Jump to content

Chlorine

Edit links
Page semi-protected
From Wikipedia, the free encyclopedia

This is an old revision of this page, as edited by Kleuske (talk | contribs) at 10:45, 16 October 2016 (Undid revision 744612712 by Vijay Patel 786 (talk)). The present address (URL) is a permanent link to this revision, which may differ significantly from the current revision.

This article is about the chemical element. For the bleach, see Sodium hypochlorite. For the 2013 American film, see Chlorine (2013 film). For the 2015 Italian film, see Chlorine (2015 film).
"Cl" and "Cl2" redirect here. For other uses, see CL (disambiguation) and CL2 (disambiguation).

Chlorine, 17Cl
A glass container filled with chlorine gas
Chlorine
Pronunciation/ˈklɔːriːn, -aɪn/ ​(KLOR-een, -⁠eyen)
Appearancepale yellow-green gas
Standard atomic weight Ar°(Cl)
Chlorine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
F
↑
Cl
↓
Br
sulfur ← chlorine → argon
Atomic number (Z)17
Groupgroup 17 (halogens)
Periodperiod 3
Block  p-block
Electron configuration[Ne] 3s2 3p5
Electrons per shell2, 8, 7
Physical properties
Phase at STPgas
Melting point(Cl2) 171.6 K ​(−101.5 °C, ​−150.7 °F)
Boiling point(Cl2) 239.11 K ​(−34.04 °C, ​−29.27 °F)
Density (at STP)3.2 g/L
when liquid (at b.p.)1.5625 g/cm3[3]
Triple point172.22 K, ​1.392 kPa[4]
Critical point416.9 K, 7.991 MPa
Heat of fusion(Cl2) 6.406 kJ/mol
Heat of vaporization(Cl2) 20.41 kJ/mol
Molar heat capacity(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
Oxidation statescommon: −1, +1, +3, +5, +7
+2,[5] +4,[5] +6[5]
ElectronegativityPauling scale: 3.16
Ionization energies
  • 1st: 1251.2 kJ/mol
  • 2nd: 2298 kJ/mol
  • 3rd: 3822 kJ/mol
  • (more)
Covalent radius102±4 pm
Van der Waals radius175 pm
Color lines in a spectral range
Spectral lines of chlorine
Other properties
Natural occurrenceprimordial
Crystal structure ​orthorhombic (oS8)
Lattice constants
Orthorhombic crystal structure for chlorine
a = 630.80 pm
b = 455.83 pm
c = 815.49 pm (at triple point)[6]
Thermal conductivity8.9×10−3 W/(m⋅K)
Electrical resistivity>10 Ω⋅m (at 20 °C)
Magnetic orderingdiamagnetic[7]
Molar magnetic susceptibility−40.5×10−6 cm3/mol[8]
Speed of sound206 m/s (gas, at 0 °C)
CAS NumberCl2: 7782-50-5
History
Discovery and first isolationCarl Wilhelm Scheele (1774)
Recognized as an element byHumphry Davy (1808)
Isotopes of chlorine
Main isotopes[9] Decay
abun­dance half-life (t1/2) mode pro­duct
35Cl 76% stable
36Cl trace 3.01×105 y β− 36Ar
ε 36S
37Cl 24% stable
 Category: Chlorine
| references

Chlorine is a chemical element with symbol Cl and atomic number 17. It has a relative atomic mass of about 35.5. Chlorine is in the halogen group (17) and is the second lightest halogen, following fluorine. The element is a yellow-green diatomic gas under standard conditions. Chlorine has the highest electron affinity and the third highest electronegativity of all the reactive elements. For this reason, chlorine is a strong oxidizing agent. Free chlorine is rare on Earth, and is usually a result of direct or indirect oxidation by oxygen.

The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630, chlorine gas was first synthesized in a chemical reaction, but not recognized as a fundamentally important substance. Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, romanized: khlôros, lit. 'pale green'.

Nearly all chlorine in the Earth's crust is in the form of ionic chloride compounds, which includes table salt. It is the second most abundant halogen and 21st most abundant chemical element in Earth's crust. Elemental chlorine is commercially produced from brine by electrolysis. The high oxidizing potential of elemental chlorine led to the development of commercial bleaches and disinfectants, and a reagent for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, and many intermediates for production of plastics and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary.

In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils as part of the immune response against bacteria.

Elemental chlorine at high concentrations is extremely dangerous and poisonous for all living organisms, and was used in World War I as the first gaseous chemical warfare agent.

Characteristics

Physical characteristics of chlorine and its compounds

Chlorine, liquefied under a pressure of 7.4 bar at room temperature, displayed in a quartz ampule embedded in acrylic glass.

At standard temperature and pressure, two chlorine atoms form the diatomic molecule Cl2,[10] a yellow-green gas with a strong, distinctive odor, familiar to most people from common household bleach.[11] The bonding between the two atoms is relatively weak (only 242.580 ± 0.004 kJ/mol), which makes the Cl2 molecule highly reactive. The boiling point at standard pressure is around −34 ˚C, but it liquefies at room temperature at 740 kPa (107 psi).[12]

Elemental chlorine is yellow-green, but the chloride ion, in common with other halide ions, has no color in either minerals or solutions (example, table salt). As with other halogens, chlorine atoms impart no color to organic chlorides when they replace hydrogen atoms organic compounds, such as tetrachloromethane. The melting point and density of these compounds is increased by substitution of chlorine in place of hydrogen. Compounds of chlorine with other halogens are colored, as are many chlorine oxides.

