Spectral lines of sodium
|Name, symbol||sodium, Na|
|Appearance||silvery white metallic|
|Sodium in the periodic table|
|Atomic number (Z)||11|
|Group, block||group 1 (alkali metals), s-block|
|Element category||alkali metal|
|Standard atomic weight (±) (Ar)||22.98976928(2)|
|Electron configuration||[Ne] 3s1|
|2, 8, 1|
|Melting point||370.944 K (97.794 °C, 208.029 °F)|
|Boiling point||1156.090 K (882.940 °C, 1621.292 °F)|
|Density near r.t.||0.968 g/cm3|
|when liquid, at m.p.||0.927 g/cm3|
|Critical point||2573 K, 35 MPa (extrapolated)|
|Heat of fusion||2.60 kJ/mol|
|Heat of vaporization||97.42 kJ/mol|
|Molar heat capacity||28.230 J/(mol·K)|
|Oxidation states||+1, −1 (a strongly basic oxide)|
|Electronegativity||Pauling scale: 0.93|
|Ionization energies||1st: 495.8 kJ/mol
2nd: 4562 kJ/mol
3rd: 6910.3 kJ/mol
|Atomic radius||empirical: 186 pm|
|Covalent radius||166±9 pm|
|Van der Waals radius||227 pm|
|Crystal structure||body-centered cubic (bcc)|
|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)|
|Young's modulus||10 GPa|
|Shear modulus||3.3 GPa|
|Bulk modulus||6.3 GPa|
|Brinell hardness||0.69 MPa|
|Discovery and first isolation||Humphry Davy (1807)|
|Most stable isotopes of sodium|
|* = excited state|
Sodium is a chemical element with symbol Na (from Latin natrium) and atomic number 11. It is a soft, silver-white, highly reactive metal. Sodium is an alkali metal, being in group 1 of the periodic table, and hence it has a single electron in its outer shell that it readily donates, creating a positively charged atom—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, 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 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 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.
- 1 Characteristics
- 2 Chemistry
- 3 History
- 4 Occurrence
- 5 Commercial production
- 6 Applications
- 7 Biological role
- 8 Safety and precautions
- 9 See also
- 10 References
- 11 Bibliography
- 12 External links
Sodium at standard temperature and pressure is a soft silvery metal that oxidizes to grayish white unless immersed in oil or inert gas, which are the conditions it is usually stored in. It is the third-least dense of all elemental metals, behind only its congeners potassium and lithium. 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. Sodium gas is at first green, then turns purple at higher temperatures.
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 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.
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. Two nuclear isomers have been discovered, the longer-lived one being 24mNa with a half-life of around 20.2 milliseconds. Acute neutron radiation, as from a nuclear criticality accident, converts some of the stable 23Na in human blood to 24Na; the neutron radiation dosage of a victim can be calculated by measuring the concentration of 24Na relative to 23Na.
Sodium is generally less reactive than potassium and more reactive than lithium. 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 air, it forms primarily sodium peroxide with some sodium oxide. Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts, though potassium and lithium have even more negative potentials.
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.
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. 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.
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 catalyst.
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. Lower concentrations of sodium may be determined by atomic absorption spectrophotometry or by potentiometry using ion-selective electrodes.
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 by the addition of cryptands to solutions of sodium in ammonia via disproportionation.
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 agent.
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 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 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.
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.
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. 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. 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.
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 aluminium:
- 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.
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.
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. 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 analyse compounds. 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.
Liquid sodium is used as a heat transfer fluid in some fast 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 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.
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. 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 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 in 2004 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 a couple grams per day tends to lower systolic blood pressure by about two or three mm Hg. It has been estimated that such a decrease in sodium intake would lead to between 9 and 17% fewer cases of hypertension. 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, 2.3g sodium is about 5.9g, or 2.7ml of salt—about a US teaspoon.) The American Heart Association recommends no more than 1.5 grams of sodium per day (but this doesn't apply to people who lose large amounts of sodium in sweat or to people with certain disorders).
On the other hand, a study in 2016 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 grams 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 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. 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. Sodium is the most prevalent metallic ion in extracellular fluid.
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.
In C4 plants, sodium is a micronutrient that aids in 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 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. 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; excess sodium may also be stored in old plant tissue, limiting the damage to new growth.
Safety and precautions
|The fire diamond hazard sign for sodium metal|
Sodium generates flammable hydrogen and caustic sodium hydroxide on contact with water; contact with moisture on skin, eyes or mucous membranes can cause severe burns. Ingestion can cause burns in gastrointestinal tract. Sodium may spontaneously explode in the presence of an oxidizer. It is very reactive with water. 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 sand.
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