–114|title=Applied illumination engineering|last=Lindsey|first=Jack L.|date=1997|publisher=Fairmont Press|isbn=978-0-88173-212-2|oclc=22184876}}</ref> By itself or [[NaK|with potassium]], sodium is a [[desiccant]]; it gives an intense blue colouration with [[benzophenone]] when the desiccate is dry.<ref>{{cite book|last=Lerner|first=Leonid|title=Small-Scale Synthesis of Laboratory Reagents with Reaction Modeling|url=https://books.google.com/books?id=VqosZeMjNjEC&pg=PA91|pages=91–92|date=2011-02-16|publisher=CRC Press|isbn=978-1-4398-1312-6|oclc=669160695}}</ref> In [[organic synthesis]], sodium is used in various reactions such as the [[Birch reduction]], and the [[sodium fusion test]] is conducted to qualitatively analyse compounds.<ref>{{cite book|last=Sethi|first=Arun|title=Systematic Laboratory Experiments in Organic Chemistry|url=https://books.google.com/books?id=x77djyQHX8UC&pg=PA32|pages=32–35|publisher=New Age International|isbn=978-81-224-1491-2|oclc=86068991|date=1 January 2006}}</ref> Lasers emitting light at the D line, utilising sodium, are used to create artificial [[laser guide star]]s that [[FASOR (laser physics)|assist]] in the [[adaptive optics]] for land-based visible light telescopes.{{citation needed|date=April 2012}}

====Heat transfer====
[[File:Phase diagram potassium sodium s l.svg|400px|thumb|NaK [[phase diagram]], showing the melting point of sodium as a function of potassium concentration. NaK with 77% potassium is [[eutectic system|eutectic]] and has the lowest melting point of the NaK alloys at −12.6 °C.<ref>{{cite journal|doi=10.1002/zaac.19120740115|last1=van Rossen|first1=G. L. C. M.|last2=van Bleiswijk|first2=H.|title=Über das Zustandsdiagramm der Kalium-Natriumlegierungen|journal=Zeitschrift für anorganische Chemie|volume=74|date=1912|pages=152–156}}</ref>]]
Liquid sodium is used as a [[Coolant|heat transfer fluid]] in [[sodium-cooled fast reactor|some fast reactors]],<ref name="sodiumcoolant">[http://www.ne.doe.gov/pdfFiles/SodiumCoolant_NRCpresentation.pdf Sodium as a Fast Reactor Coolant] presented by Thomas H. Fanning. Nuclear Engineering Division. U.S. Department of Energy. U.S. Nuclear Regulatory Commission. Topical Seminar Series on Sodium Fast Reactors. May 3, 2007</ref> due to its high thermal conductivity and low neutron absorption [[cross section (physics)|cross section]], which is required to achieve a high neutron flux; the high boiling point allows the reactor to operate at ambient pressure. Drawbacks of using sodium include its opacity, which hinders visual maintenance, and its explosive properties. Radioactive [[sodium-24]] may be formed by [[neutron activation]] during operation, posing a slight radiation hazard; the radioactivity stops within a few days after removal from the reactor. If a reactor needs to be frequently shut down, [[NaK]] is used; due to it being liquid at room temperature, cooling pipes do not freeze. In this case, the [[pyrophoricity]] of potassium means extra precautions against leaks need to be taken. Another heat transfer application is in high-performance internal combustion engines with [[poppet valves]], where valve stems partially filled with sodium are used as a [[heat pipe]] to cool the valves.