Chemical characteristics

Along with fluorine, bromine, iodine, and astatine, chlorine is a member of the halogen series that forms the group 17 (formerly VII, VIIA, or VIIB) of the periodic table. Chlorine combines with almost all other elements to produce compounds that are usually called chlorides. Chlorine gas reacts with most organic compounds, and sluggishly supports the combustion of hydrocarbons.[13]

Compounds

See also: Category:Chlorine compounds

Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state and +4 in chlorine dioxide, with respective oxidation states of 0 and 4+ (see table below, and also structures in chlorite).[14] Chlorine typically has a −1 oxidation state in compounds, except for compounds containing fluorine, oxygen, and nitrogen, all of which are even more electronegative than chlorine. Progressing through the states, hydrochloric acid can be oxidized using manganese dioxide, or hydrogen chloride gas oxidized catalytically by air to form elemental chlorine gas.[15]

Oxidation
state		 Name Formula Characteristic compounds
−1 chlorides Cl− ionic chlorides, organic chlorides, hydrochloric acid
0 chlorine Cl2 elemental chlorine
+1 hypochlorites ClO− sodium hypochlorite, calcium hypochlorite, dichlorine monoxide
+3 chlorites ClO−
2		 sodium chlorite
+4 chlorine(IV) ClO
2		 chlorine dioxide
+5 chloryl, chlorates ClO−
3 ClO+
2		 potassium chlorate, chloric acid, dichloryl trisulfate [ClO2]2[S3O10].
+6 chlorine(VI) Cl
2O
6		 dichlorine hexoxide (gas). In liquid or solid disproportionates to mix of +5 and +7 oxidation states, as ionic chloryl perchlorate [ClO
2]+
[ClO
4]−
+7 perchlorates ClO−
4		 perchloric acid, perchlorate salts such as magnesium perchlorate, dichlorine heptoxide

Hydrolysis of free chlorine or disproportionation in water

At 25 °C and atmospheric pressure, one liter of water dissolves 3.26 g or 1.125 L of gaseous chlorine.[16] Chlorine dissolves in water to produce dissolved chlorine (Cl2), and by reversible reaction, hydrochloric acid, and hypochlorous acid:

Cl2 + H2O ⇌ HCl + HClO

The conversion to the right is called disproportionation because the ingredient chlorine both increases and decreases in formal oxidation state. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide, and in this way, chlorine bleach is produced.[17]

Cl2 + 2 OH− → ClO− + Cl− + H2O

Dissolved elemental chlorine gas is present only in a neutral or acidic solution.

Chlorides
Main article: Chloride

Chlorine combines with almost all elements to produce chlorides. Compounds with oxygen, nitrogen, xenon, and krypton are known, but are not formed by direct reaction of the elements.[18] Chloride is one of the most common anions in nature. Hydrogen chloride and its aqueous solution, hydrochloric acid, are commercially produced by the megaton annually as valued intermediates and as undesirable pollutants.

Chlorine oxides

Chlorine forms a variety of oxides: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine hexoxide (Cl2O6), and dichlorine heptoxide (Cl2O7). The anionic derivatives of these same oxides are also well known and include hypochlorite (ClO−), chlorite (ClO−
2), chlorate (ClO−
3), and perchlorate (ClO−
4). The acid forms of these anions are hypochlorous acid (HOCl), chlorous acid (HClO2), chloric acid (HClO3) and perchloric acid (HClO4), respectively. The chloroxy cation chloryl (ClO2+) is known and has the same structure as chlorite but with a positive charge and chlorine in the +5 oxidation state.[19] The compound "chlorine trioxide" does not occur, but is found in gas form as the dimeric dichlorine hexoxide (Cl2O6) with a +6 oxidation state. This compound in liquid or solid form disproportionates to a mixture of +5 and +7 oxidation states, occurring as the ionic compound chloryl perchlorate, [ClO
2]+
[ClO
4]−
.[20]

In hot concentrated alkali solution, hypochlorite disproportionates:

2 ClO− → Cl− + ClO−
2
ClO− + ClO−
2 → Cl− + ClO−
3

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated to a high temperature, they undergo a further, final disproportionation:

4 ClO−
3 → Cl− + 3 ClO−
4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[21]

Reaction Electrode
potential
Cl− + 2 OH− → ClO− + H2O + 2 e− +0.89 volts
ClO− + 2 OH− → ClO−
2 + H2O + 2 e−		 +0.67 volts
ClO−
2 + 2 OH− → ClO−
3 + H2O + 2 e−		 +0.33 volts
ClO−
3 + 2 OH− → ClO−
4 + H2O + 2 e−		 +0.35 volts

Each step is accompanied at the cathode by

2 H2O + 2 e− → 2 OH− + H2 (−0.83 volts)
Interhalogen compounds

Chlorine forms a variety of interhalogen compounds, such as the chlorine fluorides, chlorine monofluoride (ClF), chlorine trifluoride (ClF
3), chlorine pentafluoride (ClF
5). Chlorides of bromine and iodine are also known. Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively, but cannot oxidize fluoride salts to fluorine.[22]

Organochlorine compounds
Main article: Organochloride

Chlorine often imparts desired properties to an organic compound, in part owing to the relative stability of the C-Cl bond. Organic chloride compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromide or iodide derivatives, but they tend to be cheaper. Chlorine is used extensively in polymers and in organic chemistry.

