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| pale yellow-green gas|
|Name, symbol, number||chlorine, Cl, 17|
|Pronunciation|| // KLOHR-een|
or // KLOHR-ən
|Group, period, block||17, 3, p|
|Standard atomic weight||35.45(1) g·mol−1|
|Electron configuration||[Ne] 3s2 3p5|
|Electrons per shell||2, 8, 7 (Image)|
|Density|| (0 °C, 101.325 kPa)|
|Liquid density at b.p.||1.5625 g·cm−3|
|Melting point||171.6 K, -101.5 °C, -150.7 °F|
|Boiling point||239.11 K, -34.04 °C, -29.27 °F|
|Critical point||416.9 K, 7.991 MPa|
|Heat of fusion||(Cl2) 6.406 kJ·mol−1|
|Heat of vaporization||(Cl2) 20.41 kJ·mol−1|
|Specific heat capacity|| (25 °C) (Cl2)|
|Oxidation states|| 7, 6, 5, 4, 3, 2, 1, -1|
(strongly acidic oxide)
|Electronegativity||3.16 (Pauling scale)|
| Ionization energies
||1st: 1251.2 kJ·mol−1|
|2nd: 2298 kJ·mol−1|
|3rd: 3822 kJ·mol−1|
|Covalent radius||102±4 pm|
|Van der Waals radius||175 pm|
|Electrical resistivity||(20 °C) > 10 Ω·m|
|Thermal conductivity||(300 K) 8.9×10−3 W·m−1·K−1|
|Speed of sound||(gas, 0 °C) 206 m/s|
|CAS registry number||7782-50-5|
|Most stable isotopes|
|Main article: Isotopes of chlorine|
Chlorine (chemical symbol Cl) is a chemical element. Its atomic number (which is the number of protons in it) is 17, and its atomic mass is 35.45. It is part of the 7th column (halogens) on the periodic table of elements.
Chlorine is a very irritating and greenish-yellow gas. It has a strong smell like bleach. It is toxic. It can be made into a liquid when cooled. It is heavier than air.
Chlorine is highly reactive. It is more reactive than bromine but less reactive than fluorine. It reacts with most things to make chlorides. It can even burn things instead of oxygen. It dissolves in water to make a mixture of hypochlorous acid and hydrochloric acid. The more acidic it is, the more chlorine is made; the more basic it is, the more hypochlorous acid (normally turned into hypochlorite) and hydrochloric acid (normally turned into chlorides) are there. Chlorine reacts with bromides and iodides to make bromine and iodine.
Chlorine exists in several oxidation states: -1, +1, +3, +4, +5, and +7. The -1 state is most often in chloride. Chlorides are not reactive. Compounds containing chlorine in its +1 oxidation state are hypochlorites. Only one is common. They are a strong oxidizing agent, as are all + oxidation state compounds. +3 is in chlorites. +4 is in chlorine dioxide, a common chlorine compound that is not a chloride. +5 is in chlorates. +7 is in perchlorates. Hypochlorites are most reactive, while perchlorates are the least reactive.
Chlorine oxides can be made, but most of them are very reactive and unstable.
Chlorine is not found as an element. Sodium chloride is the most common chlorine ore. It is in the ocean (sea salt) and in the ground (rock salt). There are some organic compounds that have chlorine in them, too.
It is made by electrolysis (the passing of electricity through a solution to make chemical reactions happen) of sodium chloride. This is known as the chloralkali process. It can also be made by reacting hydrogen chloride with oxygen and a catalyst. It can be made in the laboratory by reacting manganese dioxide with hydrochloric acid. It is made when sodium hypochlorite reacts with hydrochloric acid. This is a dangerous reaction that can happen without anyone knowing.
Chlorine is used widely to purify water (usually in a swimming pool), as a disinfectant and bleach, and in the making of many important compounds including chloroform and carbon tetrachloride. It was used as a poison gas in some wars.
It was discovered in 1774 by Carl Wilhelm Scheele who thought it had oxygen in it. Chlorine was named in 1810 by Humphry Davy who insisted it was an element. The US made all water chlorinated (added chlorine to water) by 1918.
It is poisonous in large amounts and can damage skin. When it is inhaled, it irritates the lungs, eyes, and skin badly. It can cause fire with some things because it is very reactive. It is heavier than air, so it can fill up enclosed spaces.
- Chlorine, Gas Encyclopaedia, Air Liquide
- 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. .