Sodium

Sodium,  ₁₁Na

Spectral lines of sodium
General properties
Name, symbol	sodium, Na
Appearance	silvery white metallic
Pronunciation	/ˈsoʊdiəm/ SOH-dee-əm
Sodium in the periodic table
Li ↑ Na ↓ K neon ← sodium → magnesium
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] 3s^1
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/cm^3
when liquid, at m.p.	0.927 g/cm^3
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 (more)
Atomic radius	empirical: 186 pm
Covalent radius	166±9 pm
Van der Waals radius	227 pm
Miscellanea
Crystal structure	​body-centered cubic (bcc)
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
History
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

Sodium /ˈsoʊdiəm/^[4] is a chemical element with symbol Na (from Greek Νάτριο) and atomic number 11. It is a soft, silver-white, highly reactive metal. In the Periodic table it is in column 1 (alkali metals), and like the other six elements in that column, it has a single electron in its outer shell that it readily donates, creating a positively charged atom — a cation. Its only stable isotope is ²³Na. The free metal does not occur in nature, but 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; 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 humans and cattle.

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 Na+/K+-ATPase, 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 oxidizes to grayish white unless immersed in oil or inert gas. Sodium can be easily cut with a knife and is a good conductor of electricity and heat. These properties change dramatically at elevated pressures: at 1.5 Mbar, the color changes from silvery metallic to black; at 1.9 Mbar the material becomes transparent with a red color; and at 3 Mbar, sodium is a clear and transparent solid. All of these high-pressure allotropes are insulators and electrides.^[5]

Sodium gas is at first green, then turns purple at higher temperatures.^[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 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]

Chemical

Emission spectrum for sodium, showing the D line.

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

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

Sodium is generally less reactive than potassium and more reactive than lithium.^[9] Like all the alkali metals, it 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 dry air, it forms primarily sodium peroxide with some sodium oxide. In moist air, it forms sodium hydroxide.^[10] Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts,^[11] though potassium and lithium have even more negative potentials.^[12] Extracting sodium metal from a compound such as sodium chloride uses a significant amount of energy.^[10]

Isotopes

Twenty isotopes of sodium are known, but only ²³Na is stable. 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.^[13] Two nuclear isomers have been discovered, the longer-lived one being ^24mNa with a half-life of around 20.2 miliseconds. 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.^[14]

Occurrence

²³Na is created in the carbon-burning process in stars by fusing two carbon atoms together; this requires temperatures above 600 megakelvins and a star of at least three solar masses.^[15] The Earth's crust contains 2.6% sodium by weight, making it the sixth most abundant element on Earth.^[16] Sodium's estimated crustal abundance is 2.36×10⁴ milligrams per kilogram. Sodium's estimated oceanic abundance is 1.08×10⁴ milligrams per liter. 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.^[17]

Compounds

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.^[18] 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).^[19] In compounds, sodium is usually ionically bonded to water and anions, and is viewed as a hard Lewis acid.^[20]

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.^[19] 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.^[21] 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.^[22]

Aqueous solutions

Sodium tends to form water-soluble compounds, such as halides, sulfates, nitrates, carboxylates and carbonates. The main aqueous species are the aquo complexes [Na(H₂O)_n]⁺, where n = 4–6.^[23] The high affinity of sodium for oxygen-based ligands is the basis of crown ethers; macrolide antibiotics, which interfere with Na⁺ transport in the infecting organism, are functionally related and more complex.

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).^[24] 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.^[25] Crown ethers, like 15-crown-5, may be used as a phase-transfer catalyst.^[26]

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.^[27] Lower concentrations of sodium may be determined by atomic absorption spectrophotometry^[28] or by potentiometry using ion-selective electrodes.^[29]

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. Cryptands, like crown ethers and other ionophores, have a high affinity for the sodium ion; derivatives of the alkalide Na⁻ are obtainable^[30] by the addition of cryptands to solutions of sodium in ammonia via disproportionation.^[31]

Organosodium compounds

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

Many organosodium compounds have been prepared. Because of the high polarity of the C-Na bonds, they behave like sources of carbanions (salts with organic anions). Some well known derivatives include sodium cyclopentadienide (NaC₅H₅) and trityl sodium ((C₆H₅)₃CNa).^[32]

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.^[33] The chemical abbreviation for sodium was first published by Jöns Jakob Berzelius in his system of atomic symbols,^[34] and is a contraction of the element's New Latin name natrium, which refers to the Egyptian natron,^[33] a natural mineral salt primarily made of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds. Although sodium, sometimes called soda, had long been recognized in compounds, the metal itself was not isolated until 1807 by Sir Humphry Davy through the electrolysis of sodium hydroxide.^[35][36]

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:^[37]

In a corner of our 60 m³ room farthest away from the apparatus, we exploded 3 mg. of sodium chlorate with milk sugar while observing the nonluminous flame before the slit. After a while, it glowed a bright yellow and showed a strong sodium line that disappeared only after 10 minutes. From the weight of the sodium salt and the volume of air in the room, we easily calculate that one part by weight of air could not contain more than 1/20 millionth weight of sodium.