Several general approaches exist for the formation of C-Cl bonds. Like the other halogens, chlorine undergoes electrophilic addition reactions. Notable is the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Industries often avoid using chlorine directly, preferring cheaper reagents such as hydrogen chloride, as deployed in hydrohalogenation and oxychlorination. The organochlorine compound produced on the largest scale is polyvinyl chloride, 23 billion kilograms of which were produced in 2000.[23]

Occurrence

See also: Category:Halide minerals

Essentially no chlorine was created in the Big Bang. Chlorine in the universe is created and distributed through the interstellar medium from nucleosynthesis in supernovae, via the r-process.[24] This source provided the chlorine found in the Solar System.

In meteorites and on Earth, chlorine is found primarily as the chloride ion in minerals. In the Earth's crust, chlorine is present at average concentrations of about 126 parts per million,[25] predominantly in such minerals as halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate).

Chloride is a component of the salt deposited in the earth and dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Higher concentrations of chloride are found in the Dead Sea and in underground brine deposits and evaporites. Most chloride salts are soluble in water; therefore, chloride-containing minerals are found in abundance only in dry climates or deep underground.

More than 2000 naturally occurring organic chlorine compounds are known.[26]

Isotopes

Main article: Isotopes of chlorine

Chlorine has a wide range of isotopes. The two stable isotopes are 35Cl (75.77%) and 37Cl (24.23%).[27] Together they give chlorine a relative atomic mass of 35.4527 g/mol. The half-integer value for chlorine's weight caused some confusion in the early days of chemistry, when chemists were working on the postulate that atoms were composed of units of hydrogen (see Proust's law) and the existence of isotopes was unsuspected.[28]

Radioactive 36Cl exists in the environment in a ratio with stable isotopes of about 7x10−13 to 1. 36Cl is produced in the atmosphere by spallation of 36Ar from collisions of cosmic ray protons. In the subsurface environment, 36Cl is generated primarily as a result of neutron capture by 35Cl or muon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with a combined half-life of 308,000 years. The half-life of this isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Large amounts of 36Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of 36Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and groundwater, 36Cl is also useful for dating waters less than 50 years before the present. 36Cl has seen other uses in the geological sciences, including dating ice and sediments.[27]

History

The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.[29] Around 1630, chlorine was recognized as a gas by the Flemish chemist and physician Jan Baptist van Helmont.[30]

Carl Wilhelm Scheele

Elemental chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and he is credited with the discovery.[31] He called it "dephlogisticated muriatic acid air" since it is a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid").[31] He failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory).[31] He named the new element within this oxide as muriaticum.[31] Regardless of what he thought, Scheele did isolate chlorine by reacting MnO2 (as the mineral pyrolusite) with HCl:[30]

4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2

Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow green color, and the smell similar to aqua regia.[32]

Common chemical theory at that time held that an acid is a compound that contains oxygen (still sounding in the German and Dutch names of oxygen: sauerstoff or zuurstof, both translating into English as acid substance), so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum.[33][34][35]

In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide).[31] They did not succeed and published a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, but were not convinced.[36]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it is an element, and not a compound.[31] He named this new element "chlorine", from the Greek word χλωρος (chlōros), meaning green-yellow.[37] The name "halogen", meaning "salt producer", was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger. This term was later used as a generic term to describe all the elements in the chlorine family (fluorine, bromine, iodine), after a suggestion by Jöns Jakob Berzelius in 1842.[38][39] In 1823, Michael Faraday liquefied chlorine for the first time,[40][41] and demonstrated that what was then known as "solid chlorine" had a structure of chlorine hydrate (Cl2·H2O).[30]

Bleaching and disinfection

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.[42][43] Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel (now part of Paris, France), by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of sodium hypochlorite. This process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of calcium hypochlorite ("chlorinated lime"), then solid calcium hypochlorite (bleaching powder).[42] These compounds produced low levels of elemental chlorine, and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidizer. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite.[44] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of most elemental chlorine and sodium hydroxide. A related low-temperature electrolysis reaction, the Hooker process, is now used to produce bleach and sodium hypochlorite.

Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti-putrifaction agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This practice was pioneered by Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations (for a more complete history, see below). Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) and public sanitation, particularly in swimming and drinking water.

Chlorine gas was first used as a weapon on April 22, 1915, at Ypres by the German Army.[45][46] The effect on the allies was devastating because the existing gas masks were difficult to deploy and had not been broadly distributed.

