This is a good article. Click here for more information.


From Wikipedia, the free encyclopedia
Jump to: navigation, search
This article is about the chemical element. For the nutrient commonly called sodium, see salt. For other uses, see sodium (disambiguation).
"Natrium" redirects here. For other uses, see Natrium (disambiguation).
Sodium,  11Na
Na (Sodium).jpg
Sodium Spectra.jpg
Spectral lines of sodium
General properties
Name, symbol sodium, Na
Pronunciation /ˈsdiəm/
Appearance silvery white metallic
Sodium in the periodic table
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (post-transition metal)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (post-transition metal)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (post-transition metal)
Tin (post-transition metal)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (post-transition metal)
Lead (post-transition metal)
Bismuth (post-transition metal)
Polonium (post-transition metal)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Ununtrium (unknown chemical properties)
Flerovium (post-transition metal)
Ununpentium (unknown chemical properties)
Livermorium (unknown chemical properties)
Ununseptium (unknown chemical properties)
Ununoctium (unknown chemical properties)


Atomic number (Z) 11
Group, block group 1 (alkali metals), s-block
Period period 3
Element category   alkali metal
Standard atomic weight (±) (Ar) 22.98976928(2)[1]
Electron configuration [Ne] 3s1
per shell
2, 8, 1
Physical properties
Phase solid
Melting point 370.944 K ​(97.794 °C, ​208.029 °F)
Boiling point 1156.090 K ​(882.940 °C, ​1621.292 °F)
Density near r.t. 0.968 g/cm3
when liquid, at m.p. 0.927 g/cm3
Critical point 2573 K, 35 MPa (extrapolated)
Heat of fusion 2.60 kJ/mol
Heat of vaporization 97.42 kJ/mol
Molar heat capacity 28.230 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 states +1, −1 ​(a strongly basic oxide)
Electronegativity Pauling scale: 0.93
Ionization energies 1st: 495.8 kJ/mol
2nd: 4562 kJ/mol
3rd: 6910.3 kJ/mol
Atomic radius empirical: 186 pm
Covalent radius 166±9 pm
Van der Waals radius 227 pm
Crystal structure body-centered cubic (bcc)
Body-centered cubic crystal structure for sodium
Speed of sound thin rod 3200 m/s (at 20 °C)
Thermal expansion 71 µm/(m·K) (at 25 °C)
Thermal conductivity 142 W/(m·K)
Electrical resistivity 47.7 nΩ·m (at 20 °C)
Magnetic ordering paramagnetic[2]
Young's modulus 10 GPa
Shear modulus 3.3 GPa
Bulk modulus 6.3 GPa
Mohs hardness 0.5
Brinell hardness 0.69 MPa
CAS Number 7440-23-5
Discovery and first isolation Humphry Davy (1807)
Most stable isotopes of sodium
iso NA half-life DM DE (MeV) DP
22Na trace 2.602 y β+γ 0.5454 22Ne*
1.27453(2)[3] 22Ne
ε→γ 22Ne*
1.27453(2) 22Ne
β+ 1.8200 22Ne
23Na 100% 23Na is stable with 12 neutrons
* = excited state
| references

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, and hence 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.



Emission spectrum for sodium, showing the D line.

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 weight 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]

A positive flame test for sodium has a bright yellow color.

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 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.[8]


Main article: Isotopes of sodium

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]


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 reduction of sodium ions requiring −2.71 volts,[15] though potassium and lithium have even more negative potentials.[16]

Salts and oxides[edit]

Structure of sodium chloride, showing octahedral coordination around Na+ and Cl centres. This framework disintegrates when dissolved in water and reassembles when the water evaporates.

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]

Two equivalent images of the chemical structure of sodium stearate, a typical soap.

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[edit]

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[edit]

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[edit]

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).[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[edit]

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]


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.