==Biological role==
{{main|Sodium in biology}}
In humans, sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and [[pH]]; the minimum physiological requirement for sodium is 500 milligrams per day.<ref name=r31>{{cite web|url=http://nuinfo-proto4.northwestern.edu/nutrition/factsheets/sodium.pdf|archiveurl=https://web.archive.org/web/20110823114818/http://nuinfo-proto4.northwestern.edu/nutrition/factsheets/sodium.pdf|archivedate=2011-08-23|title=Sodium|publisher=Northwestern University|accessdate=2011-11-21}}</ref> [[Sodium chloride]] is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for [[pickling]] and [[Jerky (food)|jerky]]; most of it comes from processed foods.<ref>{{cite web|url=http://health.ltgovernors.com/sodium-and-potassium-health-facts.html|title=Sodium and Potassium Quick Health Facts|work=health.ltgovernors.com}}</ref> The [[Dietary Reference Intake|UL]] for sodium is 2.3&nbsp;grams per day,<ref>{{cite web|title=Dietary Reference Intakes: Water, Potassium, Sodium, Chloride, and Sulfate|url=http://www.iom.edu/Reports/2004/Dietary-Reference-Intakes-Water-Potassium-Sodium-Chloride-and-Sulfate.aspx|publisher=Food and Nutrition Board, [[Institute of Medicine]], [[United States National Academies]]|date=11 February 2004}}</ref> the threshold which could lead to [[hypertension]] when exceeded,<ref>{{cite journal|pmid=15369026|date=2004|last1=Geleijnse|first1=J. M.|last2=Kok|first2=F. J.|last3=Grobbee|first3=D. E.|title=Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations|volume=14|issue=3|pages=235–239|journal=European Journal of Public Health|doi=10.1093/eurpub/14.3.235}}</ref> but on average people in the United States consume 3.4&nbsp;grams per day.<ref>{{cite book|author1=U.S. Department of Agriculture|author2=U.S. Department of Health and Human Services|authorlink1=United States Department of Agriculture|authorlink2=United States Department of Health and Human Services|title=Dietary Guidelines for Americans, 2010|page=22|edition=7th|date=December 2010|archiveurl=http://wayback.archive.org/web/20110206111821/http://www.cnpp.usda.gov/Publications/DietaryGuidelines/2010/PolicyDoc/PolicyDoc.pdf|url=http://www.cnpp.usda.gov/Publications/DietaryGuidelines/2010/PolicyDoc/PolicyDoc.pdf|archivedate=6 February 2011|format=PDF|accessdate=2011-11-23|isbn=978-0-16-087941-8|oclc=738512922}}</ref> Hypertension causes 7.6 million premature deaths worldwide each year.<ref>{{cite journal|pmid=18456100|date=2008|last1=Lawes|first1=C. M.|last2=Vander Hoorn|first2=S.|last3=Rodgers|first3=A.|author4=International Society of Hypertension|title=Global burden of blood-pressure-related disease, 2001|volume=371|issue=9623|pages=1513–1518|doi=10.1016/S0140-6736(08)60655-8|journal=Lancet}}</ref> (Note that salt contains about 39.3% sodium<ref>{{cite book|author=Armstrong, James |title=General, Organic, and Biochemistry: An Applied Approach |url=https://books.google.com/books?id=bcU8AAAAQBAJ&pg=PA48 |date=2011 |publisher=Cengage Learning |isbn=1-133-16826-4 |pages=48–}}</ref>{{--}}the rest being chlorine and other trace chemicals; thus the UL of 2.3g sodium would be about 5.9g, or 2.7ml of salt{{--}}about half a [[Teaspoon (metric)|US teaspoon]]<ref>[http://www.traditionaloven.com/culinary-arts/cooking/table-salt/convert-gram-g-to-tea-spoon-tsp.html Table Salt Conversion]. Traditionaloven.com. Retrieved on 2015-11-11.</ref>)

The [[renin-angiotensin system]] regulates the amount of fluids and sodium in the body. Reduction of blood pressure and sodium concentration in the kidney result in the production of [[renin]], which in turn produces [[aldosterone]] and [[angiotensin]], retaining sodium in the urine. Because of the increase in sodium concentration, the production of renin decreases, and the sodium concentration returns to normal.<ref>{{cite book|last1=McGuire|first1=Michelle|last2=Beerman|first2=Kathy A.|title=Nutritional Sciences: From Fundamentals to Food|date=2011|publisher=Cengage Learning|isbn=978-0-324-59864-3|page=546|oclc=472704484}}</ref> Sodium is also important in [[neuron]] function and osmoregulation between cells and the [[extracellular fluid]], their distribution mediated in all animals by [[Na+/K+-ATPase|Na<sup>+</sup>/K<sup>+</sup>-ATPase]], an active transporter pumping ions against the gradient, and sodium/potassium channels.<ref>{{cite book|last=Campbell|first=Neil|title=Biology|date=1987|isbn=0-8053-1840-2|page=795|publisher=Benjamin/Cummings}}</ref> Sodium is the most prominent metal ion in extracellular fluid.<ref>{{cite book|last=Srilakshmi|first=B.|title=Nutrition Science|date=2006|publisher=New Age International|url=https://books.google.com/books?id=f_i7j4_cMLIC&pg=PA318|isbn=978-81-224-1633-6|edition=2nd|page=318|oclc=173807260}}</ref>