Historically important chlorine compounds

In 1826, silver chloride was used to produce photographic images for the first time.[47] Chloroform was first used as an anesthetic in 1847.[47]

Polyvinyl chloride (PVC) was invented in 1912, initially without a purpose.[47]

Production

Liquid chlorine analysis
Main articles: Chlorine production and Chloralkali process

In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. This method, the chloralkali process industrialized in 1892, now provides most industrial chlorine gas.[48] Along with chlorine, the method yields hydrogen gas and sodium hydroxide, which is the most valuable product. The process proceeds according to the following chemical equation:[15]

2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH

The electrolysis of chloride solutions all proceed according to the following equations:

Cathode: 2 H+(aq) + 2 e− → H2(g)
Anode: 2 Cl−(aq) → Cl2(g) + 2 e−

Overall process: 2 NaCl (or KCl) + 2 H2O → Cl2 + H2 + 2 NaOH (or KOH)

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.[49] The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali, but they are not burdened with the problem of mercury disposal and they are more energy efficient.

Membrane cell electrolysis employs permeable membrane as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.[50] This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine.

Membrane cell process for chloralkali production

Oxidation of HCl

In the Deacon Process, hydrogen chloride recovered from the production of organochlorine compounds is recovered as chlorine. The process relies on oxidation using oxygen:

4 HCl + O2 → 2 Cl2 + 2 H2O

The reaction requires a catalyst. As introduced by Deacon, early catalysts were based on copper. Commercial processes, such as the Mitsui MT-Chlorine Process, have switched to chromium and ruthenium-based catalysts.[51]

Laboratory methods

Small batches of chlorine gas are prepared in the laboratory by combining hydrochloric acid and manganese dioxide. In a second method, a strong acid such as sulfuric acid or hydrochloric acid is reacted with sodium hypochlorite solution to release chlorine gas. In a third method, one of the same strong acids is reacted with sodium chlorate to produce chlorine gas and chlorine dioxide gas.

Applications

Production of industrial and consumer products

Chlorine is primarily used in the production of a wide range of industrial and consumer products and materials,[52][53] including plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals, pharmaceuticals, insecticides, dyestuffs, and household cleaning products.

Many important industrial products are produced via organochlorine intermediates. Examples include polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose, and propylene oxide. Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organic compounds. When applied to organic substrates, reaction is often—but not invariably—non-regioselective, and, hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions are common. If the different reaction products are easily separated (by distillation or the like), substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride, chloroform, and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene, and tetrachloroethylene from 1,2-dichloroethane.

Quantitatively, of all elemental chlorine produced, about 63% is used in the manufacture of organic compounds, and 18% in the manufacture of inorganic chlorine compounds.[48] About 15,000 chlorine compounds are used commercially.[32] The remaining 19% of chlorine produced is used for bleaches and disinfection products.[48] The most significant of organic compounds in terms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes, and trichlorobenzenes. The major inorganic compounds include HCl, Cl2O, HOCl, NaClO3, chlorinated isocyanurates, AlCl3, SiCl4, SnCl4, PCl3, PCl5, POCl3, AsCl3, SbCl3, SbCl5, BiCl3, S2Cl2, SCl2, SOCI2, ClF3, ICl, ICl3, TiCl3, TiCl4, MoCl5, FeCl3, ZnCl2, etc.[48][54]

In the past, pulp bleaching was often done with elemental chlorine, producing organochlorine pollution; today, that process is prohibited by environmental laws. Chlorine is used either in chlorine dioxide and sodium hypochlorite stages in elemental chlorine free (ECF) bleaching, or not at all (total chlorine free or TCF bleaching).

Sanitation, disinfection, and antisepsis

Main articles: Water chlorination, Bleach, and Antoine Labarraque

Combating putrefaction

Antoine-Germain Labarraque

In France (as elsewhere), animal intestines were processed to make musical instrument strings, Goldbeater's skin and other products. This was done in "gut factories" (boyauderies), and it was an odiferous and unhealthy process. In or about 1820, the Société d'encouragement pour l'industrie nationale offered a prize for the discovery of a method, chemical or mechanical, for separating the peritoneal membrane of animal intestines without putrefaction.[55][56] The prize was won by Antoine-Germain Labarraque, a 44-year-old French chemist and pharmacist who had discovered that Berthollet's chlorinated bleaching solutions ("Eau de Javel") not only destroyed the smell of putrefaction of animal tissue decomposition, but also actually retarded the decomposition.[56][57]

Labarraque's research resulted in the use of chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) in the boyauderies. The same chemicals were found to be useful in the routine disinfection and deodorization of latrines, sewers, markets, abattoirs, anatomical theatres, and morgues.[58] They were successful in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and they were beneficial during exhumations,[59] embalming, outbreaks of epidemic disease, fever, and blackleg in cattle.[55]

Against infection and contagion

Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", presumed to be transmitted by "miasmas"), and to treat putrefaction of existing wounds, including septic wounds.[60] In his 1828 work, Labarraque recommended that doctors breathe chlorine, wash their hands in chlorinated lime, and even sprinkle chlorinated lime about the patients' beds in cases of "contagious infection". In 1828, the contagion of infections was well known, even though the agency of the microbe was not discovered until more than half a century later.

