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== Natrium == |
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{{about|the chemical element|the nutrient commonly called sodium|salt|other uses|sodium (disambiguation)}} |
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{{Infobox natrium}} |
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{{Unsur|Natrium|Na|11}} |
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'''Sodium''' is a [[chemical element]] with symbol '''Na''' (from Latin ''natrium'') and [[atomic number]] 11. It is a soft, silvery-white, highly reactive [[metal]]. Sodium is an [[alkali metal]], being in group 1 of the periodic table, because it has a single electron in its outer shell that it readily donates, creating a positively charged atom—the Na<sup>+</sup> [[cation]]. Its only stable [[isotope]] is <sup>23</sup>Na. The free metal does not occur in nature, but must be prepared from compounds. Sodium is the [[Abundance of elements in Earth's crust|sixth most abundant 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 [[Leaching (chemistry)|leached]] by the action of water from the [[Earth|Earth's]] [[mineral]]s over eons, and thus sodium and [[chlorine]] are the most common dissolved elements by weight in the oceans. |
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Sodium was first isolated by [[Humphry Davy]] in 1807 by the [[electrolysis]] of [[sodium hydroxide]]. Among many other useful sodium compounds, [[sodium hydroxide]] ([[lye]]) is used in [[Soap|soap manufacture]], and [[sodium chloride]] ([[edible salt]]) is a [[De-ice|de-icing]] agent and a nutrient for animals including humans. |
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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 ECF compartment volume. Loss of water from the ECF compartment increases the sodium concentration, a condition called [[hypernatremia]]. [[Tonicity|Isotonic]] loss of water and sodium from the ECF compartment decreases the size of that compartment in a condition called ECF [[hypovolemia]]. |
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By means of the [[Na+/K+-ATPase|sodium-potassium pump]], living human cells pump three sodium ions out of the cell in exchange for two potassium ions pumped in; comparing ion concentrations across the cell membrane, inside to outside, [[potassium]] measures about 40:1, and sodium, about 1:10. In [[Neuron|nerve cells]], the electrical charge across the cell membrane enables transmission of the nerve impulse—an [[action potential]]—when the charge is dissipated; sodium plays a key role in that activity. |
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==Characteristics== |
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===Physical=== |
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[[File:Na-D-sodium D-lines-589nm.jpg|thumb|left|[[Emission spectrum]] for sodium, showing the [[Fraunhofer lines|D line]].]] |
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Sodium at [[standard temperature and pressure]] is a soft silvery metal that combines with oxygen in air and forms grayish white [[sodium oxide]] unless immersed in oil or inert gas, which are the conditions it is usually stored in. Sodium metal can be easily cut with a knife and is a good conductor of electricity and heat because it has only one electron in its valence shell, resulting in weak [[metallic bonding]] and [[free electron]]s, which carry energy. Due to having low atomic mass and large atomic radius, sodium is third-least dense of all elemental metals and is one of only three metals that can float on water, the other two being lithium and potassium.<ref name=Greenwood75>Greenwood and Earnshaw, p. 75</ref> The melting (98 °C) and boiling (883 °C) points of sodium are lower than those of lithium but higher than those of the heavier alkali metals potassium, rubidium, and caesium, following periodic trends down the group.<ref>{{cite web|URL=http://www.encyclopedia.com/science-and-technology/chemistry/compounds-and-elements/alkali-metals|title="Alkali Metals." Science of Everyday Things|work=Encyclopedia.com|accessdate=15 October 2016}}</ref> 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> |
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[[File:Flametest--Na.swn.jpg|thumb|left|upright|A positive [[flame test]] for sodium has a bright yellow color.]] |
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In a [[flame test]], sodium and its compounds glow 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 [[D2 line|D line]] at about 589.3 nm. [[Spin-orbit interaction]]s involving the electron in the 3p orbital split the D line into two, at 589.0 and 589.6 nm; [[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|pages=1507–1512|issue=4|bibcode=1977PhRvA..16.1507C|last4=Stroud|first4=C.}}</ref> |
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===Isotopes=== |
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{{main article|Isotopes of sodium}} |
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Twenty isotopes of sodium are known, but only <sup>23</sup>Na is stable. <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> Two [[radioactive decay|radioactive]], [[cosmogenic]] isotopes are the byproduct of [[cosmic ray spallation]]: <sup>22</sup>Na has a [[half-life]] of 2.6 years and <sup>24</sup>Na, 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 milliseconds. Acute neutron radiation, as from a nuclear [[criticality accident]], converts some of the stable <sup>23</sup>Na in human blood to <sup>24</sup>Na; the neutron radiation dosage of a victim can be calculated by measuring the concentration of <sup>24</sup>Na relative to <sup>23</sup>Na.<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> |
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==Chemistry== |
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Sodium atoms have 11 electrons, one more than the extremely stable configuration of the [[noble gas]] [[neon]]. Because of this and its low first [[ionization energy]] of 495.8 kJ/mol, the sodium atom is much more likely to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge.<ref>{{cite book|title=Biology: Threads of Life|author=Sobrasua Ibim|publisher=Xlibris Corporation, 2010|isbn=1-4535-2068-6|page=27}}</ref> This process requires so little energy that sodium is readily oxidized by giving up its 11th electron. In contrast, the second ionization energy is very high (4562 kJ/mol), because the 10th electron is closer to the nucleus than the 11th electron. As a result, sodium usually forms [[ionic compound]]s involving the Na<sup>+</sup> cation.<ref>{{cite book|title=Cambridge International AS and A Level Chemistry Coursebook|author=Lawrie Ryan|author2= Roger Norris|publisher=Cambridge University Press, 2014|edition=illustrated|isbn=1-107-63845-3|page=36}}</ref> |