Commercial production

Employed only in rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.^[18] Metallic sodium was first produced commercially in 1855 by carbothermal reduction of sodium carbonate at 1100 °C, as the first step of the Deville process for the production of aluminium:^[38][39][40]

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.^[38]

Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924.^[41][42] This is done in a Downs cell in which the NaCl is mixed with calcium chloride to lower the melting point below 700 °C. As calcium is less electropositive than sodium, no calcium will be deposited at the cathode. This method is less expensive than the previous Castner process (electrolysis sodium hydroxide).

Reagent-grade sodium in tonne quantities sold for about US$3.30/kg in 2009; lower purity metal sells for considerably less. The market for sodium is volatile due to the difficulty in its storage and shipping; it must be stored under a dry inert gas atmosphere or anhydrous mineral oil to prevent the formation of a surface layer of sodium oxide or sodium superoxide. These oxides can react violently in the presence of organic materials. Smaller quantities of sodium cost far more, in the range of US$165/kg; the high cost is partially due to the expense of shipping hazardous material.^[43]

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.

Free element

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.^[18] Sodium is also used as an alloying metal, an anti-scaling agent,^[44] 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.^[45] By itself or with potassium, sodium is a desiccant; it gives an intense blue coloration with benzophenone when the desiccate is dry.^[46] 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.^[47] 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.

Heat transfer

NaK phase diagram, showing the melting point of sodium as a function of potassium concentration. NaK with 77% potassium is eutectic and has the lowest melting point of the NaK alloys at −12.6 °C.^[48]

Liquid sodium is used as a heat transfer fluid in some fast reactors^[49] because it has the high thermal conductivity and low neutron absorption cross section required to achieve a high neutron flux in the reactor. The high boiling point of sodium allows the reactor to operate at ambient (normal) pressure, but the drawbacks include its opacity, which hinders visual maintenance, and its explosive properties. 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. 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. In this case, the pyrophoricity of potassium requires extra precautions to prevent and detect leaks. 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.

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.^[50] 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.^[51] The UL standard for sodium intake is 2.3 grams per day;^[52] and exceeding that threshold can lead to hypertension,^[53] but the average person in the United States consumes 3.4 grams per day.^[54] Hypertension causes 7.6 million premature deaths worldwide each year.^[55] (Note that salt contains about 39.3% sodium^[56]—the rest being chlorine and trace chemicals; thus, 2.3g sodium is about 5.9g, or 2.7ml of salt—about half a US teaspoon.^[57])

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.^[58] Sodium is important 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.^[59] Sodium is the most prevalent metallic ion in extracellular fluid.^[60]

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.^[61]

In C4 plants, sodium is a micronutrient that aids in metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll.^[62] In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata.^[63] 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.^[64] 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;^[65] excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

Precautions

NFPA 704 "fire diamond"
1 3 2 W
The fire diamond hazard sign for sodium metal[66]

Care is required in handling elemental sodium because it generates flammable hydrogen and caustic sodium hydroxide on contact with water; powdered sodium may spontaneously explode in the presence of an oxidizer. Excess sodium can be safely removed by hydrolysis in a ventilated cabinet; this is typically done by sequential treatment with isopropanol, ethanol and water. Isopropanol reacts very slowly, generating the corresponding alkoxide and hydrogen.^[67] Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane lose their effectiveness when they dissipate.

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, which comprises approximately 5% Saran in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop. Other effective agents include Lith-X, which has graphite powder and an organophosphate flame retardant, and dry sand.

