Sodium is a chemical element with symbol Na (from Latin natrium)
and atomic number 11. It is a soft, silvery-white, highly
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 ion—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
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.
2.1 Salts and oxides
2.2 Aqueous solutions
2.3 Electrides and sodides
2.4 Organosodium compounds
2.5 Intermetallic compounds
4.1 Astronomical observations
5 Commercial production
6.1 Heat transfer
7 Biological role
8 Safety and precautions
9 See also
12 External links
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 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. 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.
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.
A positive flame test for sodium has a bright yellow color.
In a flame test, sodium and its compounds glow yellow 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.
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. 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. 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
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. 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.
The most common oxidation state for sodium is +1. It is generally less
reactive than potassium and more reactive than lithium. Sodium
metal is highly reducing, with the standard reduction potential for
the Na+/Na couple being −2.71 volts, though potassium and
lithium have even more negative potentials.
Salts and oxides
See also: Category:
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. 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). In
compounds, sodium is usually ionically bonded to water and anions, and
is viewed as a hard Lewis acid.
Two equivalent images of the chemical structure of sodium stearate, a
Most soaps are sodium salts of fatty acids.
Sodium soaps have a higher
melting temperature (and seem "harder") than potassium soaps.
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.
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.
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). 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.
Crown ethers, like 15-crown-5, may be used as a phase-transfer
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.
rubidium do not interfere with this reaction, but potassium and
lithium do. Lower concentrations of sodium may be determined by
atomic absorption spectrophotometry or by potentiometry using
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. For example, 15-crown-5
has high affinity for sodium because the cavity size of
1.7–2.2 Å, which is enough to fit sodium ion
(1.9 Å). Cryptands, like crown ethers and other
ionophores, also have a high affinity for the sodium ion; derivatives
of the alkalide Na− are obtainable by the addition of cryptands
to solutions of sodium in ammonia via disproportionation.
The structure of the complex of sodium (Na+, shown in yellow) and the
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).
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
Sodium forms alloys with many metals, such as potassium, calcium,
lead, and the group 11 and 12 elements.
Sodium and potassium form KNa2
NaK is 40–90% potassium and it is liquid at ambient
temperature. It is an 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
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 electrolytically 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.
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. 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. In 1809, the German
physicist and chemist
Ludwig Wilhelm Gilbert
Ludwig Wilhelm Gilbert proposed the names
Natronium for Humphry Davy's "sodium" and Kalium for Davy's
"potassium". The chemical abbreviation for sodium was first
published in 1814 by
Jöns Jakob Berzelius
Jöns Jakob Berzelius in his system of atomic
symbols, and is an abbreviation of the element's New Latin
name natrium, which refers to the Egyptian natron, 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.
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:
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
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.
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.
Atomic sodium has a very strong spectral line in the yellow-orange
part of the spectrum (the same line as is used in sodium vapour street
lights). This appears as an absorption line in many types of stars,
including the Sun. The line was first studied in 1814 by Joseph von
Fraunhofer during his investigation of the lines in the solar
spectrum, now known as the Fraunhofer lines. Fraunhofer named it the
'D line', although it is now known to actually be a group of closely
spaced lines split by fine and hyperfine structure.
The strength of the D line means it has been detected in many other
astronomical environments. In stars, it is seen in any whose surfaces
are cool enough for sodium to exist in atomic form (rather than
ionised). This corresponds to stars of roughly F-type and cooler. Many
other stars appear to have a sodium absorption line, but this is
actually caused by gas in the foreground interstellar medium. The two
can be distinguished via high-resolution spectroscopy, because
interstellar lines are much narrower than those broadened by stellar
Sodium has also been detected in numerous
Solar System environments,
including Mercury's atmosphere, the exosphere of the
Moon, and numerous other bodies. Some comets have a sodium
tail, which was first detected in observations of
Sodium has even been detected in the atmospheres of some
extrasolar planets via transit spectroscopy.
