Chlorine

17 sulfur ← chlorine → argon F ↑ Cl ↓ Br periodic table
General
Name, Symbol, Number	chlorine, Cl, 17
Chemical series	halogens
Group, Period, Block	17, 3, p
Appearance	yellowish green
Atomic mass	35.453(2) g/mol
Electron configuration	[Ne] 3s2 3p5
Electrons per shell	2, 8, 7
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa) 3.2 g/L
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) 33.949 J/(mol·K)
Vapor pressure P/Pa 1 10 100 1 k 10 k 100 k at T/K 128 139 153 170 197 239
Atomic properties
Crystal structure	orthorhombic
Oxidation states	±1, 3, 5, 7 (strongly acidic oxide)
Electronegativity	3.16 (Pauling scale)
Ionization energies (more)	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
Miscellaneous
Magnetic ordering	nonmagnetic
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
Notable isotopes
Main article: Isotopes of chlorine iso NA half-life DM DE (MeV) DP 35Cl 75.77% Cl is stable with 18 neutrons 36Cl syn 3.01×105 y β- 0.709 36Ar ε - 36S 37Cl 24.23% Cl is stable with 20 neutrons

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.

Occurrence

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).

Discovery

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 1810, Sir Humphry Davy's experiments indicated that this gas was an element, not a compound. He named it chlorine, from the Greek word χλωρóς (chloros), meaning greenish yellow.

Notable characteristics

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).

Isotopes

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.

Safety

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 production

Chlorine gas can be produced by several methods. Three industrial methods involve electrolysis of sodium chloride (NaCl) dissolved in water.

Mercury cell electrolysis

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.

Diaphragm cell electrolysis

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.

Membrane cell electrolysis

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:

2NaCl + 2H₂0 ---> Cl₂ + H₂ + 2 NaOH

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.

Other methods

• Before the use of electrolytic methods, chlorine was produced by the Deacon process, in which hydrogen chloride was oxidized with oxygen or air, in the presence of CuCl₂ used as a catalyst. In this case, the reaction mixture is extremely corrosive, and industrial use of this method is difficult. The reaction is:

4HCl + O₂ → 2Cl₂ + 2H₂O

• When Carl Wilhelm Scheele isolated chlorine in a laboratory, he heated brine with acid and manganese dioxide. The reaction can be written as:

2NaCl + 2H₂SO₄ + MnO₂ → Na₂SO₄ + MnSO₄ + 2H₂O + Cl₂

• Small amounts of chlorine gas can be produced in a laboratory by adding concentrated hydrochloric acid to a solution of sodium chlorate.

Compounds

Chlorine forms a wide range of compounds with metals and nonmetals.

• Salts: In salts with metallic cations, chlorine can take the form of various anions, including chloride (Cl⁻), chlorate (ClO₃⁻), chlorite (ClO₂⁻), hypochlorite (ClO⁻), and perchlorate (ClO₄⁻).
• Fluorides: Chlorine can combine with fluorine to form chlorine monofluoride (ClF), chlorine trifluoride (ClF₃), and chlorine pentafluoride (ClF₅).
• Oxides: Chlorine forms compounds with oxygen, including chlorine dioxide (ClO₂), dichlorine monoxide (Cl₂O), and dichlorine heptoxide (Cl₂O₇).
• Acids: Acids containing chlorine include hydrochloric acid (HCl), chloric acid (HClO₃), and perchloric acid (HClO₄).
• Organochlorides: Many organic compounds contain one or more chlorine atoms. They include solvents such as carbon tetrachloride, plastics such as polyvinyl chloride (PVC), and pesticides such as DDT (dichloro diphenyl trichloroethane).

Applications

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.

• Chlorine gas, also known as bertholite, was first used as a chemical weapon against humans during World War I. The German chemical conglomerate IG Farben had been producing chlorine as a by-product of their dye manufacturing process. In cooperation with Fritz Haber of the Kaiser Wilhelm Institute for Chemistry in Berlin, they developed methods of discharging chlorine gas against entrenched enemy.

• In the form of hypochlorous acid (HOCl), chlorine is used to kill bacteria and other microbes from drinking water supplies and swimming pools. Even small water supplies are now routinely chlorinated. Chlorine, however, persists in the water after treatment and can form organochlorine compounds. Some municipalities therefore use ozone to kill bacteria in water supplies.

• Chlorine is used extensively as an oxidizing agent in chemical reactions.

• It is used in the process of extracting bromine from seawater.

• Chlorine atoms can be substituted for hydrogen atoms in organic compounds, thus generating products with various properties. Examples are chloroform, carbon tetrachloride, polyvinyl chloride, and synthetic rubber.

• Chloride ions have important physiological roles. For instance, in the central nervous system, the inhibitory action of glycine and some of the action of the neurotransmitter GABA (gamma-aminobutyric acid) rely on the entry of chloride ions into specific neurons. Also, the chloride-bicarbonate exchanger biological transport protein relies on the chloride ion to increase the blood's capacity for carbon dioxide (in the form of the bicarbonate ion).

See also

• Chemical element
• Periodic table
• Periodic table, main group elements

References

ISBN links support NWE through referral fees

• Los Alamos National Laboratory – Chlorine Retrieved February 20, 2008.
• WebElements.com – Chlorine Retrieved February 20, 2008.

External links

All links retrieved February 15, 2017.

• The Chlorine Institute (trade organization, located in Arlington Va. USA, consisting of chlorine producers, packagers, responders, end users and associated members)
• Chlorine and compounds National Pollutant Inventory

Agents of Chemical Warfare
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

E numbers
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)