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The most common oxidation state for sodium is +1. It 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> Sodium metal is highly reducing, with the [[standard reduction potential]] for the Na<sup>+</sup>/Na couple being −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> though 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> |
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===Salts and oxides=== |
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{{Category see also|Sodium compounds}} |
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[[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 when dissolved in water and reassembles when the water evaporates.]] |
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Sodium compounds are of immense commercial importance, being particularly central to industries producing [[glass]], [[paper]], [[soap]], and [[textile]]s.<ref name=Ullmann/> The most important sodium compounds are [[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 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> |
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[[File:StericAcidChemStr.png|thumb|center|300px|Two equivalent images of the chemical structure of [[sodium stearate]], a typical soap.]] |
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Most [[soap]]s are sodium salts of [[fatty acid]]s. Sodium soaps have a higher melting temperature (and seem "harder") 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> |
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Like all the [[alkali metal]]s, sodium reacts [[Exothermic reaction|exothermically]] with water, and sufficiently large pieces melt to a sphere and may explode. The reaction produces caustic soda ([[sodium hydroxide]]) and flammable [[hydrogen]] gas. When burned in air, it forms primarily [[sodium peroxide]] with some [[sodium oxide]].<ref name=Greenwood84>Greenwoood and Earnshaw, p. 84</ref> |
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===Aqueous solutions=== |
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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'' = 4–8; with ''n'' = 6 indicated from X-ray diffraction data and computer simulations.<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> |
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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> |
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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|pages=3233–3237|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> |
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===Electrides and sodides=== |
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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>, with the positive charge counterbalanced by [[electride|electrons as anions]]; [[cryptand]]s permit the isolation of these complexes as crystalline solids. Sodium forms complexes with crown ethers, cryptands and other ligands.<ref>{{cite book|title=Applications of the Newer Techniques of Analysis|editor=Ivor L. Simmons|publisher=Springer Science & Business Media, 2012|isbn=1-4684-3318-0|page=160}}</ref> For example, [[15-crown-5]] has high affinity for sodium because the cavity size of 15-crown-5 is 1.7–2.2 Å, which is enough to fit sodium ion (1.9 Å).<ref>{{cite book|title=Design, Fabrication, Properties and Applications of Smart and Advanced Materials|editor=Xu Hou|publisher=CRC Press, 2016|edition=illustrated|isbn=1-4987-2249-0|page=175}}</ref><ref>{{cite book|title=Anionic Polymerization: Principles, Practice, Strength, Consequences and Applications|editor=Nikos Hadjichristidis|editor2=Akira Hirao| |
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publisher=Springer, 2015|edition=illustrated|isbn=4-431-54186-1|page=349}}</ref> Cryptands, like crown ethers and other [[ionophore]]s, also 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> |
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===Organosodium compounds=== |
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[[File:Monensin2.png|thumb|The structure of the complex of sodium (Na<sup>+</sup>, shown in yellow) and the antibiotic [[monensin]]-A.]] |
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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> Because of the large size and very low polarising power of the Na<sup>+</sup> cation, it can stabilize large, aromatic, polarisable radical anions, such as in [[sodium naphthalenide]], Na<sup>+</sup>[C<sub>10</sub>H<sub>8</sub>•]<sup>−</sup>, a strong reducing agent.<ref>Greenwood and Earnshaw, p. 111</ref> |
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===Intermetallic compounds=== |
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Sodium forms alloys with many metals, such as potassium, [[calcium]], [[lead]], and the [[group 11 element|group 11]] and [[group 12 elements|12]] elements. Sodium and potassium form KNa<sub>2</sub> and [[NaK]]. NaK is 40–90% potassium and it is liquid at [[ambient temperature]]. It is excellent thermal and electrical conductor. Sodium-calcium alloys are by-products of electrolytic production of sodium from binary salt mixture of NaCl-CaCl<sub>2</sub> and ternary mixture NaCl-CaCl<sub>2</sub>-BaCl<sub>2</sub>. Calcium is only partially [[miscible]] with sodium. In liquid state, sodium is completely miscible with lead. There are several methods to make sodium-lead alloys. One is to melt them together and another is to deposit sodium electrolycally on molten lead cathodes. NaPb<sub>3</sub>, NaPb, Na<sub>9</sub>Pb<sub>4</sub>, Na<sub>5</sub>Pb<sub>2</sub>, and Na<sub>15</sub>Pb<sub>4</sub> are some of the known sodium-lead alloys. Sodium also forms alloys with [[gold]] (NaAu<sub>2</sub>) and [[silver]] (NaAg<sub>2</sub>). Group 12 metals ([[zinc]], [[cadmium]] and [[Mercury (element)|mercury]]) are known to make alloys with sodium. NaZn<sub>13</sub> and NaCd<sub>2</sub> are alloys of zinc and cadmium. Sodium and mercury form NaHg, NaHg<sub>4</sub>, NaHg<sub>2</sub>, Na<sub>3</sub>Hg<sub>2</sub>, and Na<sub>3</sub>Hg.<ref>{{cite book|title=Alloys: Preparation, Properties, Applications|last=Habashi|first=Fathi|publisher=John Wiley & Sons, 2008|isbn=3-527-61192-4|pages=278–280}}</ref> |
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==History== |
