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From Wikipedia, the free encyclopedia
Aluminium chloride
Aluminium chloride
Aluminium trichloride dimer
IUPAC name
aluminium chloride
Other names
aluminium(III) chloride
aluminum trichloride
CAS Registry Number
7446-70-0 (anhydrous) Yes
10124-27-3 (hydrate)
7784-13-6 (hexahydrate)
ATC code D10AX01
ChEBI CHEBI:30114 Yes
ChemSpider 22445 Yes
Jmol-3D images Image
PubChem 24012
RTECS number BD0530000
Chemical formula
Molar mass 133.34 g/mol (anhydrous)
241.43 g/mol (hexahydrate)
Appearance white or pale yellow solid,
Density 2.48 g/cm3 (anhydrous)
1.3 g/cm3 (hexahydrate)
Melting point 192.4 °C (378.3 °F; 465.5 K)
100 °C (212 °F; 373 K)
180 °C (356 °F; 453 K)
Boiling point 120 °C (248 °F; 393 K) (hexahydrate)
Solubility in water
43.9 g/100 ml (0 °C)
44.9 g/100 ml (10 °C)
45.8 g/100 ml (20 °C)
46.6 g/100 ml (30 °C)
47.3 g/100 ml (40 °C)
48.1 g/100 ml (60 °C)
48.6 g/100 ml (80 °C)
49 g/100 ml (100 °C)
Solubility soluble in hydrogen chloride, ethanol, chloroform, carbon tetrachloride
slightly soluble in benzene
Vapor pressure 133.3 Pa (99 °C)
13.3 kPa (151 °C)[1]
Viscosity 0.35 cP (197 °C)
0.26 cP (237 °C)[1]
Crystal structure
Monoclinic, mS16
Space group
C12/m1, No. 12
Coordination geometry
Octahedral (solid)
Tetrahedral (liquid)
Molecular shape
Trigonal planar
(monomeric vapour)
heat capacity (C)
91 J/mol·K[1]
Std molar
entropy (So298)
111 J/mol·K[2]
Std enthalpy of
formation (ΔfHo298)
−704.2 kJ/mol[1][2]
Gibbs free energy (ΔfG˚)
-628.6 kJ/mol[1]
Safety data sheet See: data page
GHS pictograms The corrosion pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[3]
GHS signal word Danger
GHS hazard statements
GHS precautionary statements
P280, P310, P305+351+338[3]
EU classification Corrosive C
R-phrases R34
S-phrases (S1/2), S7/8, S28, S45
NFPA 704
NFPA 704 four-colored diamond
Lethal dose or concentration (LD, LC):
LD50 (Median dose)
380 mg/kg, rat (oral)
3311 mg/kg, rat (oral)
US health exposure limits (NIOSH):
PEL (Permissible)
REL (Recommended)
2 mg/m3[4]
IDLH (Immediate danger
Related compounds
Other anions
Aluminium fluoride
Aluminium bromide
Aluminium iodide
Other cations
Boron trichloride
Gallium trichloride
Indium(III) chloride
Magnesium chloride
Related Lewis acids
Iron(III) chloride
Boron trifluoride
Supplementary data page
Structure and
Refractive index (n),
Dielectric constant (εr), etc.
Phase behaviour
Spectral data
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
verify (what is: Yes/?)
Infobox references
Aluminium chloride (AlCl3) is the main compound of aluminium and chlorine. It is white, but samples are often contaminated with iron trichloride, giving it a yellow colour. The solid has a low melting and boiling point. It is mainly produced and consumed in the production of aluminium metal, but large amounts are also used in other areas of chemical industry. The compound is often cited as a Lewis acid. It is an example of an inorganic compound that “cracks” at mild temperature, reversibly changing from a polymer to a monomer.