During the Paris cholera outbreak of 1832, large quantities of so-called chloride of lime were used to disinfect the capital. This was not simply modern calcium chloride, but chlorine gas dissolved in lime-water (dilute calcium hydroxide) to form calcium hypochlorite (chlorinated lime). Labarraque's discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and by doing so, effectively deodorised the Latin Quarter of Paris.[61] These "putrid miasmas" were thought by many to cause the spread of "contagion" and "infection" – both words used before the germ theory of infection. Chloride of lime was used for destroying odors and "putrid matter". One source claims chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well that was feeding the Broad Street pump in 1854 London,[62] though three other reputable sources that describe that famous cholera epidemic do not mention the incident.[63][64][65] One reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump—a common practice in mid-nineteenth century England.[63]: 296 

Semmelweis and experiments with antisepsis

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine and chemical base solutions was in 1847, when Ignaz Semmelweis used chlorine-water (chlorine dissolved in pure water, which was cheaper chlorinated lime solutions) to disinfect the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Long before the germ theory of disease, Semmelweis theorized that "cadaveric particles" were transmitting decay from fresh medical cadavers to living patients, and he used the well-known "Labarraque's solutions" as the only known method to remove the smell of decay and tissue decomposition (which he found that soap did not). The solutions proved to be far more effective antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis's celebrated success in stopping the transmission of childbed fever ("puerperal fever") in the maternity wards of Vienna General Hospital in Austria in 1847.[66]

Much later, during World War I in 1916, a standardized and diluted modification of Labarraque's solution containing hypochlorite (0.5%) and boric acid as an acidic stabilizer, was developed by Henry Drysdale Dakin (who gave full credit to Labarraque's prior work in this area). Called Dakin's solution, the method of wound irrigation with chlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modern antibiotic era. A modified version of this solution continues to be employed in wound irrigation in modern times, where it remains effective against bacteria that are resistant to multiple antibiotics (see Century Pharmaceuticals).

Public sanitation

Liquid Pool Chlorine

By 1918, the US Department of Treasury called for all drinking water to be disinfected with chlorine. Chlorine is presently an important chemical for water purification (such as in water treatment plants), in disinfectants, and in bleach. As a disinfectant in water, chlorine is more than three times as effective against Escherichia coli as bromine, and more than six times as effective as iodine.[67]

Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather sodium hypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. The drawback of using chlorine in swimming pools is that the chlorine reacts with the proteins in human hair and skin (see Hypochlorous acid), and becomes chemically bonded. Even small water supplies are now routinely chlorinated.[13]

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include hypochlorite solutions, which gradually release chlorine into the water, and compounds like sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred to as "dichlor", and trichloro-s-triazinetrione, sometimes referred to as "trichlor". These compounds are stable while solid and may be used in powdered, granular, or tablet form. When added in small amounts to pool water or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forming hypochlorous acid (HOCl), which acts as a general biocide, killing germs, micro-organisms, algae, and so on.[68][69]

Use as a weapon

World War I

Main article: Chemical weapons in World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres.[70][71] As described by the soldiers, it had the distinctive smell of a mixture of pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine reacts with water in the mucosa of the lungs to form hydrochloric acid, destructive to living tissue and potentially lethal. Human respiratory systems can be protected from chlorine gas by gas masks with activated charcoal or other filters, which makes chlorine gas much less lethal than other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, which developed methods for discharging chlorine gas against an entrenched enemy. Some say that Haber's role in the development of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide.[72] After its first use, both sides in the conflict used chlorine as a chemical weapon, but it was soon replaced by the more deadly phosgene and mustard gas.[73] Theodore Gray wrote in his book The Elements: A Visual Exploration of Every Atom in the Universe, "Chlorine was used as a poison gas during the grueling trench-warfare phase. Soldiers would position a line of gas cylinders at the front lines, wait for the wind to shift towards the enemy, then open the valves and run like hell. This practice---sometimes overseen personally by Fritz Haber, a man whose positive contributions to humanity are listed under nitrogen (7)—was slowly phased out as experience showed that roughly equal numbers of soldiers on both sides died regardless of who set off the gas."[74]

Iraq and Syria

Main article: Chlorine bombings in Iraq

Chlorine gas was also used during the Iraq War in Anbar Province in 2007, with insurgents packing truck bombs with mortar shells and chlorine tanks. The attacks killed two people from the explosives, and sickened more than 350. Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. In some bombings, over a hundred civilians were hospitalized due to breathing difficulties. The Iraqi authorities tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.[75][76]

Syrian Civil War

The Chemical Weapons Convention, signed in 1993, prohibits the use of chlorine gas in war.

There have been allegations of chlorine gas attacks during the Syrian Civil War such as the 2014 Kafr Zita chemical attack. The Organisation for the Prohibition of Chemical Weapons has run investigations into the use of chemical attacks in Syria.

Islamic State of Iraq and the Levant
See also: Islamic State of Iraq and the Levant § Use of chemical weapons

On October 24, 2014, it was reported that the Islamic State of Iraq and the Levant had used chlorine gas in the town of Duluiyah, Iraq.[77] Laboratory analysis of clothing and soil samples confirmed the use of chlorine gas against Kurdish Peshmerga Forces in a vehicle-borne improvised explosive device attack on January 23, 2015 at the Highway 47 Kiske Junction near Mosul.[78]

Health effects and hazards

NFPA 704
safety square
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
4
0
0
OX

Chlorine is a toxic gas that attacks the respiratory system, eyes, and skin.[79] Because it is denser than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.[80][81]

Chlorine is detectable with measuring devices in concentrations as low as 0.2 parts per million (ppm), and by smell at 3 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can be fatal after a few deep breaths of the gas.[32] The IDLH (immediately dangerous to life and health) concentration is 10 ppm.[82] Breathing lower concentrations can aggravate the respiratory system and exposure to the gas can irritate the eyes.[83] The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations greater than 30 ppm, it reacts with water and cellular fluid, producing hydrochloric acid (HCl) and hypochlorous acid (HClO).