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Because of its importance in human metabolism, salt has long been an important commodity, 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> Although sodium, sometimes called ''soda'', had long been recognized 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> In 1809, the German physicist and chemist [[Ludwig Wilhelm Gilbert]] proposed the names ''Natronium'' for Humphry Davy's "sodium" and ''Kalium'' for Davy's "potassium".<ref>Humphry Davy (1809) "Ueber einige neue Erscheinungen chemischer Veränderungen, welche durch die Electricität bewirkt werden; insbesondere über die Zersetzung der feuerbeständigen Alkalien, die Darstellung der neuen Körper, welche ihre Basen ausmachen, und die Natur der Alkalien überhaupt" (On some new phenomena of chemical changes that are achieved by electricity; particularly the decomposition of flame-resistant alkalis [i.e., alkalies that cannot be reduced to their base metals by flames], the preparation of new substances that constitute their [metallic] bases, and the nature of alkalies generally), ''Annalen der Physik'', '''31''' (2) : 113–175 ; [https://books.google.com/books?id=vyswAAAAYAAJ&pg=PA157#v=onepage&q&f=false see footnote p. 157.] From p. 157: ''"In unserer deutschen Nomenclatur würde ich die Namen ''Kalium'' und ''Natronium'' vorschlagen, wenn man nicht lieber bei den von Herrn Erman gebrauchten und von mehreren angenommenen Benennungen ''Kali-Metalloid'' and ''Natron-Metalloid'', bis zur völligen Aufklärung der chemischen Natur dieser räthzelhaften Körper bleiben will. Oder vielleicht findet man es noch zweckmässiger fürs Erste zwei Klassen zu machen, ''Metalle'' und ''Metalloide'', und in die letztere ''Kalium'' und ''Natronium'' zu setzen. — Gilbert."'' (In our German nomenclature, I would suggest the names ''Kalium'' and ''Natronium'', if one would not rather continue with the appellations ''Kali-metalloid'' and ''Natron-metalloid'' which are used by Mr. Erman and accepted by several [people], until the complete clarification of the chemical nature of these puzzling substances. Or perhaps one finds it yet more advisable for the present to create two classes, ''metals'' and ''metalloids'', and to place ''Kalium'' and ''Natronium'' in the latter — Gilbert.)</ref> The chemical abbreviation for sodium was first published in 1814 by [[Jöns Jakob Berzelius]] in his system of atomic symbols,<ref>J. Jacob Berzelius, ''Försök, att, genom användandet af den electrokemiska theorien och de kemiska proportionerna, grundlägga ett rent vettenskapligt system för mineralogien'' [Attempt, by the use of electrochemical theory and chemical proportions, to found a pure scientific system for mineralogy] (Stockholm, Sweden: A. Gadelius, 1814), [https://archive.org/stream/bub_gb_Uw0-AAAAcAAJ#page/n91/mode/2up p. 87.]</ref><ref>{{cite web|url=http://elements.vanderkrogt.net/element.php?sym=Na|title=Elementymology & Elements Multidict|first = Peter|last = van der Krogt|accessdate=2007-06-08}}</ref> and is an abbreviation of the element's [[New Latin]] name ''natrium'', which refers to the Egyptian ''[[natron]]'',<ref name=newton/> a natural mineral salt mainly consisting of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds.<ref>{{cite journal|url=http://www.sciencedirect.com/science/article/pii/S0305440305002074|title= Natron as a flux in the early vitreous materials industry: sources, beginnings and reasons for decline|author=Andrew Shortland, Lukas Schachner, Ian Freestone, and Michael Tite | doi=10.1016/j.jas.2005.09.011|volume=33|journal=Journal of Archaeological Science|pages=521–530}}</ref> |
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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> |
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<blockquote>In a corner of our 60 m<sup>3</sup> 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.</blockquote> |
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==Occurrence== |
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The Earth's crust contains 2.27% sodium, making it the [[Abundance of the chemical elements|seventh most abundant element]] on Earth and the fifth most abundant metal, behind [[aluminium]], [[iron]], [[calcium]], and [[magnesium]] and ahead of potassium.<ref name=Greenwood69>Greenwood and Earnshaw, p. 69</ref> Sodium's estimated oceanic abundance is 1.08{{e|4}} milligrams per liter.<ref name = "abundancecrc">{{Cite book|url=https://books.google.com/books?id=kTnxSi2B2FcC|title=CRC Handbook of Chemistry and Physics, 84th Edition|last=Lide|first=David R.|date=2003-06-19|publisher=CRC Press|isbn=978-0-8493-0484-2|series=[[CRC Handbook]]|at=14: Abundance of Elements in the Earth's Crust and in the Sea|language=en}}</ref> 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. |
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===Astronomical observations=== |
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In the [[interstellar medium]], sodium is identified by the D [[Fraunhofer lines|spectral line]]; though it has a high vaporization temperature, its abundance in [[Mercury (planet)|Mercury's]] atmosphere enabled its detection by [[Mariner 10|Potter and Morgan using ground-based high resolution spectroscopy]].{{citation-needed|date=March 2017}} Sodium has been detected in at least one [[comet]]; astronomers watching [[Comet Hale-Bopp]] in 1997 observed a sodium [[Comet tail|tail]] consisting of neutral atoms (not ions) and extending to some 50 million kilometres behind the head.<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> |
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==Commercial production== |
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Employed only in rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.<ref name=Ullmann/> Metallic sodium was first produced commercially in the late 19th century<ref>{{cite book|title=Speciality Chemicals: Innovations in industrial synthesis and applications|publisher=Springer Science & Business Media, 1991|editor=B. Pearson|isbn=1-85166-646-X|page=260|edition=illustrated}}</ref> by [[carbothermal reduction]] of [[sodium carbonate]] at 1100 °C, as the first step of the [[Deville process]] for the production of aluminium:<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. --> |
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:Na<sub>2</sub>CO<sub>3</sub> + 2 C → 2 Na + 3 CO |
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The high demand of aluminium created the need for the production of sodium. After the introduction of the [[Hall–Héroult process]] for the production of aluminium by [[Electrolysis|electrolysing]] a molten salt bath ended the need for large quantities of sodium. A related process based on the reduction of sodium hydroxide was developed in 1886.<ref name=kirk/> |
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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 °C. As [[calcium]] is less [[electropositive]] than sodium, no calcium will be deposited at the cathode.<ref>{{cite book|title=Sodium Metal from France|publisher=DIANE Publishing|isbn=1-4578-1780-2}}</ref> This method is less expensive than the previous [[Castner process]] (the electrolysis of [[sodium hydroxide]]).<ref>{{cite book|title=Industrial Chemistry: For Advanced Students|author=Mark Anthony Benvenuto|publisher=Walter de Gruyter GmbH & Co KG, 2015|edition=illustrated|isbn=3-11-038339-X}}</ref> |