Unusually low or high sodium levels in humans are recognized in medicine as [[hyponatremia]] and [[hypernatremia]]. These conditions may be caused by genetic factors, physical factors associated with ageing or illnesses involving vomiting or diarrhea.<ref>
{{cite book
|last1=Pohl
|first1=Hanna R.
|last2=Wheeler
|first2=John S.
|first3=H. Edward
|last3=Murray
|editor=Astrid Sigel
|editor2=Helmut Sigel
|editor3=Roland K. O. Sigel
|title=Interrelations between Essential Metal Ions and Human Diseases
|series=Metal Ions in Life Sciences
|volume=13
|date=2013
|publisher=Springer
|pages=29–47
|doi=10.1007/978-94-007-7500-8_2
}}
</ref>

In [[C4 plants]], sodium is a [[micronutrient]] that aids in metabolism, specifically in regeneration of [[phosphoenolpyruvic acid|phosphoenolpyruvate]] and synthesis of chlorophyll.<ref>{{cite web|last=Kering|first=M. K.|url=https://mospace.umsystem.edu/xmlui/bitstream/handle/10355/7201/research.pdf?sequence=3|title=Manganese Nutrition and Photosynthesis in NAD-malic enzyme C4 plants Ph.D. dissertation|publisher=University of Missouri-Columbia|date=2008|accessdate=2011-11-09}}</ref> In others, it substitutes for [[potassium]] in several roles, such as maintaining [[turgor pressure]] and aiding in the opening and closing of stomata.<ref>{{cite journal|last1=Subbarao|first1=G. V.|last2=Ito|first2=O.|last3=Berry|first3=W. L.|last4=Wheeler|first4=R. M.|title=Sodium—A Functional Plant Nutrient|journal=Critical Reviews in Plant Sciences| volume=22| pages=391–416| date=2003|doi=10.1080/07352680390243495|issue=5}}</ref> Excess sodium in the soil limits the uptake of water due to decreased [[water potential]], which may result in wilting; similar concentrations in the [[cytoplasm]] can lead to enzyme inhibition, which in turn causes necrosis and chlorosis.<ref>{{cite journal| last1=Zhu|first1=J. K.|title=Plant salt tolerance|journal=Trends in Plant Science|volume=6|issue=2|pages=66–71|date=2001|pmid=11173290|doi=10.1016/S1360-1385(00)01838-0}}</ref> To avoid these problems, plants developed mechanisms that limit sodium uptake by roots, store them in cell [[vacuole]]s, and control them over long distances;<ref>{{cite web|url=http://www.plant-biology.com/salt-ion-toxicity.php|title=Plants and salt ion toxicity|publisher=Plant Biology|accessdate=2010-11-02}}</ref> excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