Employed only in rather specialized applications, only about 100,000
tonnes of metallic sodium are produced annually. Metallic sodium
was first produced commercially in the late 19th century by
carbothermal reduction of sodium carbonate at 1100 °C, as the
first step of the
Deville process for the production of
Na2CO3 + 2 C → 2 Na + 3 CO
The high demand of aluminium created the need for the production of
sodium. 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.
Sodium is now produced commercially through the electrolysis of molten
sodium chloride, based on a process patented in 1924. 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 (the electrolysis of sodium hydroxide).
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.
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;
examples of the uses of sodium bicarbonate include baking, as a
raising agent, and sodablasting. 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.
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.
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.
Sodium is also used
as an alloying metal, an anti-scaling agent, 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. By itself or
with potassium, sodium is a desiccant; it gives an intense blue
coloration with benzophenone when the desiccate is dry. In organic
synthesis, sodium is used in various reactions such as the Birch
reduction, and the sodium fusion test is conducted to qualitatively
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. 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
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.
Liquid sodium is used as a heat transfer fluid in some types of
nuclear reactors 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
NaK is used; because
NaK is a liquid at room temperature,
the coolant does 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
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.
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. 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. The US
Institute of Medicine
Institute of Medicine set its Tolerable Upper
Intake Level for sodium at 2.3 grams per day, but the average
person in the United States consumes 3.4 grams per day.
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. It has been estimated that such a decrease in sodium
intake would lead to between 9 and 17% fewer cases of
Hypertension causes 7.6 million premature deaths worldwide each
year. (Note that salt contains about 39.3% sodium—the rest
being chlorine and trace chemicals; thus, 2.3 g sodium is about
5.9 g, or 2.7 ml of salt—about one US teaspoon.) The
American Heart Association
American Heart Association recommends no more than 1.5 g of
sodium per day.
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. The US FDA
states that adults with hypertension and prehypertension should reduce
daily intake to 1.5 g.
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. 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
Sodium is the most prevalent metallic ion in
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
In C4 plants, sodium is a micronutrient that aids metabolism,
specifically in regeneration of phosphoenolpyruvate and synthesis of
chlorophyll. In others, it substitutes for potassium in several
roles, such as maintaining turgor pressure and aiding in the opening
and closing of stomata. 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. 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; 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.
Safety and precautions
The fire diamond hazard sign for sodium metal
Sodium forms flammable hydrogen and caustic sodium hydroxide on
contact with water; ingestion and contact with moisture on skin,
eyes or mucous membranes can cause severe burns. Sodium
spontaneously explodes in the presence of water due to the formation
of hydrogen (highly explosive) and sodium hydroxide (which dissolves
in the water, liberating more surface). However, sodium exposed to air
and ignited or reaching autoignition (reported to occur when a molten
pool of sodium reaches about 290 °C) displays a relatively
mild fire. In the case of massive (non-molten) pieces of sodium the
reaction with oxygen eventually becomes slow due to formation of a
protective layer. Fire extinguishers based on water accelerate
sodium fires; those based on carbon dioxide and
bromochlorodifluoromethane should not be used on sodium fire.
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. Other effective agents include Lith-X, which
has graphite powder and an organophosphate flame retardant, and dry
Sodium fires are prevented in nuclear reactors by isolating
sodium from oxygen by surrounding sodium pipes with inert gas.
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.
View or order collections of articles
Period 3 elements
Chemical elements (sorted alphabetically)
Chemical elements (sorted by number)
Access related topics
Find out more on's
^ Meija, J.; et al. (2016). "Atomic weights of the elements 2013
(IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3):
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.
^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca
Raton, Florida: Chemical Rubber Company Publishing. pp. E110.
^ Greenwood and Earnshaw, p. 75
^ ""Alkali Metals." Science of Everyday Things". Encyclopedia.com.
Retrieved 15 October 2016.
^ 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.
^ Schumann, Walter (5 August 2008).
Minerals of the World (2nd ed.).
Sterling. p. 28. ISBN 978-1-4027-5339-8.
^ 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): 1507–1512. Bibcode:1977PhRvA..16.1507C.