When used at specified levels for water disinfection, the reaction of chlorine with water is not a major concern for human health. Other materials present in the water may generate disinfection by-products that are associated with negative effects on human health.[84][85]

In the United States, the Occupational Safety and Health Administration (OSHA) has set the permissible exposure limit for elemental chlorine at 1 ppm, or 3 mg/m3. The National Institute for Occupational Safety and Health has designated a recommended exposure limit of 0.5 ppm over 15 minutes.[82]

In the home, accidents occur when hypochlorite bleach solutions come into contact with certain acidic drain-cleaners.[86] Hyperchlorite bleach (a popular laundry additive) combined with ammonia (another popular laundry additive) produces chloramines, another toxic group of chemicals.[87]

Chlorine-induced cracking in structural materials

Chlorine "attack" on an acetal resin plumbing joint.

Chlorine is widely used for purifying water, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred from chlorine induced stress corrosion cracking of stainless steel suspension rods.[88] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic plumbing, and stress corrosion cracking caused widespread failures in the USA in the 1980s and 1990s. The picture on the right shows a fractured acetal joint in a water supply system. The cracks started at injection molding defects in the joint and slowly grew until the part failed. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.[89]

Chlorine-iron fire

The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.[90][91] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework that carries chlorine gas is made of steel.[90][91]

Organochlorine compounds as pollutants

Some organochlorine compounds, either as industrial by-products or commercial end-products, are serious pollutants that persist in the environment, including some chlorinated pesticides and chlorofluorocarbons. Chlorine is added to pesticides and pharmaceuticals to make the molecules more resistant to enzymatic degradation by bacteria, insects, and mammals, but this property has the effect of reducing biodegradability and prolonging the residence time of these compounds in the environment. In this respect, chlorinated organics resemble fluorinated organics.