==Precautions==
{{NFPA 704|Health = 3|Flammability = 1|Reactivity = 2|S= W|caption=The fire diamond hazard sign for sodium metal<ref>[http://www.ehs.neu.edu/laboratory_safety/general_information/nfpa_hazard_rating/documents/NFPAratingSZ.htm Hazard Rating Information for NFPA Fire Diamonds]. Ehs.neu.edu. Retrieved on 2015-11-11.</ref>}}
Care is required in handling elemental sodium, as it generates flammable hydrogen and caustic [[sodium hydroxide]] upon contact with water; powdered sodium may spontaneously explode in the presence of an oxidizer.{{citation needed|date=November 2012}} Excess sodium can be safely removed by [[hydrolysis]] in a ventilated cabinet; this is typically done by sequential treatment with isopropanol, ethanol and water. Isopropanol reacts very slowly, generating the corresponding [[alkoxide]] and hydrogen.<ref>{{cite book | author= Angelici, R. J.|title= Synthesis and Technique in Inorganic Chemistry |publisher = University Science Books |place = Mill Valley, CA | date = 1999 | isbn = 0-935702-48-2}}</ref> [[Fire extinguisher]]s based on water accelerate sodium fires; those based on carbon dioxide and [[bromochlorodifluoromethane]] lose their effectiveness when they dissipate.{{citation needed|date=November 2015}} An effective extinguishing agent is [[Met-L-X]], which comprises approximately 5% [[Saran (plastic)|Saran]] in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop.{{citation needed|date=November 2015}} Other materials include [[Lith+]], which has [[graphite]] powder and an [[organophosphate]] [[flame retardant]], and dry sand.{{citation needed|date=November 2015}}

==See also==
[[Sodium in biology]]

[[Alkali metal]]{{Subject bar
|portal=Chemistry
|book1=Sodium
|book2=Period 3 elements
|book3=Alkali metals
|book4=Chemical elements (sorted&nbsp;alphabetically)
|book5=Chemical elements (sorted by number)
|commons=y
|wikt=y
|wikt-search=sodium
|v=y
|v-search=Sodium atom
|b=y
|b-search=Wikijunior:The Elements/Sodium
}}

==References==
{{reflist|30em}}

==External links==
* [http://www.periodicvideos.com/videos/011.htm Sodium] at ''[[The Periodic Table of Videos]]'' (University of Nottingham)
* [http://www.balashon.com/2008/07/neter-and-nitrogen.html Etymology of "natrium" – source of symbol Na]
* [http://www.theodoregray.com/PeriodicTable/Elements/011/index.html The Wooden Periodic Table Table's Entry on Sodium]
* [http://ie.lbl.gov/education/parent/Na_iso.htm Sodium isotopes data from ''The Berkeley Laboratory Isotopes Project's'']

{{Clear}}
{{Compact periodic table}}
{{Sodium compounds}}
{{alkalimetals}}
{{Use dmy dates|date=November 2010}}

{{Authority control}}

[[Category:Sodium minerals|minerals]]
[[Category:Sodium| ]]
[[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]]

Revision as of 17:04, 6 February 2016

This page is about the chemical element. For the nutrient commonly called "sodium", see salt. For the PlayStation Home game, see Sodium (PlayStation Home). For the racehorse, see Sodium (horse).
"Natrium" redirects here. For the town in West Virginia, see Natrium, West Virginia.
Sodium, 11Na
Sodium
Appearancesilvery white metallic
Standard atomic weight Ar°(Na)
Sodium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Li
↑
Na
↓
K
neon ← sodium → magnesium
Atomic number (Z)11
Groupgroup 1: hydrogen and alkali metals
Periodperiod 3
Block  s-block
Electron configuration[Ne] 3s1
Electrons per shell2, 8, 1
Physical properties
Phase at STPsolid
Melting point370.944 K ​(97.794 °C, ​208.029 °F)
Boiling point1156.090 K ​(882.940 °C, ​1621.292 °F)
Density (at 20° C)0.9688 g/cm3[3]
when liquid (at m.p.)0.927 g/cm3
Critical point2573 K, 35 MPa (extrapolated)
Heat of fusion2.60 kJ/mol
Heat of vaporization97.42 kJ/mol
Molar heat capacity28.230 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 554 617 697 802 946 1153
Atomic properties
Oxidation statescommon: +1
−1,[4] 0[5]
ElectronegativityPauling scale: 0.93
Ionization energies
  • 1st: 495.8 kJ/mol
  • 2nd: 4562 kJ/mol
  • 3rd: 6910.3 kJ/mol
  • (more)
Atomic radiusempirical: 186 pm
Covalent radius166±9 pm
Van der Waals radius227 pm
Color lines in a spectral range
Spectral lines of sodium
Other properties
Natural occurrenceprimordial
Crystal structure ​body-centered cubic (bcc) (cI2)
Lattice constant
Body-centered cubic crystal structure for sodium
a = 428.74 pm (at 20 °C)[3]
Thermal expansion69.91×10−6/K (at 20 °C)[3]
Thermal conductivity142 W/(m⋅K)
Electrical resistivity47.7 nΩ⋅m (at 20 °C)
Magnetic orderingparamagnetic[6]
Molar magnetic susceptibility+16.0×10−6 cm3/mol (298 K)[7]
Young's modulus10 GPa
Shear modulus3.3 GPa
Bulk modulus6.3 GPa
Speed of sound thin rod3200 m/s (at 20 °C)
Mohs hardness0.5
Brinell hardness0.69 MPa
CAS Number7440-23-5
History
Discovery and first isolationHumphry Davy (1807)
Symbol"Na": from New Latin natrium, coined from German Natron, 'natron'
Isotopes of sodium
Main isotopes[8] Decay
abun­dance half-life (t1/2) mode pro­duct
22Na trace 2.6019 y β+ 22Ne
23Na 100% stable
24Na trace 14.9560 h β− 24Mg
 Category: Sodium
| references