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. 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. 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[edit]

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 in 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]


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]

Free element[edit]

Metallic sodium is used mainly 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.[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] Sodium lasers emitting light at the 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[edit]

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.[65]

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 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[edit]

Main article: Sodium in biology

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 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 in 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 limits 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 developed mechanisms to limit sodium uptake in the roots, to store it in cell vacuoles, and restriction of salt transport from roots to leaves;[92] excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

Safety and precautions[edit]

NFPA 704
"fire diamond"
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
The fire diamond hazard sign for sodium metal[93]

Sodium forms flammable hydrogen and caustic sodium hydroxide on contact with water;[94] ingestion and contact with moisture on skin, eyes or mucous membranes can cause severe burns.[95][96] Sodium spontaneously explodes in the presence of an oxidizer such as water.[97] Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane should not be used on sodium fire.[96] 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.[96] Other effective agents include Lith-X, which has graphite powder and an organophosphate flame retardant, and dry sand.[98] Sodium fires are prevented in nuclear reactors by isolating sodium from oxygen by surrounding sodium pipes with inert gas.[99] 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.[99]

See also[edit]


  1. ^ Standard Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
  2. ^ Magnetic susceptibility of the elements and inorganic compounds, in Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5. 
  3. ^ Endt, P. M. (1990). "Energy levels of A = 21–44 nuclei (VII)". Nuclear Physics A. 521: 1–400. Bibcode:1990NuPhA.521....1E. doi:10.1016/0375-9474(90)90598-G. 
  4. ^ Greenwood and Earnshaw, p. 75
  5. ^ ""Alkali Metals." Science of Everyday Things". Retrieved 15 October 2016. 
  6. ^ Gatti, M.; Tokatly, I.; Rubio, A. (2010). "Sodium: A Charge-Transfer Insulator at High Pressures". Physical Review Letters. 104 (21): 216404. arXiv:1003.0540free to read. Bibcode:2010PhRvL.104u6404G. doi:10.1103/PhysRevLett.104.216404. PMID 20867123. 
  7. ^ Schumann, Walter (5 August 2008). Minerals of the World (2nd ed.). Sterling. p. 28. ISBN 978-1-4027-5339-8. OCLC 637302667. 
  8. ^ Citron, M. L.; Gabel, C.; Stroud, C.; Stroud, C. (1977). "Experimental Study of Power Broadening in a Two-Level Atom". Physical Review A. 16 (4): 1507. Bibcode:1977PhRvA..16.1507C. doi:10.1103/PhysRevA.16.1507. 
  9. ^ Denisenkov, P. A.; Ivanov, V. V. (1987). "Sodium Synthesis in Hydrogen Burning Stars". Soviet Astronomy Letters. 13: 214. Bibcode:1987SvAL...13..214D. 