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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]].<ref>{{cite book|title=Active Pharmaceutical Ingredients: Development, Manufacturing, and Regulation, Second Edition|publisher=CRC Press, 2016|isbn=1-4398-0339-0|editor=Stanley Nusim|edition=2, illustrated, revised|page=303}}</ref> |
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==Applications== |
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Though metallic sodium has some important uses, the major applications for sodium use compounds; millions of tons of [[sodium chloride]], [[sodium hydroxide|hydroxide]], and [[sodium carbonate|carbonate]] are produced annually. 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]]; though potassium is the better ion in most cases, sodium is chosen 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> |
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Metallic sodium is used mainly for the production of [[sodium borohydride]], [[sodium azide]], [[indigo dye|indigo]], and [[triphenylphosphine]]. A once-common use was the making of [[tetraethyllead]] and titanium metal; because of the move away from TEL and new titanium production methods, 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|Sodium plasma ("vapor") lamp]]s are often used for street lighting in cities, shedding light that ranges 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 coloration 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> Sodium reacts with alcohol and gives alkoxides, and when sodium is dissolved in ammonia solution, it can be used to reduce alkynes to trans-alkenes.<ref>{{cite book|title= Organic Synthesis|last=Smith|first=Michael|publisher= Academic Press, 2011|edition=3|isbn=0-12-415884-6|page=455}}</ref><ref>{{cite book|title= Organic Chemistry|authors= Solomons & Fryhle|publisher= John Wiley & Sons, 2006|edition=8|isbn=81-265-1050-1|page=272}}</ref> Lasers emitting light at the sodium D line are used to create artificial [[laser guide star]]s that [[FASOR (laser physics)|assist]] in the [[adaptive optics]] for land-based visible light telescopes.<ref>{{cite web|url=http://www.eso.org/sci/publications/messenger/archive/no.139-mar10/messenger-no139-12-19.pdf|work=Domenico Bonaccini Calia, Yan Feng, Wolfgang Hackenberg, Ronald Holzlöhner, Luke Taylor, Steffan Lewis|title=Laser Development for Sodium Laser Guide Stars at ESO|format=PDF}}</ref> |
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===Heat transfer=== |
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[[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>]] |
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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> because it has the high thermal conductivity and low neutron absorption [[cross section (physics)|cross section]] required to achieve a high neutron flux in the reactor.<ref name="nrc">{{cite web|url=https://www.nrc.gov/docs/ML1504/ML15043A307.pdf|title=Sodium-cooled Fast Reactor (SFR)|work=Office of Nuclear Energy, U.S. Department of Energy|date= 18 February 2015|format=PDF}}</ref> The high boiling point of sodium allows the reactor to operate at ambient (normal) pressure,<ref name="nrc" /> but the drawbacks include its opacity, which hinders visual maintenance, and its explosive properties.<ref>{{cite book|title=Fire and Explosion Hazards|publisher=Research Publishing Service, 2011|isbn=981-08-7724-2|page=363}}</ref> Radioactive [[sodium-24]] may be produced by [[neutron activation|neutron bombardment]] during operation, posing a slight radiation hazard; the radioactivity stops within a few days after removal from the reactor.<ref>{{cite book|title=Simulation and Optimization Methods in Risk and Reliability Theory|publisher=Nova Science Publishers, 2009|editors=Pavel Solomonovich Knopov, Panos M. Pardalos|isbn=1-60456-658-2|page=150}}</ref> If a reactor needs to be shut down frequently, [[NaK]] is used; because NaK is a liquid at room temperature, the coolant does not solidify in the pipes.<ref>{{cite book|title=Proceedings of the Heat Transfer and Fluid Mechanics Institute|last= McKillop|first=Allan A.|publisher=Stanford University Press, 1976|page=97|isbn=0-8047-0917-3}}</ref> In this case, the [[pyrophoricity]] of potassium requires extra precautions to prevent and detect leaks.<ref>{{cite book|title=Reactor Handbook: Engineering|author=U.S. Atomic Energy Commission|publisher=Interscience Publishers|page=325|edition=2}}</ref> Another heat transfer application is [[poppet valve]]s in high-performance internal combustion engines; the valve stems are partially filled with sodium and work as a [[heat pipe]] to cool the valves.<ref>{{cite patent|country=US|number=US2949907 A| title = Coolant-filled poppet valve and method of making same| pubdate = 23 Aug 1960| inventor =Tauschek Max J}}</ref> |
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==Biological role== |
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{{main article|Sodium in biology}} |
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In humans, sodium is an essential mineral that regulates [[blood]] volume, blood pressure, [[Osmosis|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 in such commodities as [[pickling|pickled preserves]] and [[Jerky (food)|jerky]]; for Americans, most sodium chloride comes from [[Convenience food|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> Other sources of sodium are its natural occurrence in food and such food additives as [[monosodium glutamate]] (MSG), [[sodium nitrite]], sodium saccharin, [[baking soda]] (sodium bicarbonate), and [[sodium benzoate]].<ref name=medline>{{cite web|url=https://medlineplus.gov/ency/article/002415.htm|title=Sodium in diet|publisher=MedlinePlus, US National Library of Medicine|date=5 October 2016}}</ref> The US [[Institute of Medicine]] set its [[Dietary Reference Intake|Tolerable Upper Intake Level]] for sodium at 2.3 grams per day,<ref>{{cite web|title=Reference Values for Elements|url=http://www.hc-sc.gc.ca/fn-an/nutrition/reference/table/ref_elements_tbl-eng.php|website=Dietary Reference Intakes Tables|publisher=Health Canada}}</ref> but the average person in the United States consumes 3.4 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=https://web.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> Studies have found that lowering sodium intake by 2 g per day tends to lower [[systolic blood pressure]] by about two to four mm Hg.<ref name=Impact>{{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|url=http://eurpub.oxfordjournals.org/content/eurpub/14/3/235.full.pdf}}</ref> It has been estimated that such a decrease in sodium intake would lead to between 9 and 17% fewer cases of [[hypertension]].<ref name=Impact/> |