Contents [hide]
1 Structure
2 Reactions
2.1 Reactions with water
3 Synthesis
4 Uses
4.1 Anhydrous aluminium trichloride
4.2 Hydrated aluminium chlorides
5 Symmetry and Dipole moment
6 Safety
7 References
8 External links
AlCl3 adopts three different structures, depending on the temperature and the state (solid, liquid, gas). Solid AlCl3 is a sheet-like layered cubic close packed layers. In this framework, the Al centres exhibit octahedral coordination geometry.[5] In the melt, aluminium trichloride exists as the dimer Al2Cl6, with tetracoordinate aluminium. This change in structure is related to the lower density of the liquid phase (1.78 g/cm3) vs solid aluminium trichloride (2.48 g/cm3). Al2Cl6 dimers are also found in the vapour phase. At higher temperatures, the Al2Cl6 dimers dissociate into trigonal planar AlCl3, which is structurally analogous to BF3. The melt conducts electricity poorly,[6] unlike more ionic halides such as sodium chloride.

Anhydrous aluminium chloride is a powerful Lewis acid, capable of forming Lewis acid-base adducts with even weak Lewis bases such as benzophenone and mesitylene.[7] It forms tetrachloroaluminate AlCl4− in the presence of chloride ions.

Aluminium chloride reacts with calcium and magnesium hydrides in tetrahydrofuran forming tetrahydroaluminates.

Reactions with water[edit]
Aluminium chloride is hygroscopic, having a very pronounced affinity for water. It fumes in moist air and hisses when mixed with liquid water as the Cl− ions are displaced with H2O molecules in the lattice to form the hexahydrate AlCl3·6H2O (also white to yellowish in color). The anhydrous phase cannot be regained on heating as HCl is lost leaving aluminium hydroxide or alumina (aluminium oxide):

Al(H2O)6Cl3 → Al(OH)3 + 3 HCl + 3 H2O
On strong heating (~400°C), the aluminium oxide is formed from the aluminium hydroxide via:

2 Al(OH)3 → Al2O3 + 3 H2O
Aqueous solutions of AlCl3 are ionic and thus conduct electricity well. Such solutions are found to be acidic, indicative of partial hydrolysis of the Al3+ ion. The reactions can be described (simplified) as:

[Al(H2O)6]3+ ⇌ [Al(OH)(H2O)5]2+ + H+
Aqueous solutions behave similarly to other aluminium salts containing hydrated Al3+ ions, giving a gelatinous precipitate of aluminium hydroxide upon reaction with dilute sodium hydroxide:

AlCl3 + 3 NaOH → Al(OH)3 + 3 NaCl
Aluminium chloride is manufactured on a large scale by the exothermic reaction of aluminium metal with chlorine or hydrogen chloride at temperatures between 650 to 750 °C (1,202 to 1,382 °F).[6]

2 Al + 3 Cl2 → 2 AlCl3
2 Al + 6 HCl → 2 AlCl3 + 3 H2
Aluminum chloride may be formed via a single displacement reaction between copper chloride and aluminum metal.

2Al +3 CuCl2 →2AlCl3 + 3Cu
In the US in 1993, approximately 21,000 tons were produced, not counting the amounts consumed in the production of aluminium.[8]

Hydrated aluminium trichloride is prepared by dissolving aluminium oxides in hydrochloric acid. Metallic aluminum also readily dissolves in Hydrochloric acid ─ releasing hydrogen gas and generating considerable heat. Heating this solid does not produce anhydrous aluminium trichloride, the hexahydrate decomposes to aluminium oxide when heated to 300 °C (572 °F):[8]

2 AlCl3 + 3 H2O → Al2O3 + 6 HCl
Aluminium also forms a lower chloride, aluminium(I) chloride (AlCl), but this is very unstable and only known in the vapour phase.[6]

Anhydrous aluminium trichloride[edit]
AlCl3 is probably the most commonly used Lewis acid and also one of the most powerful. It finds application in the chemical industry as a catalyst for Friedel–Crafts reactions, both acylations and alkylations. Important products are detergents and ethylbenzene. It also finds use in polymerization and isomerization reactions of hydrocarbons.

The Friedel–Crafts reaction[7] is the major use for aluminium chloride, for example in the preparation of anthraquinone (for the dyestuffs industry) from benzene and phosgene.[6] In the general Friedel–Crafts reaction, an acyl chloride or alkyl halide reacts with an aromatic system as shown:[7]

Benzene Friedel-Crafts alkylation-diagram.svg
The alkylation reaction is more widely used than the acylation reaction, although its practice is more technically demanding because the reaction is more sluggish. For both reactions, the aluminium chloride, as well as other materials and the equipment, should be dry, although a trace of moisture is necessary for the reaction to proceed.[citation needed] A general problem with the Friedel–Crafts reaction is that the aluminium chloride catalyst sometimes is required in full stoichiometric quantities, because it complexes strongly with the products. This complication sometimes generates a large amount of corrosive waste. For these and similar reasons, more recyclable or environmentally benign catalysts have been sought. Thus, the use of aluminium trichloride in some applications is being displaced by zeolites.

