|
|
| Appearance |
silvery white metallic
Spectral lines of sodium |
| General properties |
| Name, symbol, number |
sodium, Na, 11 |
| Pronunciation |
/ˈsoʊdiəm/ SOH-dee-əm |
| Element category |
alkali metal |
| Group, period, block |
1, 3, s |
| Standard atomic weight |
22.98976928(2) |
| Electron configuration |
[Ne] 3s1 |
| Electrons per shell |
2,8,1 (Image) |
| Physical properties |
| Phase |
solid |
| Density (near r.t.) |
0.968 g·cm−3 |
| Liquid density at m.p. |
0.927 g·cm−3 |
| Melting point |
370.87 K, 97.72 °C, 207.9 °F |
| Boiling point |
1156 K, 883 °C, 1621 °F |
| Critical point |
(extrapolated)
2573 K, 35 MPa |
| Heat of fusion |
2.60 kJ·mol−1 |
| Heat of vaporization |
97.42 kJ·mol−1 |
| Molar heat capacity |
28.230 J·mol−1·K−1 |
| 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
(strongly basic oxide) |
| Electronegativity |
0.93 (Pauling scale) |
Ionization energies
(more) |
1st: 495.8 kJ·mol−1 |
| 2nd: 4562 kJ·mol−1 |
| 3rd: 6910.3 kJ·mol−1 |
| Atomic radius |
186 pm |
| Covalent radius |
166±9 pm |
| Van der Waals radius |
227 pm |
| Miscellanea |
| Crystal structure |
body-centered cubic |
| Magnetic ordering |
paramagnetic |
| Electrical resistivity |
(20 °C) 47.7 nΩ·m |
| Thermal conductivity |
142 W·m−1·K−1 |
| Thermal expansion |
(25 °C) 71 µm·m−1·K−1 |
| Speed of sound (thin rod) |
(20 °C) 3200 m·s−1 |
| 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 registry number |
7440-23-5 |
| Most stable isotopes |
| Main article: Isotopes of sodium |
|
|
|
· r |
Sodium (
/ˈsoʊdiəm/ SOH-dee-əm) is a chemical element with the symbol Na (from Latin: natrium) and atomic number 11. It is a soft, silvery-white, highly reactive metal and is a member of the alkali metals; its only stable isotope is 23Na. The free metal does not occur in nature, but instead must be prepared from its compounds; it was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide. Sodium is the sixth most abundant element in the Earth's crust, and exists in numerous minerals such as feldspars, sodalite and rock salt. Many salts of sodium are highly water-soluble, and their sodium has been leached by the action of water so that chloride and sodium are the most common dissolved elements by weight in the Earth's bodies of oceanic water.
Many sodium compounds are useful, such as sodium hydroxide (lye) for soapmaking, and sodium chloride for use as a deicing agent and a nutrient. Sodium is an essential element for all animals and some plants. In animals, sodium ions are used against potassium ions to build up charges on cell membranes, allowing transmission of nerve impulses when the charge is dissipated. The consequent need of animals for sodium causes it to classified as a dietary inorganic macro-mineral.
A positive
flame test for sodium has a bright yellow color.
Sodium at standard temperature and pressure is a soft metal that can be readily cut with a knife and is a good conductor of electricity. Freshly exposed, sodium has a bright, silvery luster that rapidly tarnishes, forming a white coating of sodium hydroxide and sodium carbonate. These properties change at elevated pressures: at 1.5 Mbar, the color changes to black, then to red transparent at 1.9 Mbar, and finally clear transparent at 3 Mbar. All of these allotropes are insulators and electrides.[2]
When sodium or its compounds are introduced into a flame, they turn it yellow,[3] 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.[4]
Sodium is generally less reactive than potassium and more reactive than lithium.[5] Like all the alkali metals, it reacts exothermically with water, to the point that sufficiently large pieces melt to a sphere and may explode; this reaction produces caustic sodium hydroxide and flammable hydrogen gas. When burned in dry air, it mainly forms sodium peroxide as well as some sodium oxide. In moist air, sodium hydroxide results.[6] Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts[7] but potassium and lithium have even more negative potentials.[8] Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy.[6]
20 isotopes of sodium are known, but only 23Na is stable. Two radioactive, cosmogenic isotopes are the byproduct of cosmic ray spallation: 22Na with a half-life of 2.6 years and 24Na with a half-life of 15 hours; all other isotopes have a half-life of less than one minute.[9] Two nuclear isomers have been discovered, the longer-lived one being 24mNa with a half-life of around 20.2 microseconds. Acute neutron radiation, such as from a nuclear criticality accident, converts some of the stable 23Na in human blood to 24Na; by measuring the concentration of 24Na in relation to 23Na, the neutron radiation dosage of the victim can be calculated.[10]
23Na is created in the carbon-burning process by fusing two carbon atoms together; this requires temperatures above 600 megakelvins and a star with at least three solar masses.[11] The Earth's crust has 2.6% sodium by weight, making it the sixth most abundant element on Earth.[12] 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 line; though it has a high vaporization temperature, its abundance allowed it to be detected by Mariner 10 in Mercury's atmosphere.[13]
Structure of
sodium chloride, showing octahedral coordination around Na
+ and Cl
- centres. This framework disintegrates upon dissolution in water and reassembles upon evaporation.