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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 trace chemicals; thus, 2.3 g sodium is about 5.9 g, or 2.7 ml of salt{{--}}about half a [[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><ref name="fda">{{cite web | url=http://www.fda.gov/Food/IngredientsPackagingLabeling/LabelingNutrition/ucm315393.htm | title=Sodium in Your Diet: Use the Nutrition Facts Label and Reduce Your Intake | publisher=US Food and Drug Administration | date=2 June 2016 | accessdate=15 October 2016}}</ref>) The [[American Heart Association]] recommends no more than 1.5 g of sodium per day.<ref>{{cite web | url=http://sodiumbreakup.heart.org/sodium-411/how-much-sodium-do-you-need/ | title=How much sodium should I eat per day? | publisher=American Heart Association | date=2016 | accessdate=15 October 2016}}</ref> |
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One study found that people with or without hypertension who excreted less than 3 grams of sodium per day in their urine (and therefore were taking in less than 3 g/d) had a ''higher'' risk of death, stroke, or heart attack than those excreting 4 to 5 grams per day. Levels of 7 g per day or more in people with hypertension were associated with higher mortality and cardiovascular events, but this was not found to be true for people without [[hypertension]].<ref>{{cite journal|author1=Andrew Mente|author2=''et al.''|title=Associations of urinary sodium excretion with cardiovascular events in individuals with and without hypertension: a pooled analysis of data from four studies|journal=The Lancet|date=2016|doi=10.1016/S0140-6736(16)30467-6|pmid=27216139|volume=388|pages=465–75}}</ref> The [[Food and Drug Administration|US FDA]] states that adults with hypertension and prehypertension should reduce daily intake to 1.5 g.<ref name=fda/> |
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The [[renin-angiotensin system]] regulates the amount of fluid and sodium concentration 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. When the concentration of sodium increases, 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> The sodium ion (Na<sup>+</sup>) is an important electrolyte in [[neuron]] function, and in osmoregulation between cells and the [[extracellular fluid]]. This is accomplished 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 prevalent metallic 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> |
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Unusually low or high sodium levels in humans are recognized in medicine as [[hyponatremia]] and [[hypernatremia]]. These conditions may be caused by genetic factors, ageing, or prolonged vomiting or diarrhea.<ref> |
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{{cite book |
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|last1=Pohl |
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|first1=Hanna R. |
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|last2=Wheeler |
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|first2=John S. |
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|first3=H. Edward |
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|last3=Murray |
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|editor=Astrid Sigel |
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|editor2=Helmut Sigel |
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|editor3=Roland K. O. Sigel |
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|title=Interrelations between Essential Metal Ions and Human Diseases |
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|series=Metal Ions in Life Sciences |
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|volume=13 |
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|date=2013 |
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|publisher=Springer |
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|pages=29–47 |
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|doi=10.1007/978-94-007-7500-8_2 |
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}} |
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</ref> |
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In [[C4 plants]], sodium is a [[micronutrient]] that aids 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 [[stoma]]ta.<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 can limit the uptake of water by decreasing the [[water potential]], which may result in plant wilting; excess 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> In response, some plants have developed mechanisms to limit sodium uptake in the roots, to store it in cell [[vacuole]]s, and restrict salt transport from roots to leaves;<ref name="halo">{{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. [[Halophytes]] have adapted to be able to flourish in sodium rich environments.<ref name="halo"/> |
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==Safety and precautions== |
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{{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>}} |
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Sodium forms flammable hydrogen and caustic [[sodium hydroxide]] on contact with water;<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> ingestion and contact with moisture on skin, eyes or [[mucous membrane]]s can cause severe burns.<ref>{{cite book|title=Sodium Explosion Critically Burns Firefighters: Newton, Massachusetts|work=U. S. Fire Administration|last=Routley|first=J. Gordon|publisher=FEMA, 2013}}</ref><ref name="prudent">{{cite book|title=Prudent Practices in the Laboratory: Handling and Disposal of Chemicals|work=National Research Council (U.S.). Committee on Prudent Practices for Handling, Storage, and Disposal of Chemicals in Laboratories|publisher=National Academies, 1995|page=390}}</ref> Sodium spontaneously explodes in the presence of an oxidizer such as water.<ref>{{Cite web|url=http://www.heart.org/HEARTORG/HealthyLiving/HealthyEating/Nutrition/Sodium-and-Salt_UCM_303290_Article.jsp|title=Sodium and Salt|website=www.heart.org|access-date=2016-09-05}}</ref> [[Fire extinguisher]]s based on water accelerate sodium fires; those based on carbon dioxide and [[bromochlorodifluoromethane]] should not be used on sodium fire.<ref name="prudent"/> Metal fires are [[Fire extinguisher|Class D]], but not all Class D extinguishers are workable with sodium. An effective extinguishing agent for sodium fires is [[Met-L-X]].<ref name="prudent"/> Other effective agents include Lith-X, which has [[graphite]] powder and an [[organophosphate]] [[flame retardant]], and dry sand.<ref>{{cite book|title=Industrial fire prevention and protection|last= Ladwig|first=Thomas H.|publisher=Van Nostrand Reinhold, 1991|isbn=0-442-23678-6|page=178}}</ref> Sodium fires are prevented in nuclear reactors by isolating sodium from oxygen by surrounding sodium pipes with inert gas.<ref name="fission">{{cite book|title=Sustainable and Safe Nuclear Fission Energy: Technology and Safety of Fast and Thermal Nuclear Reactors|author=Günter Kessler|publisher=Springer Science & Business Media, 2012|isbn=3-642-11990-5|page=446|edition=illustrated}}</ref> Pool-type sodium fires are prevented using different design measures called catch pan systems. They collect leaking sodium into a leak-recovery tank where it is isolated from oxygen.<ref name="fission"/> |