  10. ^ Audi, Georges; Bersillon, O.; Blachot, J.; Wapstra, A.H. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. 729: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001. 
  11. ^ Sanders, F. W.; Auxier, J. A. (1962). "Neutron Activation of Sodium in Anthropomorphous Phantoms". HealthPhysics. 8 (4): 371–379. doi:10.1097/00004032-196208000-00005. PMID 14496815. 
  12. ^ Sobrasua Ibim. Biology: Threads of Life. Xlibris Corporation, 2010. p. 27. ISBN 1-4535-2068-6. 
  13. ^ Lawrie Ryan; Roger Norris. Cambridge International AS and A Level Chemistry Coursebook (illustrated ed.). Cambridge University Press, 2014. p. 36. ISBN 1-107-63845-3. 
  14. ^ De Leon, N. "Reactivity of Alkali Metals". Indiana University Northwest. Retrieved 2007-12-07. 
  15. ^ Atkins, Peter W.; de Paula, Julio (2002). Physical Chemistry (7th ed.). W. H. Freeman. ISBN 978-0-7167-3539-7. OCLC 3345182. 
  16. ^ Davies, Julian A. (1996). Synthetic Coordination Chemistry: Principles and Practice. World Scientific. p. 293. ISBN 978-981-02-2084-6. OCLC 717012347. 
  17. ^ a b c 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
  18. ^ a b Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. pp. 931–943. ISBN 3-11-007511-3. 
  19. ^ Cowan, James A. (1997). Inorganic Biochemistry: An Introduction. Wiley-VCH. p. 7. ISBN 978-0-471-18895-7. OCLC 34515430. 
  20. ^ Greenwoood and Earnshaw, p. 84
  21. ^ Lincoln, S.F.; Richens, D.T.; Sykes, A.G. (2004). "Metal Aqua Ions". Comprehensive Coordination Chemistry II. p. 515. doi:10.1016/B0-08-043748-6/01055-0. ISBN 978-0-08-043748-4. 
  22. ^ Dean, John Aurie; Lange, Norbert Adolph (1998). Lange's Handbook of Chemistry. McGraw-Hill. ISBN 0-07-016384-7. 
  23. ^ Burgess, J. (1978). Metal Ions in Solution. New York: Ellis Horwood. ISBN 0-85312-027-7. 
  24. ^ Starks, Charles M.; Liotta, Charles L.; Halpern, Marc (1994). Phase-Transfer Catalysis: Fundamentals, Applications, and Industrial Perspectives. Chapman & Hall. p. 162. ISBN 978-0-412-04071-9. OCLC 28027599. 
  25. ^ Barber, H. H.; Kolthoff, I. M. (1929). "Gravimetric Determination of Sodium by the Uranyl Zinc Acetate Method. Ii. Application in the Presence of Rubidium, Cesium, Potassium, Lithium, Phosphate or Arsenate". J. Am. Chem. Soc. 51 (11): 3233. doi:10.1021/ja01386a008. 
  26. ^ Kingsley, G. R.; Schaffert, R. R. (1954). "Micro-flame Photometric Determination of Sodium, Potassium and Calcium in Serum with Solvents". J. Biol. Chem. 206 (2): 807–15. PMID 13143043. 
  27. ^ Levy, G. B. (1981). "Determination of Sodium with Ion-Selective Electrodes". Clinical Chemistry. 27 (8): 1435–1438. PMID 7273405. 
  28. ^ Ivor L. Simmons (ed.). Applications of the Newer Techniques of Analysis. Springer Science & Business Media, 2012. p. 160. ISBN 1-4684-3318-0. 
  29. ^ Xu Hou (ed.). Design, Fabrication, Properties and Applications of Smart and Advanced Materials (illustrated ed.). CRC Press, 2016. p. 175. ISBN 1-4987-2249-0. 
  30. ^ Nikos Hadjichristidis; Akira Hirao (eds.). Anionic Polymerization: Principles, Practice, Strength, Consequences and Applications (illustrated ed.). Springer, 2015. p. 349. ISBN 4-431-54186-1. 
  31. ^ Dye, J. L.; Ceraso, J. M.; Mei Lok Tak; Barnett, B. L.; Tehan, F. J. (1974). "Crystalline Salt of the Sodium Anion (Na)". J. Am. Chem. Soc. 96 (2): 608–609. doi:10.1021/ja00809a060. 