^ Denisenkov, P. A.; Ivanov, V. V. (1987). "
Sodium Synthesis in
Hydrogen Burning Stars". Soviet Astronomy Letters. 13: 214.
^ 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.
^ 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.
^ 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.
^ De Leon, N. "Reactivity of Alkali Metals". Indiana University
Northwest. Retrieved 2007-12-07.
^ Atkins, Peter W.; de Paula, Julio (2002). Physical Chemistry (7th
ed.). W. H. Freeman. ISBN 978-0-7167-3539-7.
^ Davies, Julian A. (1996). Synthetic Coordination Chemistry:
Principles and Practice. World Scientific. p. 293.
ISBN 978-981-02-2084-6. OCLC 717012347.
^ a b c Alfred Klemm, Gabriele Hartmann, Ludwig Lange, "
Sodium Alloys" in Ullmann's Encyclopedia of Industrial Chemistry 2005,
Wiley-VCH, Weinheim. doi:10.1002/14356007.a24_277
^ 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.
^ Cowan, James A. (1997). Inorganic Biochemistry: An Introduction.
Wiley-VCH. p. 7. ISBN 978-0-471-18895-7.
^ Greenwoood and Earnshaw, p. 84
^ Lincoln, S. F.; Richens, D. T.; Sykes, A. G. (2004). "
Ions". Comprehensive Coordination Chemistry II. p. 515.
doi:10.1016/B0-08-043748-6/01055-0. ISBN 978-0-08-043748-4.
^ Dean, John Aurie; Lange, Norbert Adolph (1998). Lange's Handbook of
Chemistry. McGraw-Hill. ISBN 0-07-016384-7.
^ Burgess, J. (1978).
Metal Ions in Solution. New York: Ellis Horwood.
^ 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.
^ 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
J. Am. Chem. Soc. 51 (11): 3233–3237.
^ Kingsley, G. R.; Schaffert, R. R. (1954). "Micro-flame Photometric
Determination of Sodium,
Calcium in Serum with
Solvents". J. Biol. Chem. 206 (2): 807–15. PMID 13143043.
^ Levy, G. B. (1981). "Determination of
Sodium with Ion-Selective
Electrodes". Clinical Chemistry. 27 (8): 1435–1438.
^ Ivor L. Simmons (ed.). Applications of the Newer Techniques of
Analysis. Springer Science & Business Media, 2012. p. 160.
^ 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.
^ Nikos Hadjichristidis; Akira Hirao (eds.). Anionic Polymerization:
Principles, Practice, Strength, Consequences and Applications
(illustrated ed.). Springer, 2015. p. 349.
^ Dye, J. L.; Ceraso, J. M.; Mei Lok Tak; Barnett, B. L.; Tehan, F. J.
Salt of the
Anion (Na−)". J. Am. Chem.
Soc. 96 (2): 608–609. doi:10.1021/ja00809a060.
^ Holleman, A. F.; Wiberg, E.; Wiberg, N. (2001). Inorganic Chemistry.
Academic Press. ISBN 978-0-12-352651-9. OCLC 48056955.
^ Renfrow, Jr., W. B.; Hauser, C. R. (1943). "Triphenylmethylsodium".
Organic Syntheses. ; Collective Volume, 2, p. 607
^ Greenwood and Earnshaw, p. 111
^ Habashi, Fathi. Alloys: Preparation, Properties, Applications. John
Wiley & Sons, 2008. pp. 278–280.
^ a b Newton, David E. (1999). Baker, Lawrence W., ed. Chemical
Elements. ISBN 978-0-7876-2847-5. OCLC 39778687.
^ 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.
^ 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.
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-
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.)
^ 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.
^ van der Krogt, Peter. "Elementymology & Elements Multidict".
^ Andrew Shortland, Lukas Schachner, Ian Freestone, and Michael Tite
Natron as a flux in the early vitreous materials industry:
sources, beginnings and reasons for decline". Journal of
Archaeological Science. 33 (4): 521–530.
doi:10.1016/j.jas.2005.09.011. CS1 maint: Multiple names: authors
^ Kirchhoff, G.; Bunsen, R. (1860). "Chemische Analyse durch
Annalen der Physik und Chemie. 186 (6):
^ Greenwood and Earnshaw, p. 69
^ 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.