See also

References

  1. ^ "Standard Atomic Weights: Chlorine". CIAAW. 2009.
  2. ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ Chlorine, Gas Encyclopaedia, Air Liquide
  4. ^ Ambrose, D; Hall, D.J; Lee, D.A; Lewis, G.B; Mash, C.J (1979). "The vapour pressure of chlorine". The Journal of Chemical Thermodynamics. 11: 1089–1094. doi:10.1016/0021-9614(79)90139-3.
  5. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
  6. ^ Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
  7. ^ 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.
  8. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
  9. ^ Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
  10. ^ Barbalace, Kenneth (1995-10-22). "Chemical Database: Chlorine". EnvironmentalChemistry.com. Retrieved 2011-09-19.
  11. ^ "Facts About Chlorine". Emergency Preparedness and Response. United States Centers for Disease Control and Prevention. Retrieved 2011-09-19.
  12. ^ "SRD134 Chlorine". NIST. Retrieved 2012-09-12.
  13. ^ a b Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics (81st ed.). CRC press. ISBN 0-8493-0481-4.
  14. ^ Greenwood 1997, p. 806.
  15. ^ a b Wiberg 2001, p. 408.
  16. ^ Wiberg 2001, p. 409.
  17. ^ Greenwood 1997, pp. 857–858.
  18. ^ Windholz, Martha; et al., eds. (1976). Merck Index of Chemicals and Drugs (9th ed.). Rahway, N.J.: Merck & Co. ISBN 0-911910-26-3.
  19. ^ Greenwood 1997, pp. 844–850.
  20. ^ Greenwood 1997, p. 849.
  21. ^ Cotton, F. Albert; Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry (2nd ed.). John Wiley & sons. p. 568. {{cite book}}: Unknown parameter |last-author-amp= ignored (|name-list-style= suggested) (help)
  22. ^ Emeléus, H. J (1961). Advances in inorganic chemistry and radiochemistry. pp. 133–143. ISBN 978-0-12-023603-9.
  23. ^ Allsopp, M. W. and Vianello, G. (2012) "Poly(Vinyl Chloride)" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim. doi:10.1002/14356007.a21_717.
  24. ^ Cameron, A.G.W. (June 1957). "Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis" (PDF). CRL-41.
  25. ^ Greenwood 1997, p. 795.
  26. ^ "Risk assessment and the cycling of natural organochlorines" (PDF). Euro Chlor. Archived from the original (PDF) on 2011-07-26. Retrieved 2007-08-12.
  27. ^ a b Georges, Audi; Bersillon, O.; Blachot, J.; Wapstra, A.H. (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. 729. Atomic Mass Data Center: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001.
  28. ^ Gamow, George (1967). One Two Three... Infinity: Facts and Speculations of Science (Bantam Science and Mathematics ed.). Bantam. pp. 151–154.
  29. ^ "The earliest salt production in the world: an early Neolithic exploitation in Poiana Slatinei-Lunca, Romania". Archived from the original on April 30, 2011. Retrieved 2008-07-10. {{cite web}}: Unknown parameter |deadurl= ignored (|url-status= suggested) (help)
  30. ^ a b c Greenwood 1997, p. 790.
  31. ^ a b c d e f "17 Chlorine". Elements.vanderkrogt.net. Retrieved 2008-09-12.
  32. ^ a b c Greenwood 1997, p. 793.
  33. ^ Greenwood 1997, p. 792.
  34. ^ Ihde, Aaron John (1984). The development of modern chemistry. Courier Dover Publications. p. 158. ISBN 0-486-64235-6.
  35. ^ Weeks, Mary Elvira (1932). "The discovery of the elements. XVII. The halogen family". Journal of Chemical Education. 9 (11): 1915. Bibcode:1932JChEd...9.1915W. doi:10.1021/ed009p1915.
  36. ^ Gay-Lussac, Joseph Louis; Thénard, Louis-Jacques (1809). "On the nature and the properties of muriatic acid and of oxygenated muriatic acid". Mémoires de Physique et de Chimie de la Société d'Arcueil. 2: 339–358. In fact, oxygenated muriatic acid is not decomposed by charcoal, and it might be supposed, from this fact and those that are communicated in this Memoir, that this gas is a simple body. The phenomena that it presents can be explained well enough on this hypothesis; we shall not seek to defend it, however, as it appears to us that they are still better explained by regarding oxygenated muriatic acid as a compound body.
  37. ^ Sir Humphry Davy (1811). "On a Combination of Oxymuriatic Gas and Oxygene Gas". Philosophical Transactions of the Royal Society. 101: 155–162. doi:10.1098/rstl.1811.0008.
  38. ^ Greenwood 1997, p. 789.
  39. ^ Snelders, H. A. M. (1971). "J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter". Isis. 62 (3): 328. doi:10.1086/350763. JSTOR 229946.
  40. ^ Chodos, Alan (ed.). "This Month in Physics History September 4, 1821 and August 29, 1831: Faraday and Electromagnetism". American Physical Society. Retrieved 2010-05-08.
  41. ^ O'Connor J. J.; Robertson E. F. "Michael Faraday". School of Mathematics and Statistics, University of St Andrews, Scotland. Retrieved 2010-05-08.
  42. ^ a b "Bleaching". Encyclopædia Britannica (9th Edition (1875) and 10th Edition (1902) ed.). Retrieved 2012-05-02.
  43. ^ Aspin, Chris (1981). The Cotton Industry. Shire Publications Ltd. p. 24. ISBN 0-85263-545-1.Template:Inconsistent citations{{cite book}}: CS1 maint: postscript (link)
  44. ^ "How Products Are Made". May 2011.
  45. ^ "Chlorine – History" (PDF). Archived from the original (PDF) on 21 February 2007. Retrieved 2008-07-10. {{cite web}}: Unknown parameter |deadurl= ignored (|url-status= suggested) (help)
  46. ^ "Weaponry: Use of Chlorine Gas Cylinders in World War I". historynet.com. Retrieved 2008-07-10.
  47. ^ a b c Jacqueline Brazin. "Chlorine & its Consequences" (PDF). Archived from the original (PDF) on September 18, 2006. Retrieved 2008-07-10. {{cite web}}: Unknown parameter |deadurl= ignored (|url-status= suggested) (help)
  48. ^ a b c d Greenwood 1997, p. 798.
  49. ^ "The diaphragm cell process". Euro Chlor. Retrieved 2007-08-15.
  50. ^ "The membrane cell process". Euro Chlor. Retrieved 2007-08-15.