  32. ^ Holleman, A. F.; Wiberg, E.; Wiberg, N. (2001). Inorganic Chemistry. Academic Press. ISBN 978-0-12-352651-9. OCLC 48056955. 
  33. ^ Renfrow, Jr., W. B.; Hauser, C. R. (1943). "Triphenylmethylsodium". Org. Synth. ; Coll. Vol., 2, p. 607 
  34. ^ Greenwood and Earnshaw, p. 111
  35. ^ Habashi, Fathi. Alloys: Preparation, Properties, Applications. John Wiley & Sons, 2008. pp. 278–280. ISBN 3-527-61192-4. 
  36. ^ a b Newton, David E. (1999). Baker, Lawrence W., ed. Chemical Elements. ISBN 978-0-7876-2847-5. OCLC 39778687. 
  37. ^ Davy, Humphry (1808). "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". Philosophical Transactions of the Royal Society of London. 98: 1–44. doi:10.1098/rstl.1808.0001. 
  38. ^ Weeks, Mary Elvira (1932). "The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium". Journal of Chemical Education. 9 (6): 1035. Bibcode:1932JChEd...9.1035W. doi:10.1021/ed009p1035. 
  39. ^ 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 ; 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.)
  40. ^ 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), p. 87.
  41. ^ van der Krogt, Peter. "Elementymology & Elements Multidict". Retrieved 2007-06-08. 
  42. ^ "Natron as a flux in the early vitreous materials industry: sources, beginnings and reasons for decline". Andrew Shortland, Lukas Schachner, Ian Freestone, and Michael Tite. 
  43. ^ Kirchhoff, G.; Bunsen, R. (1860). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie. 186 (6): 161–189. Bibcode:1860AnP...186..161K. doi:10.1002/andp.18601860602. 
  44. ^ Greenwood and Earnshaw, p. 69
  45. ^ Lide, David R. (2003-06-19). CRC Handbook of Chemistry and Physics, 84th Edition. CRC Handbook. CRC Press. 14: Abundance of Elements in the Earth's Crust and in the Sea. ISBN 978-0-8493-0484-2. 
  46. ^ Cremonese, G; Boehnhardt, H; Crovisier, J; Rauer, H; Fitzsimmons, A; Fulle, M; Licandro, J; Pollacco, D; et al. (1997). "Neutral Sodium from Comet Hale–Bopp: A Third Type of Tail". The Astrophysical Journal Letters. 490 (2): L199–L202. arXiv:astro-ph/9710022free to read. Bibcode:1997ApJ...490L.199C. doi:10.1086/311040. 
  47. ^ B. Pearson (ed.). Speciality Chemicals: Innovations in industrial synthesis and applications (illustrated ed.). Springer Science & Business Media, 1991. p. 260. ISBN 1-85166-646-X. 
  48. ^ a b Eggeman, Tim; Updated By Staff (2007). "Sodium and Sodium Alloys". Kirk-Othmer Encyclopedia of Chemical Technology. John Wiley & Sons. doi:10.1002/0471238961.1915040912051311.a01.pub3. ISBN 0-471-23896-1. 
  49. ^ Oesper, R. E.; Lemay, P. (1950). "Henri Sainte-Claire Deville, 1818–1881". Chymia. 3: 205–221. doi:10.2307/27757153. JSTOR 27757153. 
  50. ^ Banks, Alton (1990). "Sodium". Journal of Chemical Education. 67 (12): 1046. Bibcode:1990JChEd..67.1046B. doi:10.1021/ed067p1046. 
  51. ^ Pauling, Linus, General Chemistry, 1970 ed., Dover Publications
  52. ^ "Los Alamos National Laboratory – Sodium". Retrieved 2007-06-08. 
  53. ^ Sodium Metal from France. DIANE Publishing. ISBN 1-4578-1780-2. 
  54. ^ Mark Anthony Benvenuto. Industrial Chemistry: For Advanced Students (illustrated ed.). Walter de Gruyter GmbH & Co KG, 2015. ISBN 3-11-038339-X. 