Sodium compounds are of immense commercial importance, being particularly central to industries producing glass, paper, soap, and textiles.[14] The sodium compounds that are the most important are common 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).[15] In its compounds, sodium is usually ionically bonded to water and anions, and is viewed as a hard Lewis acid.[16]
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.[17] 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).[18] 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.[19] Crown ethers, like 15-crown-5, may be used as a phase-transfer catalyst.[20]
Sodium content in bulk may be determined by treating with a large excess of uranyl zinc acetate; the hexahydrate (UO2)2ZnNa(CH3CO2)·6H2O precipitates, which can be weighed. Caesium and rubidium do not interfere with this reaction, but potassium and lithium do.[21] Lower concentrations of sodium may be determined by atomic absorption spectrophotometry[22] or by potentiometry using ion-selective electrodes.[23]
Like the other alkali metals, sodium dissolves in ammonia and some amines to give deeply coloured solutions; evaporation of these solutions leaves a shiny film of metallic sodium. The solutions contain the coordination complex (Na(NH3)6)+, whose positive charge is counterbalanced by electrons as anions; cryptands permit the isolation of these complexes as crystalline solids. Cryptands, like crown ethers and other ionophores, have a high affinity for the sodium ion; derivatives of the alkalide Na- are obtainable[24] by the addition of cryptands to solutions of sodium in ammonia via disproportionation.[25]
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).[26]
Salt has been an important commodity in human activities, as shown by the English word salary, which derives from salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages. In medieval Europe, a compound of sodium with the Latin name of sodanum was used as a headache remedy. The name sodium is thought to originate from the Arabic suda, meaning headache, as the headache-alleviating properties of sodium carbonate or soda were well known in early times.[27] The chemical abbreviation for sodium was first published by Jöns Jakob Berzelius in his system of atomic symbols,[28] and is a contraction of the element's new Latin name natrium, which refers to the Egyptian natron,[27] a natural mineral salt primarily made of hydrated sodium carbonate. Natron historically had several important industrial and household uses, later eclipsed by other sodium compounds. Although sodium, sometimes called soda, had long been recognised in compounds, the metal itself was not isolated until 1807 by Sir Humphry Davy through the electrolysis of sodium hydroxide.[29][30]
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:[31]
| “ |
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. |
” |
Enjoying rather specialized applications, only about 100,000 tonnes of metallic sodium are produced annually.[14] Metallic sodium was first produced commercially in 1855 by carbothermal reduction of sodium carbonate at 1100 °C, in what is known as the Deville process:[citation needed][32][33][34]
- Na2CO3 + 2 C → 2 Na + 3 CO
A related process based on the reduction of sodium hydroxide was developed in 1886.[32]
Sodium is now produced commercially through the electrolysis of molten sodium chloride, based on a process patented in 1924.[35][36] 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 formed at the anode. This method is less expensive than the previous Castner process of electrolyzing sodium hydroxide.
Reagent-grade sodium in tonne quantities sold for about US$3.30/kg in 2009; lower purity metal sells for considerably less. 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. Sodium will also burn violently when heated in air.[37] 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.[38]
Though metallic sodium has some important uses, the major applications of sodium use it in its many compounds; millions of tons of the chloride, hydroxide, and carbonate are produced annually.
Metallic sodium is mainly used for the production of sodium borohydride, sodium azide, indigo, and triphenylphosphine. Previous uses were for the making of tetraethyllead and titanium metal; because applications for these chemicals were discontinued, the production of sodium declined after 1970.[14] Sodium is also used as an alloying metal, an anti-scaling agent,[39] and as a reducing agent for metals when other materials are ineffective. Sodium vapor lamps are often used for street lighting in cities and give colours ranging from yellow-orange to peach as the pressure increases.[40] By itself or with potassium, sodium is a desiccant; it gives an intense blue colouration with benzophenone when the desiccate is dry.[41] 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.[42] Lasers emitting light at the D line, utilising sodium, are used to create artificial laser guide stars that assist in the adaptive optics for land-based visible light telescopes.[citation needed]
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.