Aluminium chloride can also be used to introduce aldehyde groups onto aromatic rings, for example via the Gattermann-Koch reaction which uses carbon monoxide, hydrogen chloride and a copper(I) chloride co-catalyst.[9]

AlCl3 formylation.gif
Aluminium chloride finds a wide variety of other applications in organic chemistry.[10] For example, it can catalyse the “ene reaction”, such as the addition of 3-buten-2-one (methyl vinyl ketone) to carvone:[11]

AlCl3 ene rxn.gif
AlCl3 is also widely used for polymerization and isomerization reactions of hydrocarbons. Important examples include the manufacture of ethylbenzene, which used to make styrene and thus polystyrene, and also production of dodecylbenzene, which is used for making detergents.[6]

Aluminium chloride combined with aluminium in the presence of an arene can be used to synthesize bis(arene) metal complexes, e.g. bis(benzene)chromium, from certain metal halides via the so-called Fischer-Hafner synthesis.

Hydrated aluminium chlorides[edit]
The hexahydrate has few applications, but aluminium chlorohydrate is a common component in antiperspirants at low concentrations.[8] Hyperhidrosis sufferers need a much higher concentration (12% or higher), sold under such brand names as Xeransis, Drysol, DryDerm, sunsola, Maxim, Odaban, CertainDri, B+Drier, Chlorhydrol, Anhydrol Forte and Driclor.

Symmetry and Dipole moment[edit]
Aluminium chloride belongs to the point group D3h in its monomeric form and D2h in its dimeric form. Both forms of aluminium chloride, however, do not possess a dipole moment because the bond dipole moments cancel each other out.

Anhydrous AlCl3 reacts vigorously with bases, so suitable precautions are required. It can cause irritation to the eyes, skin, and the respiratory system if inhaled or on contact.[12]

Aluminum chloride has been established as a neurotoxin.

