|Name, Symbol, Number||chlorine, Cl, 17|
|Group, Period, Block||17, 3, p|
|Atomic mass||35.453(2) g/mol|
|Electron configuration||[Ne] 3s2 3p5|
|Electrons per shell||2, 8, 7|
|Density||(0 °C, 101.325 kPa)
|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|
|Heat of vaporization||(Cl2) 20.41 kJ/mol|
|Heat capacity||(25 °C) (Cl2)
|Oxidation states||±1, 3, 5, 7
(strongly acidic oxide)
|Electronegativity||3.16 (Pauling scale)|
|1st: 1251.2 kJ/mol|
|2nd: 2298 kJ/mol|
|3rd: 3822 kJ/mol|
|Atomic radius||100 pm|
|Atomic radius (calc.)||79 pm|
|Covalent radius||99 pm|
|Van der Waals radius||175 pm|
|Electrical resistivity||(20 °C) > 10 Ω·m|
|Thermal conductivity||(300 K) 8.9 mW/(m·K)|
|Speed of sound||(gas, 0 °C) 206 m/s|
|CAS registry number||7782-50-5|
Chlorine (chemical symbol Cl, atomic number 17) is a nonmetal that belongs to a group of chemical elements known as halogens. At ordinary temperatures and pressures, pure chlorine is a highly reactive, poisonous gas with a greenish-yellow color and an unpleasant odor. Its chemical formula is Cl2. Given its high reactivity, the free element is not found in nature.
On the other hand, chloride (Cl−) ions are abundant in nature and necessary for most forms of life, including human life. They are part of various salts and are found in solution in naturally occurring waters. Common salt or table salt is the compound sodium chloride (NaCl).
Chlorine, its ions, and its compounds are widely used in the manufacture of many products, including paper, bleach, antiseptics, dyestuffs, pesticides, paints, solvents, plastics, medicines, and textiles. Drinking water supplies and swimming pools are usually chlorinated as a way to kill bacteria.
As noted above, elemental chlorine is not found in nature. Rather, chlorine is found mainly in the form of the chloride ion, a component of salts deposited in the earth or dissolved in the oceans. About 1.9 percent of the mass of seawater is chloride ions. Higher concentrations of chloride are found in the Dead Sea and in underground brine deposits.
Given that most chloride salts are soluble in water, the abundance of chloride-containing minerals is higher in regions with dry climates and deep underground, where the salts seldom come in contact with water. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate).
Chlorine was discovered in 1774 by Swedish chemist Carl Wilhelm Scheele, who observed the greenish-yellow gas when experimenting with seawater. He mistakenly thought it was a compound of oxygen and hydrochloric acid (HCl), and he called it dephlogisticated marine acid. He used the term "dephlogisticated" to indicate that the material could not burn, and the term "marine acid" was the name for hydrochloric acid. (According to the "phlogiston theory" commonly held at that time, phlogiston was an invisible, weightless substance released by burning materials. When a material could no longer burn, it was said to be "dephlogisticated.")
In the periodic table, chlorine is located in group 17 (former group 7A), the halogen family, between fluorine and bromine. In addition, it lies in period three, between sulfur and argon. Chlorine gas is about 2.5 times as dense as air.
Elemental chlorine (Cl2) combines readily with nearly all other elements. It is not, however, as extremely reactive as fluorine. It reacts with various metals to form chloride salts, which contain the chloride anion (Cl−). A familiar example is table salt, or sodium chloride, with the chemical formula NaCl. In water, NaCl dissolves to produce Na+ and Cl− ions.
Chlorine also forms compounds in which it is covalently bound to various nonmetals. For example, it is covalently bound to carbon atoms in many organic compounds. Under suitable conditions, it combines with other halogens—fluorine, bromine, and iodine—to form "interhalogen" compounds such as ClF, ClF3, ClF5, ClBr, and ICl.
One liter of water dissolves 3.10 liters of gaseous chlorine at ten °C, but the same amount of water dissolves only 1.77 liters chlorine gas at 30 °C. In water, it exists as a mixture of chlorine (Cl2), hydrochloric acid (HCl), and hypochlorous acid (HOCl).
Chlorine has nine isotopes, with atomic mass numbers ranging from 32 to 40. Of these, the two main, stable isotopes are 35Cl (75.77 percent) and 37Cl (24.23 percent). As the relative proportions of these two isotopes are three to one respectively, chlorine atoms in bulk have an apparent atomic weight of 35.5 atomic mass units.