  51. ^ Peter Schmittinger et al. "Chlorine," Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH Verlag GmbH & Co, 2006, doi:10.1002/14356007.a06_399.pub2
  52. ^ "What is chlorine used for?". Euro Chlor. October 2010. Archived from the original on 2011-07-26. Retrieved 2007-08-20.
  53. ^ "Chlorine Tree". Chlorine Tree. Retrieved 2007-08-20.
  54. ^ Wiberg 2001, p. 412.
  55. ^ a b Hoefer, Jean Chrétien Ferdinand (ed.). "Labarraque, Antoine-Germain". Nouvelle biographie universelle. Vol. 28. pp. 323–324. OL 24229911M.
  56. ^ a b Knight, Charles (1867). Arts and sciences. Vol. 1. Bradbury, Evans & Co. p. 427.
  57. ^ Bouvet, Maurice (1950). "Les grands pharmaciens: Labarraque (1777–1850)". Revue d'histoire de la pharmacie (in French). 38 (128): 97–107. doi:10.3406/pharm.1950.8662. {{cite journal}}: Unknown parameter |trans_title= ignored (|trans-title= suggested) (help)
  58. ^ Gédéon, Andras (2006). Science and technology in medicine. Springer. pp. 181–2. ISBN 9780387278742.
  59. ^ Labarraque, Antoine Germain (1828). On the disinfecting properties of Labarraque's preparations of chlorine. Translated by James Scott. p. 8.
  60. ^ Scott, James (trans.). On the disinfecting properties of Labarraque's preparations of chlorine (S. Highley, 1828) Accessed Nov 1, 2011.
  61. ^ Corbin, Alain. The Foul and the Fragrant: Odor and the French Social Imagination (Harvard University Press, 1988) pp. 121–2.
  62. ^ Black and Veatch Corp. (2010). White's Handbook of Chlorination and Alternative Disinfectants. Hoboken, NJ:Wiley. first page, Chapter 9.
  63. ^ a b Vinten-Johansen, Peter, Howard Brody, Nigel Paneth, Stephen Rachman and Michael Rip. (2003). Cholera, Chloroform, and the Science of Medicine. New York:Oxford University.
  64. ^ Hemphill, Sandra. (2007). The Strange Case of the Broad Street Pump: John Snow and the Mystery of Cholera. Los Angeles:University of California
  65. ^ Johnson, Steven. (2006). The Ghost Map: The Story of London's Most Terrifying Epidemic and How It Changed Science, Cities, and the Modern World. New York :Riverhead Books
  66. ^ "Chlorine Story". americanchemistry. Retrieved 2008-07-10.
  67. ^ Koski T. A.; Stuart L. S.; Ortenzio L. F. (1966). "Comparison of chlorine, bromine, iodine as disinfectants for swimming pool water". Applied Microbiology. 14 (2): 276–279. PMC 546668. PMID 4959984.
  68. ^ Greenwood 1997, p. 860.
  69. ^ Wiberg 2001, p. 411.
  70. ^ "Battle of Ypres" The Canadian Encyclopedia
  71. ^ Everts, Sarah (February 23, 2015). "When Chemicals Became Weapons of War". Chemical & Engineering News. 93 (8).
  72. ^ Smil, Vaclav (2004-04-01). Enriching the Earth: Fritz Haber, Carl Bosch, and the Transformation of World Food Production. p. 226. ISBN 978-0-262-69313-4.
  73. ^ "Weapons of War: Poison Gas". First World War.com. Retrieved 2007-08-12.
  74. ^ "The Elements". google.com. Retrieved 25 August 2015.
  75. ^ Mahdi, Basim (2007-03-17). "Iraq gas attack makes hundreds ill". CNN. Retrieved 2007-03-17.
  76. ^ "'Chlorine bomb' hits Iraq village". BBC News. 2007-05-17. Retrieved 2007-05-17.
  77. ^ "ISIS chlorine attack on security forces confirmed – Iraq officials". RT. October 24, 2014.
  78. ^ "Lab report on chlorine gas usage" (PDF). Kurdistan Region Security Council. March 14, 2015.
  79. ^ "CDC | Facts About Chlorine". www.bt.cdc.gov. Retrieved 2016-04-12.
  80. ^ "Chlorine MSDS" (PDF). 1997-10-23. Archived from the original (PDF) on 2007-09-26.
  81. ^ NOAA Office of Response and Restoration, US GOV. "CHLORINE". noaa.gov. Retrieved 25 August 2015.
  82. ^ a b NIOSH Pocket Guide to Chemical Hazards. "#0115". National Institute for Occupational Safety and Health (NIOSH).
  83. ^ Winder, Chris (2001). "The Toxicology of Chlorine". Environmental Research. 85 (2): 105–114. Bibcode:2001ER.....85..105W. doi:10.1006/enrs.2000.4110. PMID 11161660.
  84. ^ "What's in your Water?: Disinfectants Create Toxic By-products". ACES News. College of Agricultural, Consumer and Environmental Sciences – University of Illinois at Urbana-Champaign. 2009-03-31. Retrieved 2009-03-31.
  85. ^ Richardson, Susan D.; Plewa, Michael J.; Wagner, Elizabeth D.; Schoeny, Rita; DeMarini, David M. (2007). "Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: A review and roadmap for research". Mutation Research/Reviews in Mutation Research. 636 (1–3): 178–242. doi:10.1016/j.mrrev.2007.09.001. PMID 17980649.
  86. ^ Berezow, Alex. "Why You Should Never Mix Different Drain Cleaners". Forbes. Retrieved 2016-04-12.
  87. ^ "Bleach Mixing Dangers :: Washington State Dept. of Health". www.doh.wa.gov. Retrieved 2016-04-12.
  88. ^ Bertolini, Luca; Elsener, Bernhard; Pedeferri, Pietro; Polder, Rob B. (2004). Corrosion of steel in concrete: prevention, diagnosis, repair. Wiley-VCH. p. 148. ISBN 3-527-30800-8.
  89. ^ Lewis, P.R. (1 January 2000). Polymer Product Failure. iSmithers Rapra Publishing. pp. 19–. ISBN 978-1-85957-192-7. Retrieved 2011-04-30.
  90. ^ a b "Chlorine: Product Datasheet" (PDF). Bayer MaterialScience AG. 2008-04-21. Archived from the original (PDF) on September 15, 2012. Retrieved 2013-12-17. {{cite web}}: Unknown parameter |deadurl= ignored (|url-status= suggested) (help)
  91. ^ a b Sanders, Roy E. (2004). Chemical Process Safety: Learning from Case Histories, 3rd Revised edition. Oxford: Elsevier Science & Technology. p. 92. ISBN 9780750677493.

Bibliography

  • Greenwood, Norman N; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0-08-037941-9. {{cite book}}: Invalid |ref=harv (help)
  • Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. Academic Press. ISBN 0-12-352651-5. {{cite book}}: Invalid |ref=harv (help)
Retrieved from "https://en.wikipedia.org/w/index.php?title=Chlorine&oldid=744612777"