From Wikipedia, the free encyclopedia
This article is about the chemical element. For the bleach, see Sodium hypochlorite. For the film, see Chlorine (film).
“Cl” and “Cl2” redirect here. For other uses, see CL and CL2.
Chlorine, 17Cl
Chlorine ampoule.jpg
A glass container filled with chlorine gas
Chlorine spectrum visible.png
Emission line spectra; 400–700 nm
General properties
Name, symbol chlorine, Cl
Appearance pale yellow-green gas
Pronunciation /ˈklɔəriːn/ or /ˈklɔərɨn/
klohr-een or klohr-ən
Chlorine in the periodic table
Hydrogen (diatomic nonmetal)
Helium (noble gas)
Lithium (alkali metal)
Beryllium (alkaline earth metal)
Boron (metalloid)
Carbon (polyatomic nonmetal)
Nitrogen (diatomic nonmetal)
Oxygen (diatomic nonmetal)
Fluorine (diatomic nonmetal)
Neon (noble gas)
Sodium (alkali metal)
Magnesium (alkaline earth metal)
Aluminium (post-transition metal)
Silicon (metalloid)
Phosphorus (polyatomic nonmetal)
Sulfur (polyatomic nonmetal)
Chlorine (diatomic nonmetal)
Argon (noble gas)
Potassium (alkali metal)
Calcium (alkaline earth metal)
Scandium (transition metal)
Titanium (transition metal)
Vanadium (transition metal)
Chromium (transition metal)
Manganese (transition metal)
Iron (transition metal)
Cobalt (transition metal)
Nickel (transition metal)
Copper (transition metal)
Zinc (transition metal)
Gallium (post-transition metal)
Germanium (metalloid)
Arsenic (metalloid)
Selenium (polyatomic nonmetal)
Bromine (diatomic nonmetal)
Krypton (noble gas)
Rubidium (alkali metal)
Strontium (alkaline earth metal)
Yttrium (transition metal)
Zirconium (transition metal)
Niobium (transition metal)
Molybdenum (transition metal)
Technetium (transition metal)
Ruthenium (transition metal)
Rhodium (transition metal)
Palladium (transition metal)
Silver (transition metal)
Cadmium (transition metal)
Indium (post-transition metal)
Tin (post-transition metal)
Antimony (metalloid)
Tellurium (metalloid)
Iodine (diatomic nonmetal)
Xenon (noble gas)
Caesium (alkali metal)
Barium (alkaline earth metal)
Lanthanum (lanthanide)
Cerium (lanthanide)
Praseodymium (lanthanide)
Neodymium (lanthanide)
Promethium (lanthanide)
Samarium (lanthanide)
Europium (lanthanide)
Gadolinium (lanthanide)
Terbium (lanthanide)
Dysprosium (lanthanide)
Holmium (lanthanide)
Erbium (lanthanide)
Thulium (lanthanide)
Ytterbium (lanthanide)
Lutetium (lanthanide)
Hafnium (transition metal)
Tantalum (transition metal)
Tungsten (transition metal)
Rhenium (transition metal)
Osmium (transition metal)
Iridium (transition metal)
Platinum (transition metal)
Gold (transition metal)
Mercury (transition metal)
Thallium (post-transition metal)
Lead (post-transition metal)
Bismuth (post-transition metal)
Polonium (post-transition metal)
Astatine (metalloid)
Radon (noble gas)
Francium (alkali metal)
Radium (alkaline earth metal)
Actinium (actinide)
Thorium (actinide)
Protactinium (actinide)
Uranium (actinide)
Neptunium (actinide)
Plutonium (actinide)
Americium (actinide)
Curium (actinide)
Berkelium (actinide)
Californium (actinide)
Einsteinium (actinide)
Fermium (actinide)
Mendelevium (actinide)
Nobelium (actinide)
Lawrencium (actinide)
Rutherfordium (transition metal)
Dubnium (transition metal)
Seaborgium (transition metal)
Bohrium (transition metal)
Hassium (transition metal)
Meitnerium (unknown chemical properties)
Darmstadtium (unknown chemical properties)
Roentgenium (unknown chemical properties)
Copernicium (transition metal)
Ununtrium (unknown chemical properties)
Flerovium (post-transition metal)
Ununpentium (unknown chemical properties)
Livermorium (unknown chemical properties)
Ununseptium (unknown chemical properties)
Ununoctium (unknown chemical properties)


sulfur ← chlorine → argon
Atomic number 17
Standard atomic weight (Ar) 35.45[1] (35.446–35.457)[2]
Element category diatomic nonmetal
Group, block group 17 (halogens), p-block
Period period 3
Electron configuration [Ne] 3s2 3p5
per shell
2, 8, 7
Physical properties
Phase gas
Melting point 171.6 K ​(−101.5 °C, ​−150.7 °F)
Boiling point 239.11 K ​(−34.04 °C, ​−29.27 °F)
Density at stp (0 °C and 101.325 kPa) 3.2 g/L
when liquid, at b.p. 1.5625 g/cm3[3]
Critical point 416.9 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 states 7, 6, 5, 4, 3, 2, 1, −1 ​(a strongly acidic oxide)
Electronegativity Pauling scale: 3.16
Ionization energies 1st: 1251.2 kJ/mol
2nd: 2298 kJ/mol
3rd: 3822 kJ/mol
Covalent radius 102±4 pm
Van der Waals radius 175 pm
Crystal structure ​orthorhombic Orthorhombic crystal structure for chlorine
Speed of sound 206 m/s (gas, at 0 °C)
Thermal conductivity 8.9×10−3 W/(m·K)
Electrical resistivity >10 Ω·m (at 20 °C)
Magnetic ordering diamagnetic[4]
CAS Registry Number 7782-50-5
Discovery and first isolation Carl Wilhelm Scheele (1774)
Recognized as an element by Humphry Davy (1808)
Most stable isotopes
Main article: Isotopes of chlorine
iso NA half-life DM DE (MeV) DP
35Cl 75.77% 35Cl is stable with 18 neutrons
36Cl trace 3.01×105 y β− 0.709 36Ar
ε – 36S
37Cl 24.23% 37Cl is stable with 20 neutrons
view talk edit · references
Chlorine is a chemical element with symbol Cl and atomic number 17. Chlorine is in the halogen group (17) and is the second lightest halogen following fluorine. The element is a yellow-green gas under standard conditions, where it forms diatomic molecules. 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. Characterization of chlorine gas was made in 1774 by Carl Wilhelm Scheele, who supposed 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: χλωρóς (khlôros) “pale green”.