Sodium /ˈsoʊdiəm/[9] is a chemical element with symbol Na (from Ancient Greek Νάτριο) and atomic number 11. It is a soft, silver-white, highly reactive metal. In the Periodic table it is in column 1 (alkali metals), and shares with the other six elements in that column that it has a single electron in its outer shell, which it readily donates, creating a positively charged atom - a cation. Its only stable isotope is 23Na. The free metal does not occur in nature, but instead must be prepared from its compounds. Sodium is the sixth most abundant element in the Earth's crust, and exists in numerous minerals such as feldspars, sodalite and rock salt (NaCl). Many salts of sodium are highly water-soluble: sodium ions have been leached by the action of water from the earth's minerals over eons, so that sodium (and chlorine) are the most common dissolved elements by weight in the oceans.

Sodium was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide. Many sodium compounds are useful, such as sodium hydroxide (lye) for soap-making and sodium chloride for use as a de-icing agent and a nutrient (edible salt).

Sodium is an essential element for all animals and some plants. Sodium ions are the major cation in the extracellular fluid (ECF) and as such are the major contributor to the ECF osmotic pressure, and thus ECF compartment volume. Loss of only water from the ECF compartment increases the sodium concentration: hypernatremia. Isotonic loss of water and sodium from the ECF compartment decreases the size of that compartment: ECF hypovolemia.

Sodium ions are pumped out of cells by Na+/K+-ATPase in exchange for potassium. It pumps three sodium ions out of the cell for every two potassium ions pumped in, contributing to the fact that about forty times as much potassium is inside cells compared to outside, and about ten times as much sodium is outside cells compared to inside. In nerve cells, a buildup of electrical charge across the cell membrane allows transmission of a nerve impulse - an action potential - when the charge is dissipated.

Characteristics

Physical

Sodium at standard temperature and pressure is a soft silvery metal that oxidizes to grayish white unless immersed in oil or inert gas. Sodium can be readily cut with a knife and is a good conductor of electricity and heat. These properties change dramatically at elevated pressures: at 1.5 Mbar, the color changes from silvery metallic to black; at 1.9 Mbar the material becomes transparent with a red color; and at 3 Mbar sodium is a clear and transparent solid. All of these high-pressure allotropes are insulators and electrides.[10]

When sodium or its compounds are introduced into a flame, they turn it yellow,[11] because the excited 3s electrons of sodium emit a photon when they fall from 3p to 3s; the wavelength of this photon corresponds to the D line at 589.3 nm. Spin-orbit interactions involving the electron in the 3p orbital split the D line into two; hyperfine structures involving both orbitals cause many more lines.[12]

Chemical

Emission spectrum for sodium, showing the D line.
A positive flame test for sodium has a bright yellow color.

When freshly cut, sodium has a bright, silvery luster. If exposed to air, the surface rapidly tarnishes, darkening at first and then forming a white coating of sodium hydroxide and sodium carbonate.