The environment also contains trace amounts of the radioactive isotope 36Cl. It decays to sulfur-36 and argon-36, with a combined half-life of 308,000 years. Based on its half-life, high solubility in water, and nonreactive nature, this isotope is suitable for geologic dating in the range of 60,000 to one million years.
Chlorine is a toxic gas that irritates the respiratory systems. Being heavier (denser) than air, it tends to accumulate at the bottom of poorly ventilated spaces. In addition, chlorine gas is a potential oxidizer that may react with flammable materials.
Chlorine gas can be produced by several methods. Three industrial methods involve electrolysis of sodium chloride (NaCl) dissolved in water.
The first electrolytic method of producing chlorine on an industrial scale involved the use of a mercury cell. Liquid mercury formed the cathode, titanium was used for a set of anodes, and a solution of sodium chloride was positioned between the electrodes. When an electrical current was applied, chlorine was released at the anodes and sodium dissolved in the mercury to form an amalgam.
Mercury was regenerated from the amalgam by reacting it with water. The reaction produced hydrogen and sodium hydroxide—two useful byproducts. This method, however, consumed vast amounts of energy, and it led to concerns about mercury emissions.
In this method, the electrolysis of sodium chloride solution is performed in a cell with an iron grid cathode with an asbestos diaphragm on it. The diaphragm prevents the chlorine (which forms at the anode) and the sodium hydroxide (which forms at the cathode) from re-mixing. This method avoids the use of mercury and consumes less energy than the mercury cell. The sodium hydroxide, however, is not as easily concentrated and precipitated into a useful substance.
The electrolytic cell is divided into two by a membrane acting as an ion exchanger. Saturated sodium chloride solution is passed through the anode compartment, leaving at a lower concentration. Sodium hydroxide solution is circulated through the cathode compartment, exiting at a higher concentration. The electrolytic process generates chlorine in the anode compartment and hydrogen and sodium hydroxide in the cathode compartment. A portion of the concentrated sodium hydroxide solution leaving the cathode is diverted as product, while the remainder is diluted with deionized water and passed through the electrolyzer again.
The overall chemical equation is:
This method is nearly as efficient as the diaphragm cell and produces very pure sodium hydroxide, but it requires a very pure solution of sodium chloride.
Chlorine forms a wide range of compounds with metals and nonmetals.
Chlorine, its ions, and its compounds are widely used in the manufacture of many products, including paper, bleach, antiseptics, dyestuffs, pesticides, paints, solvents, plastics, medicines, and textiles. Some specific uses are mentioned below.
All links retrieved May 16, 2013.
|Blood agents:||Cyanogen chloride (CK) – Hydrogen cyanide (AC)|
|Blister agents:||Lewisite (L) – Sulfur mustard gas (HD, H, HT, HL, HQ) – Nitrogen mustard gas (HN1, HN2, HN3)|
|Nerve agents:||G-Agents: Tabun (GA) – Sarin (GB) – Soman (GD) – Cyclosarin (GF) | V-Agents: VE – VG – VM – VX|
|Pulmonary agents:||Chlorine – Chloropicrin (PS) – Phosgene (CG) – Diphosgene (DP)|
|Incapacitating agents:||Agent 15 (BZ) – KOLOKOL-1|
|Riot control agents:||Pepper spray (OC) – CS gas – CN gas (mace) – CR gas|
|Colours (E100-199) • Preservatives (E200-299) • Antioxidants & Acidity regulators (E300-399) • Thickeners, stabilisers & emulsifiers (E400-499) • pH regulators & anti-caking agents (E500-599) • Flavour enhancers (E600-699) • Miscellaneous (E900-999) • Additional chemicals (E1100-1599)|
|Waxes (E900-909) • Synthetic glazes (E910-919) • Improving agents (E920-929) • Packaging gases (E930-949) • Sweeteners (E950-969) • Foaming agents (E990-999)|
|L-cysteine (E920) • L-cystine (E921) • Potassium persulfate (E922) • Ammonium persulfate (E923) • Potassium bromate (E924) • Chlorine (E925) • Chlorine dioxide (E926) • Azodicarbonamide (E927) • Carbamide (E927b) • Benzoyl peroxide (E928) • Calcium peroxide (E930)|
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