Nearly all chlorine in the Earth’s crust occurs as chloride in various ionic compounds, including 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 commercially to free chlorine’s bleaching and disinfectant uses, as well as its many uses of an essential reagent 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, as well as 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.

Contents [hide]
1 Characteristics
1.1 Physical characteristics of chlorine and its compounds
1.2 Chemical characteristics
1.2.1 Hydrolysis of free chlorine or disproportionation in water
1.2.2 Chemistry and compounds Chlorides Chlorine oxides Interhalogen compounds Organochlorine compounds
1.3 Occurrence
1.4 Isotopes
2 History
3 Production
3.1 Laboratory methods
4 Applications
4.1 Production of industrial and consumer products
4.2 Public sanitation, disinfection, and antisepsis
4.2.1 Combating putrefaction
4.2.2 Against infection and contagion
4.2.3 Semmelweis and experiments with antisepsis
4.2.4 Public sanitation
4.3 Use as a weapon
4.3.1 World War I
4.3.2 Iraq War
4.3.3 Syrian Civil War
4.3.4 Islamic State of Iraq and the Levant (ISIL/ISIS)
5 Health effects and hazards
5.1 Chlorine induced cracking in structural materials
5.2 Chlorine-iron fire
6 Organochlorine compounds as pollutants
7 See also
8 References
9 Bibliography
10 External links
Physical characteristics of chlorine and its compounds[edit]

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.[5] This is a yellow-green gas that has a distinctive strong odor, familiar to most from common household bleach.[6] 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 can be liquefied at room temperature with pressures above 740 kPa (107 psi).[7]

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

Chemical characteristics[edit]
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 forms compounds with almost all of the elements to give compounds that are usually called chlorides. Chlorine gas reacts with most organic compounds, and will even sluggishly support the combustion of hydrocarbons.[8]

Hydrolysis of free chlorine or disproportionation in water[edit]
At 25 °C and atmospheric pressure, one liter of water dissolves 3.26 g or 1.125 L of gaseous chlorine.[9] Solutions of chlorine in water contain chlorine (Cl2), hydrochloric acid, and hypochlorous acid:

Cl2 + H2O is in equilibrium with HCl + HClO
This 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.[10]

Cl2 + 2 OH− → ClO− + Cl− + H2O
Chlorine gas only exists in a neutral or acidic solution.

Chemistry and compounds[edit]
See also: Category:Chlorine compounds.
Chlorine exists in all odd numbered oxidation states from −1 to +7, as well as the elemental state of zero and four in chlorine dioxide (see table below, and also structures in chlorite).[11] 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.[12]

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
6 dichlorine hexoxide (gas). In liquid or solid disproportionates to mix of +5 and +7 oxidation states, as ionic chloryl perchlorate [ClO
+7 perchlorates ClO−
4 perchloric acid, perchlorate salts such as magnesium perchlorate, dichlorine heptoxide
Main article: Chloride
Chlorine combines with almost all elements to give chlorides. Compounds with oxygen, nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.[13] Chloride is one of the most common anions in nature. Hydrogen chloride and its aqueous solution, hydrochloric acid, are produced on megaton scale annually both as valued intermediates but sometimes as undesirable pollutants.