Sodium is generally less reactive than potassium and more reactive than lithium.[13] Like all the alkali metals, it reacts exothermically with water, to the point that sufficiently large pieces melt to a sphere and may explode; this reaction produces caustic soda (sodium hydroxide) and flammable hydrogen gas. When burned in dry air, it mainly forms sodium peroxide as well as some sodium oxide. In moist air, sodium hydroxide results.[14] Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts.[15] Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy.[14] However, potassium and lithium have even more negative potentials.[16]

Isotopes

Main article: Isotopes of sodium

Twenty isotopes of sodium are known, but only 23Na is stable. Two radioactive, cosmogenic isotopes are the byproduct of cosmic ray spallation: 22Na with a half-life of 2.6 years and 24Na with a half-life of 15 hours; all other isotopes have a half-life of less than one minute.[17] Two nuclear isomers have been discovered, the longer-lived one being 24mNa with a half-life of around 20.2 microseconds. Acute neutron radiation, such as from a nuclear criticality accident, converts some of the stable 23Na in human blood to 24Na; by measuring the concentration of 24Na in relation to 23Na, the neutron radiation dosage of the victim can be calculated.[18]

Occurrence

23Na is created in the carbon-burning process in stars by fusing two carbon atoms together; this requires temperatures above 600 megakelvins and a star of at least three solar masses.[19] The Earth's crust contains 2.6% sodium by weight, making it the sixth most abundant element on Earth.[20] Sodium's estimated crustal abundance is 2.36×104 milligrams per kilogram.[citation needed] Sodium's estimated oceanic abundance is 1.08×104 milligrams per liter.[citation needed] Because of its high reactivity, it is never found as a pure element. It is found in many different minerals, some very soluble, such as halite and natron, others much less soluble, such as amphibole and zeolite. The insolubility of certain sodium minerals such as cryolite and feldspar arises from their polymeric anions, which in the case of feldspar is a polysilicate. In the interstellar medium, sodium is identified by the D spectral line; though it has a high vaporization temperature, its abundance allowed it to be detected by Mariner 10 in Mercury's atmosphere.[21] Sodium has been detected in at least one comet; astronomers watching Comet Hale-Bopp in 1997 observed a sodium tail, which consisted of neutral atoms (not ions) and extended to some 50 million kilometres in length.[22]

Compounds

See also: Category:Sodium compounds
Structure of sodium chloride, showing octahedral coordination around Na+ and Cl− centres. This framework disintegrates upon dissolution in water and reassembles upon evaporation.

Sodium compounds are of immense commercial importance, being particularly central to industries producing glass, paper, soap, and textiles.[23] The sodium compounds that are the most important include table salt (NaCl), soda ash (Na2CO3), baking soda (NaHCO3), caustic soda (NaOH), sodium nitrate (NaNO3), di- and tri-sodium phosphates, sodium thiosulfate (Na2S2O3·5H2O), and borax (Na2B4O7·10H2O).[24] In its compounds, sodium is usually ionically bonded to water and anions, and is viewed as a hard Lewis acid.[25]

File:StericAcidChemStr.png
Two equivalent images of the chemical structure of sodium stearate, a typical soap.

Most soaps are sodium salts of fatty acids. Sodium soaps are harder (higher melting) soaps than potassium soaps.[24] Sodium chloride is extensively used for anti-icing and de-icing and as a preservative; sodium bicarbonate is mainly used for cooking. Along with potassium, many important medicines have sodium added to improve their bioavailability; although in most cases potassium is the better ion, sodium is selected for its lower price and atomic weight.[26] Sodium hydride is used as a base for various reactions (such as the aldol reaction) in organic chemistry, and as a reducing agent in inorganic chemistry.[27]

Aqueous solutions

Sodium tends to form water-soluble compounds, such as halides, sulfates, nitrates, carboxylates and carbonates. The main aqueous species are the aquo complexes [Na(H2O)n]+, where n = 4–6.[28] The high affinity of sodium for oxygen-based ligands is the basis of crown ethers; macrolide antibiotics, which interfere with Na+ transport in the infecting organism, are functionally related and more complex.[citation needed]