^ "D-lines spectroscopy". Encyclopedia Britannica. Retrieved 6
^ Welty, Daniel E; Hobbs, L. M; Kulkarni, Varsha P (1994). "A
high-resolution survey of interstellar NA I D1 lines". The
Astrophysical Journal. 436: 152. Bibcode:1994ApJ...436..152W.
^ Colaprete, A; Sarantos, M; Wooden, D. H; Stubbs, T. J; Cook, A. M;
Shirley, M (2015). "How surface composition and meteoroid impacts
mediate sodium and potassium in the lunar exosphere". Science. 351
(6270): 249. Bibcode:2016Sci...351..249C. doi:10.1126/science.aad2380.
^ "Cometary Neutral Tail COSMOS". astronomy.swin.edu.au. Retrieved 6
^ Cremonese, G; Boehnhardt, H; Crovisier, J; Rauer, H; Fitzsimmons, A;
Fulle, M; Licandro, J; Pollacco, D; et al. (1997). "Neutral Sodium
Comet Hale–Bopp: A Third Type of Tail". The Astrophysical
Journal Letters. 490 (2): L199–L202. arXiv:astro-ph/9710022 .
^ Redfield, Seth; Endl, Michael; Cochran, William D; Koesterke, Lars
Sodium Absorption from the Exoplanetary Atmosphere of HD
189733b Detected in the Optical Transmission Spectrum". The
Astrophysical Journal. 673: L87. Bibcode:2008ApJ...673L..87R.
^ 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.
^ 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.
^ Oesper, R. E.; Lemay, P. (1950). "Henri Sainte-Claire Deville,
1818–1881". Chymia. 3: 205–221. doi:10.2307/27757153.
^ Banks, Alton (1990). "Sodium". Journal of Chemical Education. 67
(12): 1046. Bibcode:1990JChEd..67.1046B. doi:10.1021/ed067p1046.
^ Pauling, Linus, General Chemistry, 1970 ed., Dover Publications
^ "Los Alamos National Laboratory – Sodium". Retrieved
Metal from France. DIANE Publishing.
^ Mark Anthony Benvenuto. Industrial Chemistry: For Advanced Students
(illustrated ed.). Walter de Gruyter GmbH & Co KG, 2015.
^ 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.
^ 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.
^ Wiberg, Egon; Wiberg, Nils; Holleman, A. F. (2001). Inorganic
Chemistry. Academic Press. pp. 1103–1104.
ISBN 978-0-12-352651-9. OCLC 48056955.
^ Harris, Jay C. (1949).
Metal cleaning: bibliographical abstracts,
1842–1951. American Society for Testing and Materials. p. 76.
^ Lindsey, Jack L. (1997). Applied illumination engineering. Fairmont
Press. pp. 112–114. ISBN 978-0-88173-212-2.
^ 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.
^ 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.
^ Smith, Michael. Organic Synthesis (3 ed.). Academic Press, 2011.
p. 455. ISBN 0-12-415884-6.
^ Solomons & Fryhle. Organic Chemistry (8 ed.). John Wiley &
Sons, 2006. p. 272. ISBN 81-265-1050-1. CS1 maint: Uses
authors parameter (link)
^ "Laser Development for
Sodium Laser Guide Stars at ESO" (PDF).
Domenico Bonaccini Calia, Yan Feng, Wolfgang Hackenberg, Ronald
Holzlöhner, Luke Taylor, Steffan Lewis.
^ van Rossen, G. L. C. M.; van Bleiswijk, H. (1912). "Über das
Zustandsdiagramm der Kalium-Natriumlegierungen". Zeitschrift für
anorganische Chemie. 74: 152–156.
Sodium as a Fast Reactor
Coolant Archived 13 January 2013 at the
Wayback Machine. 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,
^ a b "Sodium-cooled Fast Reactor (SFR)" (PDF). Office of Nuclear
Energy, U.S. Department of Energy. 18 February 2015.