[43]
Liquid sodium is used as a heat transfer fluid in some fast reactors,[44] due to its high thermal conductivity and low neutron absorption cross section, which is required to achieve a high neutron flux; the high boiling point allows the reactor to operate at ambient pressure. Drawbacks of using sodium include its opacity, which hinders visual maintenance, and its explosive properties. Radioactive sodium-24 may be formed by neutron activation during operation, posing a slight radiation hazard; the radioactivity stops within a few days after removal from the reactor. If a reactor needs to be frequently shut down, NaK is used; due to it being liquid at room temperature, cooling pipes do not freeze. In this case, the pyrophoricity of potassium means extra precautions against leaks need to be taken. Another heat transfer application is in high-performance internal combustion engines with poppet valves, where valve stems partially filled with sodium are used as a heat pipe to cool the valves.
Two equivalent images of the chemical structure of sodium stearate, a typical soap.
Most soaps are sodium salts of fatty acids. Sodium soaps are harder (higher melting) soaps than potassium soaps.[15] Sodium chloride is extensively used for anti-icing and de-icing and as a preservative; sodium bicarbonate is mainly used for cooking. Along with potassium, many important medicines have sodium added to improve their bioavailability; although in most cases potassium is the better ion, sodium is selected for its lower price and atomic weight.[45] 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.[46]
Sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day.[47] Sodium chloride is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods.[48] The DRI for sodium is 2.3 grams per day,[49] but on average people in the United States consume 3.4 grams per day,[50] the minimum amount that promotes hypertension;[51] this in turn causes 7.6 million premature deaths worldwide.[52]
The renin-angiotensin system regulates the amount of fluids and sodium in the body. Reduction of blood pressure and sodium concentration in the kidney result in the production of renin, which in turn produces aldosterone and angiotensin, retaining sodium in the urine. Because of the increase in sodium concentration, the production of renin decreases, and the sodium concentration returns to normal.[53] Sodium is also important in neuron function and osmoregulation between cells and the extracellular fluid, their distribution mediated in all animals by Na+/K+-ATPase;[54] hence, sodium is the most prominent cation in extracellular fluid.[55]
In C4 plants, sodium is a micronutrient that aids in metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll.[56] In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata.[57] Excess sodium in the soil limits the uptake of water due to decreased water potential, which may result in wilting; similar concentrations in the cytoplasm can lead to enzyme inhibition, which in turn causes necrosis and chlorosis.[58] To avoid these problems, plants developed mechanisms that limit sodium uptake by roots, store them in cell vacuoles, and control them over long distances;[59] excess sodium may also be stored in old plant tissue, limiting the damage to new growth.
Care is required in handling elemental sodium, as it is potentially explosive and generates flammable hydrogen and caustic sodium hydroxide upon contact with water; powdered sodium may combust spontaneously in air or oxygen.[60] 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.[61] Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane lose their effectiveness when they dissipate. An effective extinguishing agent is Met-L-X, which comprises approximately 5% Saran in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop. Other materials include Lith+, which has graphite powder and an organophosphate flame retardant, and dry sand.
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- ^ Kering, M. K. (2008). "Manganese Nutrition and Photosynthesis in NAD-malic enzyme C4 plants Ph.D. dissertation". University of Missouri-Columbia. https://mospace.umsystem.edu/xmlui/bitstream/handle/10355/7201/research.pdf?sequence=3. 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. PMID 11173290.
- ^ "Plants and salt ion toxicity". Plant Biology. http://www.plant-biology.com/salt-ion-toxicity.php. Retrieved 2010-11-02.
- ^ "Sodium Lake Explosion 1". Video.google.de. http://video.google.de/videoplay?docid=-2158222101210607510&q=sodium. Retrieved 2010-11-27.
- ^ Angelici, R. J. (1999). Synthesis and Technique in Inorganic Chemistry. Mill Valley, CA: University Science Books. ISBN 0-935702-48-2.
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Lithium
Li
Atomic Number: 3
Atomic Weight: 6.941
Melting Point: 453.85 K
Boiling Point: 1615 K
Specific mass: 0.534 g/cm3
Electronegativity: 0.98
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Sodium
Na
Atomic Number: 11
Atomic Weight: 22.98976928
Melting Point: 371.15 K
Boiling Point: 1156 K
Specific mass: 0.97 g/cm3
Electronegativity: 0.96
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Potassium
K
Atomic Number: 19
Atomic Weight: 39.0983
Melting Point: 336.5 K
Boiling Point: 1032 K
Specific mass: 0.86 g/cm3
Electronegativity: 0.82
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Rubidium
Rb
Atomic Number: 37
Atomic Weight: 85.4678
Melting Point: 312.79 K
Boiling Point: 961 K
Specific mass: 1.53 g/cm3
Electronegativity: 0.82
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Caesium
Cs
Atomic Number: 55
Atomic Weight: 132.9054519
Melting Point: 301.7 K
Boiling Point: 944 K
Specific mass: 1.93 g/cm3
Electronegativity: 0.79
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Francium
Fr
Atomic Number: 87
Atomic Weight: [223]
Melting Point: 300.15 K
Boiling Point: 950 K
Specific mass: 1.87 g/cm3
Electronegativity: 0.7
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vep:Natrii
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