Sodium

Sodium, ₁₁Na
General properties
Appearance	silvery white metallic
Standard atomic weight (Ar, standard)	22.98976928(2)
Sodium in the periodic table
	neon ← sodium → magnesium
Atomic number (Z)	11
Group	group 1 (alkali metals)
Period	period 3
Block	s-block
Element category	alkali metal
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)
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
	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
Magnetic susceptibility	+16.0·10−6 cm3/mol (298 K)
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)
Main isotopes of sodium
	view; talk; edit; \| references

Sodium (symbol Na, from the Latin name natrium) is the chemical element number 11 in the periodic table of elements. It follows that its nucleus includes 11 protons, and 11 electrons orbit around it (according to the simplified model known as "Niels Bohr atom"). Even if many isotopes can be artificially made, all decay in a short time. As a result, all sodium found in nature (mainly in sea water) has the composition ₁₁Na²³, meaning that the nucleus includes 12 neutrons. The atomic mass of sodium is 22.9898; if it is rounded, it would be 23.

Sodium pellets in a container

PropertiesEdit

Sodium is a light, silver-coloured metal. Sodium is so soft that it can be easily cut with a knife. When it is cut, the surface will become white over time. This is because it reacts with air to form sodium hydroxide and sodium carbonate. Sodium is a little lighter than water; when it reacts with water it floats. This reaction is very fast. Hydrogen and sodium hydroxide are produced. The hydrogen may ignite. Since sodium melts at a low temperature, it melts when it reacts with water. It has one valence electron which is removed easily, making it highly reactive.

Compared with other alkali metals (metals in the first column of the periodic table), sodium is usually less reactive than potassium and more reactive than lithium.^[4]

Chemical compoundsEdit

These are chemical compounds that contain sodium ions. Sodium only exists in 1 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 increase 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 toothpastes, 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, 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
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, weak reducing agent, used to preserve dried food
Sodium tellurate, strong oxidizing agent
Sodium telluride, strong reducing agent, reacts with air easily
Sodium tellurite, main tellurite compound

Discovery and nameEdit

Sodium was discovered by Sir Humphrey Davy, an English scientist, back in 1807. He made it by the electrolysis of sodium hydroxide. It is named after soda, a name for sodium hydroxide or sodium carbonate.

Use as elementEdit

It is used in the preparation of organic compounds. It is also used in the street lights that are orange, and ultra violet lights.

Use as compoundsEdit

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

Occurrence and productionEdit

Sodium does not exist as an element in nature; its easily removed valence electron is too reactive. It exists as an ion in chemical compounds. Sodium ions are found in the ocean. It is also found as sodium chloride in the earth's crust, where it is mined.

Sodium is normally made by electrolysis of very hot sodium chloride that was melted.

Use in organismsEdit

Sodium ion in the form of sodium chloride is needed in the human body, but large amounts of it cause problems, which is why one should not eat too much salt and other food items with huge sodium amount (such as biscuits with baking soda). Many organisms in the ocean depend on the proper concentration of ions in sea water to live.

Related pagesEdit

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

ReferencesEdit

↑ Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
↑ 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. ISBN 0-8493-0464-4.
↑ De Leon, N. "Reactivity of Alkali Metals". Indiana University Northwest. Retrieved 2007-12-07.

This short article about chemistry can be made longer. You can help Wikipedia by adding to it.

[CIAAW2016-1] Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.

[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] Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.

[4] De Leon, N. "Reactivity of Alkali Metals". Indiana University Northwest. Retrieved 2007-12-07.

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