^ Fire and Explosion Hazards. Research Publishing Service, 2011.
p. 363. ISBN 981-08-7724-2.
^ 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. CS1
maint: Uses editors parameter (link)
^ McKillop, Allan A. Proceedings of the Heat Transfer and Fluid
Mechanics Institute. Stanford University Press, 1976. p. 97.
^ U.S. Atomic Energy Commission. Reactor Handbook: Engineering (2
ed.). Interscience Publishers. p. 325.
^ A US US2949907 A, Tauschek Max J, "Coolant-filled poppet valve and
method of making same", published 23 Aug 1960
^ "Sodium" (PDF). Northwestern University. Archived from the original
(PDF) on 2011-08-23. Retrieved 2011-11-21.
Potassium Quick Health Facts".
Sodium in diet". MedlinePlus, US National Library of Medicine. 5
^ "Reference Values for Elements". Dietary Reference Intakes Tables.
^ 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.
^ 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.
^ Lawes, C. M.; Vander Hoorn, S.; Rodgers, A.; International Society
Hypertension (2008). "Global burden of blood-pressure-related
disease, 2001" (PDF). Lancet. 371 (9623): 1513–1518.
doi:10.1016/S0140-6736(08)60655-8. PMID 18456100.
^ Armstrong, James (2011). General, Organic, and Biochemistry: An
Applied Approach. Cengage Learning. pp. 48–.
Salt Conversion. Traditionaloven.com. Retrieved on 2015-11-11.
^ a b "Use the Nutrition Facts Label to Reduce Your Intake of Sodium
in Your Diet". US Food and Drug Administration. 3 January 2018.
Retrieved 2 February 2018.
^ "How much sodium should I eat per day?". American Heart Association.
2016. Retrieved 15 October 2016.
^ 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. 388 (10043): 465–75. doi:10.1016/S0140-6736(16)30467-6.
PMID 27216139. CS1 maint: Explicit use of et al. (link)
^ McGuire, Michelle; Beerman, Kathy A. (2011). Nutritional Sciences:
From Fundamentals to Food. Cengage Learning. p. 546.
ISBN 978-0-324-59864-3. OCLC 472704484.
^ Campbell, Neil (1987). Biology. Benjamin/Cummings. p. 795.
^ Srilakshmi, B. (2006). Nutrition Science (2nd ed.). New Age
International. p. 318. ISBN 978-81-224-1633-6.
^ Pohl, Hanna R.; Wheeler, John S.; Murray, H. Edward (2013). Astrid
Sigel; Helmut Sigel; Roland K. O. Sigel, eds. Interrelations between
Metal Ions and Human Diseases.
Metal Ions in Life Sciences.
13. Springer. pp. 29–47. doi:10.1007/978-94-007-7500-8_2.
^ Kering, M. K. (2008). "
Manganese Nutrition and Photosynthesis in
C4 plants Ph.D. dissertation" (PDF). University of
Missouri-Columbia. Retrieved 2011-11-09.
^ 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.
^ Zhu, J. K. (2001). "Plant salt tolerance". Trends in Plant Science.
6 (2): 66–71. doi:10.1016/S1360-1385(00)01838-0.
^ a b "Plants and salt ion toxicity". Plant Biology. Retrieved
^ Hazard Rating Information for NFPA Fire Diamonds Archived 17
February 2015 at the Wayback Machine.. Ehs.neu.edu. Retrieved on
^ Angelici, R. J. (1999). Synthesis and Technique in Inorganic
Chemistry. Mill Valley, CA: University Science Books.
^ Routley, J. Gordon.
Sodium Explosion Critically Burns Firefighters:
Newton, Massachusetts. U. S. Fire Administration. FEMA, 2013.
^ 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.
^ Ladwig, Thomas H. Industrial fire prevention and protection. Van
Nostrand Reinhold, 1991. p. 178. ISBN 0-442-23678-6.
^ 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 3-642-11990-5.
Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements
(2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
The Periodic Table of Videos
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
Periodic table (Large cells)
Alkaline earth metal
BNF: cb11980515x (d