Chlorine oxides[edit]
Chlorine forms a variety of oxides, as seen above: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O), dichlorine hexoxide (Cl2O6), dichlorine heptoxide (Cl2O7). The anionic derivatives of these same oxides are also well known including chlorate (ClO−
3), chlorite (ClO−
2), hypochlorite (ClO−), and perchlorate (ClO−
4). The acid derivatives of these anions are hypochlorous acid (HOCl), chloric acid (HClO3) and perchloric acid (HClO4). 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.[14] The compound “chlorine trioxide” does not occur, but rather in gas form is found 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

In hot concentrated alkali solution hypochlorite disproportionates:

2 ClO− → Cl− + ClO−
ClO− + ClO−
2 → Cl− + ClO−
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−
This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:[16]

Reaction Electrode
Cl− + 2 OH− → ClO− + H2O + 2 e− +0.89 volts
ClO− + 2 OH− → ClO−
2 + H2O + 2 e− +0.67 volts
2 + 2 OH− → ClO−
3 + H2O + 2 e− +0.33 volts
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[edit]
Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively. However, it cannot oxidize fluoride salts to fluorine. It makes 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.[17]

Organochlorine compounds[edit]
Main article: Organochloride
Chlorine is used extensively in organic chemistry in substitution and addition reactions. Chlorine often imparts many desired properties to an organic compound, in part owing to its electronegativity.

Like the other halides, chlorine undergoes electrophilic addition reactions, the most notable one being the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide.[citation needed]

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 creation in supernovae, via the r-process.[18] This chlorine provides the supply found in the Solar System.

In meteorites and on Earth, chlorine is found primarily as the chloride ion which occurs in minerals. In the Earth’s crust, chlorine is present at average concentrations of about 126 parts per million,[19] predominantly in such minerals as halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate).

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

Over 2000 naturally occurring organic chlorine compounds are known.[20]

Main article: Isotopes of chlorine
Chlorine has a wide range of isotopes. The two stable isotopes are 35Cl (75.77%) and 37Cl (24.23%).[21] Together they give chlorine an atomic weight of 35.4527 g/mol. The half-integer value for chlorine’s weight caused some confusion in the early days of chemistry, when it had been postulated that atoms were composed of even units of hydrogen (see Proust’s law), and the existence of chemical isotopes was unsuspected.[22]

Trace amounts of radioactive 36Cl exist in the environment, in a ratio of about 7×10−13 to 1 with stable isotopes. 36Cl is produced in the atmosphere by spallation of 36Ar by interactions with 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. Additionally, 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 use in other areas of the geological sciences, including dating ice and sediments.[21]

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.[23] Around 1630, chlorine was recognized as a gas by the Flemish chemist and physician Jan Baptist van Helmont.[24]

Carl Wilhelm Scheele
Elemental chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and, therefore, he is credited for its discovery.[25] 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”).[25] However, he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory).[25] He named the new element within this oxide as muriaticum.[25] Regardless of what he thought, Scheele did isolate chlorine by reacting MnO2 (as the mineral pyrolusite) with HCl:[24]

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.[26]

At the time, common chemical theory was: any 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.[27][28][29]

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).[25] 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.[30]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it is an element, and not a compound.[25] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow.[31] The name “halogen”, meaning “salt producer”, was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger. However, 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.[32][33] In 1823, Michael Faraday liquefied chlorine for the first time,[34][35] and demonstrated that what was then known as “solid chlorine” had a structure of chlorine hydrate (Cl2·H2O).[24]

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.[36][37] 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. However, 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).[36] 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.[38] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of essentially all modern elemental chlorine and sodium hydroxide production (a related low-temperature electrolysis reaction, the Hooker process, is now responsible for bleach and sodium hypochlorite production).

Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti-putrification agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This work is mainly due to Antoine-Germain Labarraque, who adapted Berthollet’s “Javel water” bleach and other chlorine preparations for the purpose (for a more complete history, see below). Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) as well as public sanitation (especially of swimming and drinking water).

In 1826, silver chloride was used to produce photographic images for the first time.[39] Chloroform was first used as an anesthetic in 1847.[39]

Polyvinyl chloride (PVC) was invented in 1912, initially without a purpose.[39]

Chlorine gas was first introduced as a weapon on April 22, 1915, at Ypres by the German Army,[40][41] and the results of this weapon were disastrous because gas masks had not been mass distributed and were tricky to get on quickly.