See also

Sodium in biology

Alkali metal

References

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. Wells, John C. (2008), Longman Pronunciation Dictionary (3rd ed.), Longman, p. 755, ISBN 9781405881180
5. Gatti, M.; Tokatly, I.; Rubio, A. (2010). "Sodium: A Charge-Transfer Insulator at High Pressures". Physical Review Letters 104 (21): 216404. arXiv:1003.0540. Bibcode:2010PhRvL.104u6404G. doi:10.1103/PhysRevLett.104.216404. PMID 20867123.
6. https://www.youtube.com/watch?v=g2rFc7DVZIo
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. De Leon, N. "Reactivity of Alkali Metals". Indiana University Northwest. Retrieved 2007-12-07.
10. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
11. Atkins, Peter W.; de Paula, Julio (2002). Physical Chemistry (7th ed.). W. H. Freeman. ISBN 978-0-7167-3539-7. OCLC 3345182.
12. Davies, Julian A. (1996). Synthetic Coordination Chemistry: Principles and Practice. World Scientific. p. 293. ISBN 978-981-02-2084-6. OCLC 717012347.
13. 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.
14. 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.
15. Denisenkov, P. A.; Ivanov, V. V. (1987). "Sodium Synthesis in Hydrogen Burning Stars". Soviet Astronomy Letters 13: 214. Bibcode:1987SvAL...13..214D.
16. Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
17. 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/9710022. Bibcode:1997ApJ...490L.199C. doi:10.1086/311040.
18. 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
19. 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.
20. Cowan, James A. (1997). Inorganic Biochemistry: An Introduction. Wiley-VCH. p. 7. ISBN 978-0-471-18895-7. OCLC 34515430.
21. 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.
22. Wiberg, Egon; Wiberg, Nils; Holleman, A. F. (2001). Inorganic Chemistry. Academic Press. pp. 1103–1104. ISBN 978-0-12-352651-9. OCLC 48056955.
23. 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.
24. Dean, John Aurie; Lange, Norbert Adolph (1998). Lange's Handbook of Chemistry. McGraw-Hill. ISBN 0-07-016384-7.
25. Burgess, J. (1978). Metal Ions in Solution. New York: Ellis Horwood. ISBN 0-85312-027-7.
26. 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.
27. 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.
28. 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.
29. Levy, G. B. (1981). "Determination of Sodium with Ion-Selective Electrodes". Clinical Chemistry 27 (8): 1435–1438. PMID 7273405.
30. 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.
31. Holleman, A. F.; Wiberg, E.; Wiberg, N. (2001). Inorganic Chemistry. Academic Press. ISBN 978-0-12-352651-9. OCLC 48056955.
32. Renfrow, Jr., W. B.; Hauser, C. R. (1943). "Triphenylmethylsodium". Org. Synth.; Coll. Vol. 2, p. 607
33. Newton, David E. (1999). Baker, Lawrence W., ed. Chemical Elements. ISBN 978-0-7876-2847-5. OCLC 39778687.
34. van der Krogt, Peter. "Elementymology & Elements Multidict". Retrieved 2007-06-08.
35. 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.
36. 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.
37. 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.
38. 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.
39. Oesper, R. E.; Lemay, P. (1950). "Henri Sainte-Claire Deville, 1818–1881". Chymia 3: 205–221. doi:10.2307/27757153. JSTOR 27757153.
40. Banks, Alton (1990). "Sodium". Journal of Chemical Education 67 (12): 1046. Bibcode:1990JChEd..67.1046B. doi:10.1021/ed067p1046.
41. Pauling, Linus, General Chemistry, 1970 ed., Dover Publications
42. "Los Alamos National Laboratory – Sodium". Retrieved 2007-06-08.
43. "007-Sodium Metal". Mcssl.com. Retrieved 2010-11-27.
44. Harris, Jay C. (1949). Metal cleaning: bibliographical abstracts, 1842–1951. American Society for Testing and Materials. p. 76. OCLC 1848092.
45. Lindsey, Jack L. (1997). Applied illumination engineering. Fairmont Press. pp. 112–114. ISBN 978-0-88173-212-2. OCLC 22184876.
46. 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.
47. 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.
48. 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.
49. 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
50. "Sodium" (PDF). Northwestern University. Archived from the original (PDF) on 2011-08-23. Retrieved 2011-11-21.
51. "Sodium and Potassium Quick Health Facts". health.ltgovernors.com.
52. "Dietary Reference Intakes: Water, Potassium, Sodium, Chloride, and Sulfate". Food and Nutrition Board, Institute of Medicine, United States National Academies. 11 February 2004.
53. Geleijnse, J. M.; Kok, F. J.; Grobbee, D. E. (2004). "Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations". European Journal of Public Health 14 (3): 235–239. doi:10.1093/eurpub/14.3.235. PMID 15369026.
54. 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.
55. 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.
56. Armstrong, James (2011). General, Organic, and Biochemistry: An Applied Approach. Cengage Learning. pp. 48–. ISBN 1-133-16826-4.
57. Table Salt Conversion. Traditionaloven.com. Retrieved on 2015-11-11.
58. McGuire, Michelle; Beerman, Kathy A. (2011). Nutritional Sciences: From Fundamentals to Food. Cengage Learning. p. 546. ISBN 978-0-324-59864-3. OCLC 472704484.
59. Campbell, Neil (1987). Biology. Benjamin/Cummings. p. 795. ISBN 0-8053-1840-2.
60. Srilakshmi, B. (2006). Nutrition Science (2nd ed.). New Age International. p. 318. ISBN 978-81-224-1633-6. OCLC 173807260.
61. 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.
62. 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.
63. 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.
64. 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.
65. "Plants and salt ion toxicity". Plant Biology. Retrieved 2010-11-02.
66. Hazard Rating Information for NFPA Fire Diamonds. Ehs.neu.edu. Retrieved on 2015-11-11.
67. Angelici, R. J. (1999). Synthesis and Technique in Inorganic Chemistry. Mill Valley, CA: University Science Books. ISBN 0-935702-48-2.

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