Direct precipitation of sodium salts from aqueous solutions is rare because sodium salts typically have a high affinity for water; an exception is sodium bismuthate (NaBiO3).[29] Because of this, sodium salts are usually isolated as solids by evaporation or by precipitation with an organic solvent, such as ethanol; for example, only 0.35 g/L of sodium chloride will dissolve in ethanol.[30] Crown ethers, like 15-crown-5, may be used as a phase-transfer catalyst.[31]

Sodium content in bulk may be determined by treating with a large excess of uranyl zinc acetate; the hexahydrate (UO2)2ZnNa(CH3CO2)·6H2O precipitates and can be weighed. Caesium and rubidium do not interfere with this reaction, but potassium and lithium do.[32] Lower concentrations of sodium may be determined by atomic absorption spectrophotometry[33] or by potentiometry using ion-selective electrodes.[34]

Electrides and sodides

Like the other alkali metals, sodium dissolves in ammonia and some amines to give deeply colored solutions; evaporation of these solutions leaves a shiny film of metallic sodium. The solutions contain the coordination complex (Na(NH3)6)+, whose positive charge is counterbalanced by electrons as anions; cryptands permit the isolation of these complexes as crystalline solids. Cryptands, like crown ethers and other ionophores, have a high affinity for the sodium ion; derivatives of the alkalide Na− are obtainable[35] by the addition of cryptands to solutions of sodium in ammonia via disproportionation.[36]

Organosodium compounds

The structure of the complex of sodium (Na+, shown in yellow) and the antibiotic monensin-A.

Many organosodium compounds have been prepared. Because of the high polarity of the C-Na bonds, they behave like sources of carbanions (salts with organic anions). Some well known derivatives include sodium cyclopentadienide (NaC5H5) and trityl sodium ((C6H5)3CNa).[37]

History

Salt has been an important commodity in human activities, as shown by the English word salary, which derives from salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages. In medieval Europe, a compound of sodium with the Latin name of sodanum was used as a headache remedy. The name sodium is thought to originate from the Arabic suda , meaning headache, as the headache-alleviating properties of sodium carbonate or soda were well known in early times.[38] The chemical abbreviation for sodium was first published by Jöns Jakob Berzelius in his system of atomic symbols,[39] and is a contraction of the element's New Latin name natrium, which refers to the Egyptian natron,[38] a natural mineral salt primarily made of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds. Although sodium, sometimes called soda, had long been recognised in compounds, the metal itself was not isolated until 1807 by Sir Humphry Davy through the electrolysis of sodium hydroxide.[40][41]

Sodium imparts an intense yellow color to flames. As early as 1860, Kirchhoff and Bunsen noted the high sensitivity of a sodium flame test, and stated in Annalen der Physik und Chemie:[42]

In a corner of our 60 m3 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 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.

Commercial production

Enjoying rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.[23] Metallic sodium was first produced commercially in 1855 by carbothermal reduction of sodium carbonate at 1100 °C,[citation needed] in what is known as the Deville process:[43][44][45]

Na2CO3 + 2 C → 2 Na + 3 CO

A related process based on the reduction of sodium hydroxide was developed in 1886.[43]

Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924.[46][47] 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 deposited at the cathode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide.

Reagent-grade sodium in tonne quantities sold for about US$3.30/kg in 2009; lower purity metal sells for considerably less.[citation needed] The market for sodium 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. Smaller quantities of sodium cost far more, in the range of US$165/kg; the high cost is partially due to the expense of shipping hazardous material.[48]

Applications

Though metallic sodium has some important uses, the major applications of sodium use is in its many compounds; millions of tons of the chloride, hydroxide, and carbonate are produced annually.