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==See also== |
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{{Subject bar |
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|portal=Chemistry |
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|book1=Sodium |
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|book2=Period 3 elements |
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==References== |
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{{reflist|30em}} |
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==Bibliography== |
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*{{Greenwood&Earnshaw2nd}} |
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==External links== |
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* [http://www.periodicvideos.com/videos/011.htm Sodium] at ''[[The Periodic Table of Videos]]'' (University of Nottingham) |
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* [http://www.balashon.com/2008/07/neter-and-nitrogen.html Etymology of "natrium" – source of symbol Na] |
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* [http://www.theodoregray.com/PeriodicTable/Elements/011/index.html The Wooden Periodic Table Table's Entry on Sodium] |
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* [http://ie.lbl.gov/education/parent/Na_iso.htm Sodium isotopes data from ''The Berkeley Laboratory Isotopes Project's''] |
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Revisi per 1 Agustus 2017 10.01
Natrium
11Na Natrium | |||||||||||||||||||||
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.jpg) Logam natrium murni | |||||||||||||||||||||
 Garis spektrum natrium | |||||||||||||||||||||
| Sifat umum | |||||||||||||||||||||
| Pengucapan | /natrium/[1] | ||||||||||||||||||||
| Penampilan | metalik putih keperakan | ||||||||||||||||||||
| Natrium dalam tabel periodik | |||||||||||||||||||||
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| Nomor atom (Z) | 11 | ||||||||||||||||||||
| Golongan | golongan 1 (logam alkali) | ||||||||||||||||||||
| Periode | periode 3 | ||||||||||||||||||||
| Blok | blok-s | ||||||||||||||||||||
| Kategori unsur | logam alkali | ||||||||||||||||||||
| Berat atom standar (Ar) |
| ||||||||||||||||||||
| Konfigurasi elektron | [Ne] 3s1 | ||||||||||||||||||||
| Elektron per kelopak | 2,8,1 | ||||||||||||||||||||
| Sifat fisik | |||||||||||||||||||||
| Fase pada STS (0 °C dan 101,325 kPa) | padat | ||||||||||||||||||||
| Titik lebur | 370,944 K (97,794 °C, 208,029 °F) | ||||||||||||||||||||
| Titik didih | 1156,090 K (882,940 °C, 1621,292 °F) | ||||||||||||||||||||
| Kepadatan mendekati s.k. | 0,968 g/cm3 | ||||||||||||||||||||
| saat cair, pada t.l. | 0,927 g/cm3 | ||||||||||||||||||||
| Titik kritis | 2573 K, 35 MPa (diekstrapolasi) | ||||||||||||||||||||
| Kalor peleburan | 2,60 kJ/mol | ||||||||||||||||||||
| Kalor penguapan | 97,42 kJ/mol | ||||||||||||||||||||
| Kapasitas kalor molar | 28,230 J/(mol·K) | ||||||||||||||||||||
Tekanan uap
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| Sifat atom | |||||||||||||||||||||
| Bilangan oksidasi | −1, +1 (oksida basa kuat) | ||||||||||||||||||||
| Elektronegativitas | Skala Pauling: 0,93 | ||||||||||||||||||||
| Energi ionisasi | ke-1: 495,8 kJ/mol ke-2: 4562 kJ/mol ke-3: 6910,3 kJ/mol (artikel) | ||||||||||||||||||||
| Jari-jari atom | empiris: 186 pm | ||||||||||||||||||||
| Jari-jari kovalen | 166±9 pm | ||||||||||||||||||||
| Jari-jari van der Waals | 227 pm | ||||||||||||||||||||
| Lain-lain | |||||||||||||||||||||
| Kelimpahan alami | primordial | ||||||||||||||||||||
| Struktur kristal | kubus berpusat badan (bcc) | ||||||||||||||||||||
| Kecepatan suara batang ringan | 3200 m/s (suhu 20 °C) | ||||||||||||||||||||
| Ekspansi kalor | 71 µm/(m·K) (suhu 25 °C) | ||||||||||||||||||||
| Konduktivitas termal | 142 W/(m·K) | ||||||||||||||||||||
| Resistivitas listrik | 47,7 nΩ·m (suhu 20 °C) | ||||||||||||||||||||
| Arah magnet | paramagnetik[2] | ||||||||||||||||||||
| Suseptibilitas magnetik molar | +16,0×10−6 cm3/mol (298 K)[3] | ||||||||||||||||||||
| Modulus Young | 10 GPa | ||||||||||||||||||||
| Modulus Shear | 3,3 GPa | ||||||||||||||||||||
| Modulus curah | 6,3 GPa | ||||||||||||||||||||
| Skala Mohs | 0,5 | ||||||||||||||||||||
| Skala Brinell | 0,69 MPa | ||||||||||||||||||||
| Nomor CAS | 7440-23-5 | ||||||||||||||||||||
| Sejarah | |||||||||||||||||||||
| Penamaan | dari Jerman Natron, 'natron' | ||||||||||||||||||||
| Penemuan | H. Davy (1807) | ||||||||||||||||||||
| Isolasi pertama | H. Davy (1807) | ||||||||||||||||||||
| Isotop natrium yang utama | |||||||||||||||||||||
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Natrium adalah unsur kimia dengan lambang Na dan nomor atom 11. Sodium is a chemical element with symbol Na (from Latin natrium) and atomic number 11. It is a soft, silvery-white, highly reactive metal. Sodium is an alkali metal, being in group 1 of the periodic table, because it has a single electron in its outer shell that it readily donates, creating a positively charged atom—the Na+ cation. Its only stable isotope is 23Na. The free metal does not occur in nature, but must be prepared from 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, and thus 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. Among many other useful sodium compounds, sodium hydroxide (lye) is used in soap manufacture, and sodium chloride (edible salt) is a de-icing agent and a nutrient for animals including humans.
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 ECF compartment volume. Loss of water from the ECF compartment increases the sodium concentration, a condition called hypernatremia. Isotonic loss of water and sodium from the ECF compartment decreases the size of that compartment in a condition called ECF hypovolemia.
By means of the sodium-potassium pump, living human cells pump three sodium ions out of the cell in exchange for two potassium ions pumped in; comparing ion concentrations across the cell membrane, inside to outside, potassium measures about 40:1, and sodium, about 1:10. In nerve cells, the electrical charge across the cell membrane enables transmission of the nerve impulse—an action potential—when the charge is dissipated; sodium plays a key role in that activity.
