|Appearance||silvery white metallic|
|Standard atomic weight Ar, std(Na)||22.98976928(2)|
|Sodium in the periodic table|
|Atomic number (Z)||11|
|Group||group 1: hydrogen and alkali metals|
|Electron configuration||[Ne] 3s1|
|Electrons per shell||2, 8, 1|
|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)|
|Oxidation states||−1, +1 (a strongly basic oxide)|
|Electronegativity||Pauling scale: 0.93|
|Atomic radius||empirical: 186 pm|
|Covalent radius||166±9 pm|
|Van der Waals radius||227 pm|
|Spectral lines of sodium|
|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)|
|Molar 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|
|Brinell hardness||0.69 MPa|
|Discovery and first isolation||Humphry Davy (1807)|
|Symbol||"Na": from New Latin natrium, coined from German Natron, 'natron'|
|Main isotopes of sodium|
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.
Properties[change | change source]
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 and reacts. This reaction is very fast. Hydrogen and sodium hydroxide are produced and lots of heat is created which usually causes the hydrogen to 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.
Chemical compounds[change | change source]
- 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 name[change | change source]
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 element[change | change source]
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 compounds[change | change source]
Sodium compounds are used in soaps, toothpaste, baking and antiacids. .
Occurrence and production[change | change source]
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.
Use in organisms[change | change source]
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 pages[change | change source]
- List of common elements
- Hyponatremia (a medical problem caused by not having enough sodium in the body).
References[change | change source]
- "Standard Atomic Weights: Sodium". CIAAW. 2005.
- 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. Archived from the original on 2018-10-16. Retrieved 2007-12-07.