  55. ^ Stanley Nusim (ed.). Active Pharmaceutical Ingredients: Development, Manufacturing, and Regulation, Second Edition (2, illustrated, revised ed.). CRC Press, 2016. p. 303. ISBN 1-4398-0339-0. 
  56. ^ Remington, Joseph P. (2006). Beringer, Paul, ed. Remington: The Science and Practice of Pharmacy (21st ed.). Lippincott Williams & Wilkins. pp. 365–366. ISBN 978-0-7817-4673-1. OCLC 60679584. 
  57. ^ Wiberg, Egon; Wiberg, Nils; Holleman, A. F. (2001). Inorganic Chemistry. Academic Press. pp. 1103–1104. ISBN 978-0-12-352651-9. OCLC 48056955. 
  58. ^ Harris, Jay C. (1949). Metal cleaning: bibliographical abstracts, 1842–1951. American Society for Testing and Materials. p. 76. OCLC 1848092. 
  59. ^ Lindsey, Jack L. (1997). Applied illumination engineering. Fairmont Press. pp. 112–114. ISBN 978-0-88173-212-2. OCLC 22184876. 
  60. ^ Lerner, Leonid (2011-02-16). Small-Scale Synthesis of Laboratory Reagents with Reaction Modeling. CRC Press. pp. 91–92. ISBN 978-1-4398-1312-6. OCLC 669160695. 
  61. ^ Sethi, Arun (1 January 2006). Systematic Laboratory Experiments in Organic Chemistry. New Age International. pp. 32–35. ISBN 978-81-224-1491-2. OCLC 86068991. 
  62. ^ Smith, Michael. Organic Synthesis (3 ed.). Academic Press, 2011. p. 455. ISBN 0-12-415884-6. 
  63. ^ Solomons & Fryhle. Organic Chemistry (8 ed.). John Wiley & Sons, 2006. p. 272. ISBN 81-265-1050-1. 
  64. ^ "Laser Development for Sodium Laser Guide Stars at ESO" (PDF). Domenico Bonaccini Calia, Yan Feng, Wolfgang Hackenberg, Ronald Holzlöhner, Luke Taylor, Steffan Lewis. 
  65. ^ van Rossen, G. L. C. M.; van Bleiswijk, H. (1912). "Über das Zustandsdiagramm der Kalium-Natriumlegierungen". Zeitschrift für anorganische Chemie. 74: 152–156. doi:10.1002/zaac.19120740115. 
  66. ^ 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
  67. ^ a b "Sodium-cooled Fast Reactor (SFR)" (PDF). Office of Nuclear Energy, U.S. Department of Energy. 18 February 2015. 
  68. ^ Fire and Explosion Hazards. Research Publishing Service, 2011. p. 363. ISBN 981-08-7724-2. 
  69. ^ Pavel Solomonovich Knopov, Panos M. Pardalos (eds.). Simulation and Optimization Methods in Risk and Reliability Theory. Nova Science Publishers, 2009. p. 150. ISBN 1-60456-658-2. 
  70. ^ McKillop, Allan A. Proceedings of the Heat Transfer and Fluid Mechanics Institute. Stanford University Press, 1976. p. 97. ISBN 0-8047-0917-3. 
  71. ^ U.S. Atomic Energy Commission. Reactor Handbook: Engineering (2 ed.). Interscience Publishers. p. 325. 
  72. ^ A US US2949907 A, Tauschek Max J, "Coolant-filled poppet valve and method of making same", published 23 Aug 1960 
  73. ^ "Sodium" (PDF). Northwestern University. Archived from the original (PDF) on 2011-08-23. Retrieved 2011-11-21. 
  74. ^ "Sodium and Potassium Quick Health Facts". 
  75. ^ "Sodium in diet". MedlinePlus, US National Library of Medicine. 5 October 2016. 
  76. ^ "Reference Values for Elements". Dietary Reference Intakes Tables. Health Canada. 
  77. ^ U.S. Department of Agriculture; U.S. Department of Health and Human Services (December 2010). Dietary Guidelines for Americans, 2010 (PDF) (7th ed.). p. 22. ISBN 978-0-16-087941-8. OCLC 738512922. Archived from the original (PDF) on 6 February 2011. Retrieved 2011-11-23. 
  78. ^ a b Geleijnse, J. M.; Kok, F. J.; Grobbee, D. E. (2004). "Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations" (PDF). European Journal of Public Health. 14 (3): 235–239. doi:10.1093/eurpub/14.3.235. PMID 15369026. 