Characteristics
Physical
Sodium at standard temperature and pressure is a soft silvery metal that combines with oxygen in air and forms grayish white sodium oxide unless immersed in oil or inert gas, which are the conditions it is usually stored in. Sodium metal can be easily cut with a knife and is a good conductor of electricity and heat because it has only one electron in its valence shell, resulting in weak metallic bonding and free electrons, which carry energy. Due to having low atomic mass and large atomic radius, sodium is third-least dense of all elemental metals and is one of only three metals that can float on water, the other two being lithium and potassium.[4] The melting (98 °C) and boiling (883 °C) points of sodium are lower than those of lithium but higher than those of the heavier alkali metals potassium, rubidium, and caesium, following periodic trends down the group.[5] 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.[6]
In a flame test, sodium and its compounds glow yellow[7] 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 about 589.3 nm. Spin-orbit interactions involving the electron in the 3p orbital split the D line into two, at 589.0 and 589.6 nm; hyperfine structures involving both orbitals cause many more lines.[8]
Isotopes
Twenty isotopes of sodium are known, but only 23Na is stable. 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.[9] Two radioactive, cosmogenic isotopes are the byproduct of cosmic ray spallation: 22Na has a half-life of 2.6 years and 24Na, a half-life of 15 hours; all other isotopes have a half-life of less than one minute.[10] Two nuclear isomers have been discovered, the longer-lived one being 24mNa with a half-life of around 20.2 milliseconds. Acute neutron radiation, as from a nuclear criticality accident, converts some of the stable 23Na in human blood to 24Na; the neutron radiation dosage of a victim can be calculated by measuring the concentration of 24Na relative to 23Na.[11]
Chemistry
Sodium atoms have 11 electrons, one more than the extremely stable configuration of the noble gas neon. Because of this and its low first ionization energy of 495.8 kJ/mol, the sodium atom is much more likely to lose the last electron and acquire a positive charge than to gain one and acquire a negative charge.[12] This process requires so little energy that sodium is readily oxidized by giving up its 11th electron. In contrast, the second ionization energy is very high (4562 kJ/mol), because the 10th electron is closer to the nucleus than the 11th electron. As a result, sodium usually forms ionic compounds involving the Na+ cation.[13]
The most common oxidation state for sodium is +1. It is generally less reactive than potassium and more reactive than lithium.[14] Sodium metal is highly reducing, with the standard reduction potential for the Na+/Na couple being −2.71 volts,[15] though potassium and lithium have even more negative potentials.[16]
Salts and oxides
Sodium compounds are of immense commercial importance, being particularly central to industries producing glass, paper, soap, and textiles.[17] The most important sodium compounds are 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).[18] In compounds, sodium is usually ionically bonded to water and anions, and is viewed as a hard Lewis acid.[19]
Most soaps are sodium salts of fatty acids. Sodium soaps have a higher melting temperature (and seem "harder") than potassium soaps.[18]
Like all the alkali metals, sodium reacts exothermically with water, and sufficiently large pieces melt to a sphere and may explode. The reaction produces caustic soda (sodium hydroxide) and flammable hydrogen gas. When burned in air, it forms primarily sodium peroxide with some sodium oxide.[20]
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 = 4–8; with n = 6 indicated from X-ray diffraction data and computer simulations.[21]
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).[22] 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.[23] Crown ethers, like 15-crown-5, may be used as a phase-transfer catalyst.[24]
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.[25] Lower concentrations of sodium may be determined by atomic absorption spectrophotometry[26] or by potentiometry using ion-selective electrodes.[27]
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)+, with the positive charge counterbalanced by electrons as anions; cryptands permit the isolation of these complexes as crystalline solids. Sodium forms complexes with crown ethers, cryptands and other ligands.[28] For example, 15-crown-5 has high affinity for sodium because the cavity size of 15-crown-5 is 1.7–2.2 Å, which is enough to fit sodium ion (1.9 Å).[29][30] Cryptands, like crown ethers and other ionophores, also have a high affinity for the sodium ion; derivatives of the alkalide Na− are obtainable[31] by the addition of cryptands to solutions of sodium in ammonia via disproportionation.[32]
Organosodium compounds
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).[33] Because of the large size and very low polarising power of the Na+ cation, it can stabilize large, aromatic, polarisable radical anions, such as in sodium naphthalenide, Na+[C10H8•]−, a strong reducing agent.[34]
Intermetallic compounds
Sodium forms alloys with many metals, such as potassium, calcium, lead, and the group 11 and 12 elements. Sodium and potassium form KNa2 and NaK. NaK is 40–90% potassium and it is liquid at ambient temperature. It is excellent thermal and electrical conductor. Sodium-calcium alloys are by-products of electrolytic production of sodium from binary salt mixture of NaCl-CaCl2 and ternary mixture NaCl-CaCl2-BaCl2. Calcium is only partially miscible with sodium. In liquid state, sodium is completely miscible with lead. There are several methods to make sodium-lead alloys. One is to melt them together and another is to deposit sodium electrolycally on molten lead cathodes. NaPb3, NaPb, Na9Pb4, Na5Pb2, and Na15Pb4 are some of the known sodium-lead alloys. Sodium also forms alloys with gold (NaAu2) and silver (NaAg2). Group 12 metals (zinc, cadmium and mercury) are known to make alloys with sodium. NaZn13 and NaCd2 are alloys of zinc and cadmium. Sodium and mercury form NaHg, NaHg4, NaHg2, Na3Hg2, and Na3Hg.[35]
History
Because of its importance in human metabolism, salt has long been an important commodity, 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.[36] Although sodium, sometimes called soda, had long been recognized in compounds, the metal itself was not isolated until 1807 by Sir Humphry Davy through the electrolysis of sodium hydroxide.[37][38] In 1809, the German physicist and chemist Ludwig Wilhelm Gilbert proposed the names Natronium for Humphry Davy's "sodium" and Kalium for Davy's "potassium".[39] The chemical abbreviation for sodium was first published in 1814 by Jöns Jakob Berzelius in his system of atomic symbols,[40][41] and is an abbreviation of the element's New Latin name natrium, which refers to the Egyptian natron,[36] a natural mineral salt mainly consisting of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds.[42]
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:[43]
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.
Occurrence
The Earth's crust contains 2.27% sodium, making it the seventh most abundant element on Earth and the fifth most abundant metal, behind aluminium, iron, calcium, and magnesium and ahead of potassium.[44] Sodium's estimated oceanic abundance is 1.08×104 milligrams per liter.[45] 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.