Free element

Metallic sodium is mainly used for the production of sodium borohydride, sodium azide, indigo, and triphenylphosphine. Previous uses were for the making of tetraethyllead and titanium metal; because applications for these chemicals were discontinued, the production of sodium declined after 1970.[23] Sodium is also used as an alloying metal, an anti-scaling agent,[49] and as a reducing agent for metals when other materials are ineffective. Note the free element is not used as a scaling agent, ions in the water are exchanged for sodium ions. Sodium vapor lamps are often used for street lighting in cities and give colours ranging from yellow-orange to peach as the pressure increases.[50] By itself or with potassium, sodium is a desiccant; it gives an intense blue colouration with benzophenone when the desiccate is dry.[51] In organic synthesis, sodium is used in various reactions such as the Birch reduction, and the sodium fusion test is conducted to qualitatively analyse compounds.[52] Lasers emitting light at the D line, utilising sodium, are used to create artificial laser guide stars that assist in the adaptive optics for land-based visible light telescopes.[citation needed]

Heat transfer

NaK phase diagram, showing the melting point of sodium as a function of potassium concentration. NaK with 77% potassium is eutectic and has the lowest melting point of the NaK alloys at −12.6 °C.[53]

Liquid sodium is used as a heat transfer fluid in some fast reactors,[54] due to its high thermal conductivity and low neutron absorption cross section, which is required to achieve a high neutron flux; the high boiling point allows the reactor to operate at ambient pressure. Drawbacks of using sodium include its opacity, which hinders visual maintenance, and its explosive properties. Radioactive sodium-24 may be formed by neutron activation during operation, posing a slight radiation hazard; the radioactivity stops within a few days after removal from the reactor. If a reactor needs to be frequently shut down, NaK is used; due to it being liquid at room temperature, cooling pipes do not freeze. In this case, the pyrophoricity of potassium means extra precautions against leaks need to be taken. Another heat transfer application is in high-performance internal combustion engines with poppet valves, where valve stems partially filled with sodium are used as a heat pipe to cool the valves.

Biological role

Main article: Sodium in biology

In humans, sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day.[55] Sodium chloride is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods.[56] The UL for sodium is 2.3 grams per day,[57] the threshold which could lead to hypertension when exceeded,[58] but on average people in the United States consume 3.4 grams per day.[59] Hypertension causes 7.6 million premature deaths worldwide each year.[60] (Note that salt contains about 39.3% sodium[61]—the rest being chlorine and other trace chemicals; thus the UL of 2.3g sodium would be about 5.9g, or 2.7ml of salt—about half a US teaspoon[62])

The renin-angiotensin system regulates the amount of fluids and sodium in the body. Reduction of blood pressure and sodium concentration in the kidney result in the production of renin, which in turn produces aldosterone and angiotensin, retaining sodium in the urine. Because of the increase in sodium concentration, the production of renin decreases, and the sodium concentration returns to normal.[63] Sodium is also important in neuron function and osmoregulation between cells and the extracellular fluid, their distribution mediated in all animals by Na+/K+-ATPase, an active transporter pumping ions against the gradient, and sodium/potassium channels.[64] Sodium is the most prominent metal ion in extracellular fluid.[65]

Unusually low or high sodium levels in humans are recognized in medicine as hyponatremia and hypernatremia. These conditions may be caused by genetic factors, physical factors associated with ageing or illnesses involving vomiting or diarrhea.[66]

In C4 plants, sodium is a micronutrient that aids in metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll.[67] In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata.[68] Excess sodium in the soil limits the uptake of water due to decreased water potential, which may result in wilting; similar concentrations in the cytoplasm can lead to enzyme inhibition, which in turn causes necrosis and chlorosis.[69] To avoid these problems, plants developed mechanisms that limit sodium uptake by roots, store them in cell vacuoles, and control them over long distances;[70] excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

Precautions

NFPA 704
safety square
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g. sodium, sulfuric acid
3
1
2
W
The fire diamond hazard sign for sodium metal[71]

Care is required in handling elemental sodium, as it generates flammable hydrogen and caustic sodium hydroxide upon contact with water; powdered sodium may spontaneously explode in the presence of an oxidizer.[citation needed] Excess sodium can be safely removed by hydrolysis in a ventilated cabinet; this is typically done by sequential treatment with isopropanol, ethanol and water. Isopropanol reacts very slowly, generating the corresponding alkoxide and hydrogen.[72] Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane lose their effectiveness when they dissipate.[citation needed] An effective extinguishing agent is Met-L-X, which comprises approximately 5% Saran in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop.[citation needed] Other materials include Lith+, which has graphite powder and an organophosphate flame retardant, and dry sand.[citation needed]

See also

Sodium in biology

Alkali metal

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