  79. ^ Lawes, C. M.; Vander Hoorn, S.; Rodgers, A.; International Society of Hypertension (2008). "Global burden of blood-pressure-related disease, 2001". Lancet. 371 (9623): 1513–1518. doi:10.1016/S0140-6736(08)60655-8. PMID 18456100. 
  80. ^ Armstrong, James (2011). General, Organic, and Biochemistry: An Applied Approach. Cengage Learning. pp. 48–. ISBN 1-133-16826-4. 
  81. ^ Table Salt Conversion. Retrieved on 2015-11-11.
  82. ^ a b "Sodium in Your Diet: Use the Nutrition Facts Label and Reduce Your Intake". US Food and Drug Administration. 2 June 2016. Retrieved 15 October 2016. 
  83. ^ "How much sodium should I eat per day?". American Heart Association. 2016. Retrieved 15 October 2016. 
  84. ^ Andrew Mente; et al. (2016). "Associations of urinary sodium excretion with cardiovascular events in individuals with and without hypertension: a pooled analysis of data from four studies". The Lancet. doi:10.1016/S0140-6736(16)30467-6. 
  85. ^ McGuire, Michelle; Beerman, Kathy A. (2011). Nutritional Sciences: From Fundamentals to Food. Cengage Learning. p. 546. ISBN 978-0-324-59864-3. OCLC 472704484. 
  86. ^ Campbell, Neil (1987). Biology. Benjamin/Cummings. p. 795. ISBN 0-8053-1840-2. 
  87. ^ Srilakshmi, B. (2006). Nutrition Science (2nd ed.). New Age International. p. 318. ISBN 978-81-224-1633-6. OCLC 173807260. 
  88. ^ Pohl, Hanna R.; Wheeler, John S.; Murray, H. Edward (2013). Astrid Sigel; Helmut Sigel; Roland K. O. Sigel, eds. Interrelations between Essential Metal Ions and Human Diseases. Metal Ions in Life Sciences. 13. Springer. pp. 29–47. doi:10.1007/978-94-007-7500-8_2. 
  89. ^ Kering, M. K. (2008). "Manganese Nutrition and Photosynthesis in NAD-malic enzyme C4 plants Ph.D. dissertation" (PDF). University of Missouri-Columbia. Retrieved 2011-11-09. 
  90. ^ Subbarao, G. V.; Ito, O.; Berry, W. L.; Wheeler, R. M. (2003). "Sodium—A Functional Plant Nutrient". Critical Reviews in Plant Sciences. 22 (5): 391–416. doi:10.1080/07352680390243495. 
  91. ^ Zhu, J. K. (2001). "Plant salt tolerance". Trends in Plant Science. 6 (2): 66–71. doi:10.1016/S1360-1385(00)01838-0. PMID 11173290. 
  92. ^ "Plants and salt ion toxicity". Plant Biology. Retrieved 2010-11-02. 
  93. ^ Hazard Rating Information for NFPA Fire Diamonds. Retrieved on 2015-11-11.
  94. ^ Angelici, R. J. (1999). Synthesis and Technique in Inorganic Chemistry. Mill Valley, CA: University Science Books. ISBN 0-935702-48-2. 
  95. ^ Routley, J. Gordon. Sodium Explosion Critically Burns Firefighters: Newton, Massachusetts. U. S. Fire Administration. FEMA, 2013. 
  96. ^ a b c Prudent Practices in the Laboratory: Handling and Disposal of Chemicals. National Research Council (U.S.). Committee on Prudent Practices for Handling, Storage, and Disposal of Chemicals in Laboratories. National Academies, 1995. p. 390. 
  97. ^ "Sodium and Salt". Retrieved 2016-09-05. 
  98. ^ Ladwig, Thomas H. Industrial fire prevention and protection. Van Nostrand Reinhold, 1991. p. 178. ISBN 0-442-23678-6. 
  99. ^ a b Günter Kessler. Sustainable and Safe Nuclear Fission Energy: Technology and Safety of Fast and Thermal Nuclear Reactors (illustrated ed.). Springer Science & Business Media, 2012. p. 446. ISBN 3642119905. 


External links[edit]