Astronomical observations
In the interstellar medium, sodium is identified by the D spectral line; though it has a high vaporization temperature, its abundance in Mercury's atmosphere enabled its detection by Potter and Morgan using ground-based high resolution spectroscopy.[butuh rujukan] Sodium has been detected in at least one comet; astronomers watching Comet Hale-Bopp in 1997 observed a sodium tail consisting of neutral atoms (not ions) and extending to some 50 million kilometres behind the head.[46]
Commercial production
Employed only in rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.[17] Metallic sodium was first produced commercially in the late 19th century[47] by carbothermal reduction of sodium carbonate at 1100 °C, as the first step of the Deville process for the production of aluminium:[48][49][50]
- Na2CO3 + 2 C → 2 Na + 3 CO
The high demand of aluminium created the need for the production of sodium. After the introduction of the Hall–Héroult process for the production of aluminium by electrolysing a molten salt bath ended the need for large quantities of sodium. A related process based on the reduction of sodium hydroxide was developed in 1886.[48]
Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924.[51][52] 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.[53] This method is less expensive than the previous Castner process (the electrolysis of sodium hydroxide).[54]
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.[55]
Applications
Though metallic sodium has some important uses, the major applications for sodium use compounds; millions of tons of sodium chloride, hydroxide, and carbonate are produced annually. 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; though potassium is the better ion in most cases, sodium is chosen for its lower price and atomic weight.[56] 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.[57]
Metallic sodium is used mainly for the production of sodium borohydride, sodium azide, indigo, and triphenylphosphine. A once-common use was the making of tetraethyllead and titanium metal; because of the move away from TEL and new titanium production methods, the production of sodium declined after 1970.[17] Sodium is also used as an alloying metal, an anti-scaling agent,[58] 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 plasma ("vapor") lamps are often used for street lighting in cities, shedding light that ranges from yellow-orange to peach as the pressure increases.[59] By itself or with potassium, sodium is a desiccant; it gives an intense blue coloration with benzophenone when the desiccate is dry.[60] 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.[61] Sodium reacts with alcohol and gives alkoxides, and when sodium is dissolved in ammonia solution, it can be used to reduce alkynes to trans-alkenes.[62][63] Lasers emitting light at the sodium D line are used to create artificial laser guide stars that assist in the adaptive optics for land-based visible light telescopes.[64]
Heat transfer
Liquid sodium is used as a heat transfer fluid in some fast reactors[66] because it has the high thermal conductivity and low neutron absorption cross section required to achieve a high neutron flux in the reactor.[67] The high boiling point of sodium allows the reactor to operate at ambient (normal) pressure,[67] but the drawbacks include its opacity, which hinders visual maintenance, and its explosive properties.[68] Radioactive sodium-24 may be produced by neutron bombardment during operation, posing a slight radiation hazard; the radioactivity stops within a few days after removal from the reactor.[69] If a reactor needs to be shut down frequently, NaK is used; because NaK is a liquid at room temperature, the coolant does not solidify in the pipes.[70] In this case, the pyrophoricity of potassium requires extra precautions to prevent and detect leaks.[71] Another heat transfer application is poppet valves in high-performance internal combustion engines; the valve stems are partially filled with sodium and work as a heat pipe to cool the valves.[72]
Biological role
In humans, sodium is an essential mineral that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day.[73] Sodium chloride is the principal source of sodium in the diet, and is used as seasoning and preservative in such commodities as pickled preserves and jerky; for Americans, most sodium chloride comes from processed foods.[74] Other sources of sodium are its natural occurrence in food and such food additives as monosodium glutamate (MSG), sodium nitrite, sodium saccharin, baking soda (sodium bicarbonate), and sodium benzoate.[75] The US Institute of Medicine set its Tolerable Upper Intake Level for sodium at 2.3 grams per day,[76] but the average person in the United States consumes 3.4 grams per day.[77] Studies have found that lowering sodium intake by 2 g per day tends to lower systolic blood pressure by about two to four mm Hg.[78] It has been estimated that such a decrease in sodium intake would lead to between 9 and 17% fewer cases of hypertension.[78]
Hypertension causes 7.6 million premature deaths worldwide each year.[79] (Note that salt contains about 39.3% sodium[80]—the rest being chlorine and trace chemicals; thus, 2.3 g sodium is about 5.9 g, or 2.7 ml of salt—about half a US teaspoon.[81][82]) The American Heart Association recommends no more than 1.5 g of sodium per day.[83]
One study found that people with or without hypertension who excreted less than 3 grams of sodium per day in their urine (and therefore were taking in less than 3 g/d) had a higher risk of death, stroke, or heart attack than those excreting 4 to 5 grams per day. Levels of 7 g per day or more in people with hypertension were associated with higher mortality and cardiovascular events, but this was not found to be true for people without hypertension.[84] The US FDA states that adults with hypertension and prehypertension should reduce daily intake to 1.5 g.[82]
The renin-angiotensin system regulates the amount of fluid and sodium concentration 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. When the concentration of sodium increases, the production of renin decreases, and the sodium concentration returns to normal.[85] The sodium ion (Na+) is an important electrolyte in neuron function, and in osmoregulation between cells and the extracellular fluid. This is accomplished in all animals by Na+/K+-ATPase, an active transporter pumping ions against the gradient, and sodium/potassium channels.[86] Sodium is the most prevalent metallic ion in extracellular fluid.[87]
Unusually low or high sodium levels in humans are recognized in medicine as hyponatremia and hypernatremia. These conditions may be caused by genetic factors, ageing, or prolonged vomiting or diarrhea.[88]
In C4 plants, sodium is a micronutrient that aids metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll.[89] In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata.[90] Excess sodium in the soil can limit the uptake of water by decreasing the water potential, which may result in plant wilting; excess concentrations in the cytoplasm can lead to enzyme inhibition, which in turn causes necrosis and chlorosis.[91] In response, some plants have developed mechanisms to limit sodium uptake in the roots, to store it in cell vacuoles, and restrict salt transport from roots to leaves;[92] excess sodium may also be stored in old plant tissue, limiting the damage to new growth. Halophytes have adapted to be able to flourish in sodium rich environments.[92]
Safety and precautions
Sodium forms flammable hydrogen and caustic sodium hydroxide on contact with water;[93] ingestion and contact with moisture on skin, eyes or mucous membranes can cause severe burns.[94][95] Sodium spontaneously explodes in the presence of an oxidizer such as water.[96] Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane should not be used on sodium fire.[95] Metal fires are Class D, but not all Class D extinguishers are workable with sodium. An effective extinguishing agent for sodium fires is Met-L-X.[95] Other effective agents include Lith-X, which has graphite powder and an organophosphate flame retardant, and dry sand.[97] Sodium fires are prevented in nuclear reactors by isolating sodium from oxygen by surrounding sodium pipes with inert gas.[98] Pool-type sodium fires are prevented using different design measures called catch pan systems. They collect leaking sodium into a leak-recovery tank where it is isolated from oxygen.[98]
See also
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Bibliography
- Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (edisi ke-2), Oxford: Butterworth-Heinemann, ISBN 0-7506-3365-4
External links
- Sodium at The Periodic Table of Videos (University of Nottingham)
- Etymology of "natrium" – source of symbol Na
- The Wooden Periodic Table Table's Entry on Sodium
- Sodium isotopes data from The Berkeley Laboratory Isotopes Project's
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| 1 | H | He | |||||||||||||||||||||||||||||||
| 2 | Li | Be | B | C | N | O | F | Ne | |||||||||||||||||||||||||
| 3 | Na | Mg | Al | Si | P | S | Cl | Ar | |||||||||||||||||||||||||
| 4 | K | Ca | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | |||||||||||||||
| 5 | Rb | Sr | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | |||||||||||||||
| 6 | Cs | Ba | La | Ce | Pr | Nd | Pm | Sm | Eu | Gd | Tb | Dy | Ho | Er | Tm | Yb | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn | |
| 7 | Fr | Ra | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | |
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