Sodium

Sodium is a chemical element with an atomic number of 11. Its symbol is Na (from its Latin name natrium). It is an alkali metal. Although sodium has many isotopes, most decay in a short time. Because of this, all sodium in nature (mainly found in seawater) is of the isotope ₁₁Na²³. The atomic mass of sodium is 22.9898.

Sodium, ₀₀Na

Sodium
Appearance	silvery white metallic
Standard atomic weight Ar°(Na)
	22.98976928(2)[1]
Sodium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Li ↑ Na ↓ K neon ← sodium → magnesium
Group	group 1: hydrogen and alkali metals
Period	period 3
Block	s-block
Electron configuration	[Ne] 3s1
Electrons per shell	2, 8, 1
Physical properties
Phase at STP	solid
Melting point	370.944 K ​(97.794 °C, ​208.029 °F)
Boiling point	1156.090 K ​(882.940 °C, ​1621.292 °F)
Density (near r.t.)	0.968 g/cm3
when liquid (at m.p.)	0.927 g/cm3
Critical point	2573 K, 35 MPa (extrapolated)
Heat of fusion	2.60 kJ/mol
Heat of vaporization	97.42 kJ/mol
Molar heat capacity	28.230 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 554 617 697 802 946 1153
Atomic properties
Oxidation states	−1, 0,[2] +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
Spectral lines of sodium
Other properties
Natural occurrence	primordial
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[3]
Molar magnetic susceptibility	+16.0·10−6 cm3/mol (298 K)[4]
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)
Symbol	"Na": from New Latin natrium, coined from German Natron, 'natron'
Isotopes of sodium v e
Main isotopes[5] Decay abun­dance half-life (t1/2) mode pro­duct 22Na trace 2.6019 y β+ 22Ne 23Na 100% stable 24Na trace 14.9560 h β− 24Mg
Category: Sodium view talk edit | references

Sodium pellets in a container

Properties

Sodium is a light-weight, silver-colored metal. Sodium is soft. It can easily be cut with a knife. When someone cuts it, the exposed part will become white over time. It reacts with air to form, at first sodium oxide, then slowly sodium hydroxide and sodium carbonate. Sodium is a little less dense than water. It floats and reacts instantly with water, producing hydrogen and sodium hydroxide. This reaction makes a lot of heat, usually causing the hydrogen to light on fire. When this happens, sodium melts because of its low melting point. Sodium is highly reactive because it has one valence electron, which is easily removed.

Compared with other alkali metals, sodium is less reactive than potassium and more reactive than lithium.^[6]

Chemical compounds

These are chemical compounds that contain sodium ions. Sodium only exists in one oxidation state: +1.

• Sodium aluminum fluoride, used to make aluminum
• Sodium amide, very strong base
• Sodium arsenite, colorless solid, very toxic
• Sodium arsenate, oxidizing agent, very toxic
• Sodium azide, used in airbags
• Sodium bicarbonate, baking soda, used in cooking
• Sodium bismuthate, oxidizing agent, used to test for manganese
• Sodium bisulfate, acidic, used to lower pH
• Sodium bromate, oxidizing agent, used to dye hair
• Sodium bromide, rare, used in some medicine
• Sodium carbonate, used to make glass
• Sodium chlorate, used in some explosives
• Sodium chlorite, used in disinfectants
• Sodium chloride, table salt
• Sodium chromate, yellow, oxidizing agent, toxic
• Sodium dichromate, orange, oxidizing agent, toxic
• Sodium fluoride, used in toothpaste, bitter, toxic in large doses
• Sodium hydroxide, lye, used in soap, strong base
• Sodium hypochlorite, bleach, disinfectant
• Sodium hypophosphite, reducing agent, poisonous
• Sodium iodate, oxidizing agent, prevents iodine deficiency
• Sodium iodide, a weak reducing agent, prevents iodine deficiency
• Sodium manganate, rare green solid
• Sodium nitrate, used in blasting powder
• Sodium nitrite, used in food preservation
• Sodium periodate, oxidizing agent
• Sodium permanganate, less common than potassium permanganate, oxidizing agent
• Sodium phosphate, various uses
• Sodium phosphide, catalyst, used to speed up chemical reactions
• Sodium phosphite, toxic, reducing agent
• Sodium selenate, strong oxidizing agent, other selenium compounds
• Sodium selenide, strong reducing agent, reactive
• Sodium selenite, weak oxidizing agent, vitamin supplement
• Sodium sulfate, bitter, laxative
• Sodium sulfite, a weak reducing agent, used to preserve dried food
• Sodium tellurate, strong oxidizing agent
• Sodium telluride, a strong reducing agent, reacts with air easily
• Sodium tellurite, main tellurite compound

Discovery and name origins

Sodium was discovered by Sir Humphrey Davy, an English scientist, in 1807. He created it by electrolyzing sodium hydroxide. Davy named the element after soda, a name for sodium hydroxide or sodium carbonate.

Uses

Scientists can use it in the creation of organic compounds. It is used in orange streetlights and lamps that emit ultraviolet light.

Sodium compounds are used in soaps, toothpaste, baking, and antacids.

The human body needs sodium ions, taken in the form of sodium chloride, to live. Too much of it can cause health problems. Many organisms in the ocean depend on the concentration of sodium ions in water to survive.

Occurrence and production

Sodium does not occur as an element in nature, because it is not stable enough. It exists only in chemical compounds. Sodium ions are found in the ocean and in the Earth's crust.

Sodium is normally made by electrolysis of sodium chloride, which is mined from the Earth's crust.

Related pages

• List of common elements
• Hyponatremia (a medical problem caused by not having enough sodium in the body)

References

1. "Standard Atomic Weights: Sodium". CIAAW. 2005.
2. The compound NaCl has been shown in experiments to exists in several unusual stoichiometries under high pressure, including Na₃Cl in which contains a layer of sodium(0) atoms; see Zhang, W.; Oganov, A. R.; Goncharov, A. F.; Zhu, Q.; Boulfelfel, S. E.; Lyakhov, A. O.; Stavrou, E.; Somayazulu, M.; Prakapenka, V. B.; Konôpková, Z. (2013). "Unexpected Stable Stoichiometries of Sodium Chlorides". Science. 342 (6165): 1502–1505. arXiv:1310.7674. Bibcode:2013Sci...342.1502Z. doi:10.1126/science.1244989. PMID 24357316. S2CID 15298372.
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.
4. Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
5. Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
6. De Leon, N. "Reactivity of Alkali Metals". Indiana University Northwest. Archived from the original on 2018-10-16. Retrieved 2007-12-07.