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 essentially all industrial chlorine gas.[42] Along with chlorine, the method yields hydrogen gas and sodium hydroxide (with sodium hydroxide actually being the most crucial of the three industrial products produced by the process). The process proceeds according to the following chemical equation:[12]

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.[43] 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 preventing mercury discharge into the environment and they are more energy efficient. Membrane cell electrolysis employ permeable membrane as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.[44] 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
Laboratory methods[edit]
Small amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid and manganese dioxide. Alternatively a strong acid such as sulfuric acid or hydrochloric acid reacts with sodium hypochlorite solution to release chlorine gas but reacts with sodium chlorate to produce chlorine gas and chlorine dioxide gas as well. In the home, accidents occur when hypochlorite bleach solutions are combined with certain acidic drain-cleaners.

Production of industrial and consumer products[edit]
Principal applications of chlorine are in the production of a wide range of industrial and consumer products.[45][46] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.

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, e.g., by distillation, 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, about 63% and 18% of all elemental chlorine produced is used in the manufacture of organic and inorganic chlorine compounds, respectively.[42] About 15,000 chlorine compounds are being used commercially.[26] The remaining 19% is used for bleaches and disinfection products.[42] 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.[42][47]

Pulp bleaching was done often with elemental chlorine in the past. However, this tends to produce organochlorine pollution, and today environmental laws make it prohibitive. 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).

Public sanitation, disinfection, and antisepsis[edit]
Main articles: Water chlorination, Bleach and Antoine Labarraque
Combating putrefaction[edit]

Antoine-Germain Labarraque
In France (as elsewhere) there was a need to process animal guts in order to make musical instrument strings, Goldbeater’s skin and other products. This was carried out in “gut factories” (boyauderies) as an odiferous and unhealthy business. In or about 1820, the Société d’encouragement pour l’industrie nationale offered a prize for the discovery of a method, chemical or mechanical, that could be used to separate the peritoneal membrane of animal intestines without causing putrefaction.[48][49] It 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 retarded the decomposition process itself.[49][50]

Labarraque’s research resulted in chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) being employed not only in the boyauderies but also for the routine disinfection and deodorisation of latrines, sewers, markets, abattoirs, anatomical theatres and morgues.[51] They were also used, with success, in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and for exhumations,[52] embalming, during outbreaks of epidemic illness, fever, blackleg in cattle, etc.[48]

Against infection and contagion[edit]
Labarraque’s chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called “contagious infection”, and presumed to be transmitted by “miasmas”) and also to treat putrefaction of existing wounds, including septic wounds.[53] In this 1828 work, Labarraque recommended for the doctor to breathe chlorine, wash his hands with chlorinated lime, and even sprinkle chlorinated lime about the patient’s bed, in cases of “contagious infection”. In 1828, it was well known that some infections were contagious, even though the agency of the microbe was not to be realized or discovered for more than half a century.

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 contained 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.[54] These “putrid miasmas” were thought by many to be responsible for the spread of “contagion” and “infection” – both words used before the germ theory of infection. The use of chloride of lime was based on destruction of odors and “putrid matter”. One source has claimed that chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well feeding the Broad Street pump in 1854 London.[55] Three reputable sources that described the famous Broad Street pump cholera epidemic do not mention Snow performing any disinfection of water from that well.[56][57][58] Instead, 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.[56]:296

Semmelweis and experiments with antisepsis[edit]

Ignaz Semmelweis
Perhaps the most famous application of Labarraque’s chlorine and chemical base solutions was in 1847, when Ignaz Semmelweis used (first) chlorine-water (simply chlorine dissolved in pure water), then cheaper chlorinated lime solutions, to deodorize the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Semmelweis, still long before the germ theory of disease, had theorized that “cadaveric particles” were somehow 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 germicide antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis’s (later) celebrated success in stopping the transmission of childbed fever (“puerperal fever”) in the maternity wards of Vienna General Hospital in Austria in 1847.[59]

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 the modern era, where it remains effective against multiply antibiotic resistant bacteria (see Century Pharmaceuticals).

Public sanitation[edit]
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. Chlorine in water is more than three times as effective as a disinfectant against Escherichia coli than an equivalent concentration of bromine, and is more than six times more effective than an equivalent concentration of iodine.[60]

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). Once the chlorine reacts with the hair and skin, it becomes chemically bonded. Even small water supplies are now routinely chlorinated.[8]

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.[61][62]

Use as a weapon[edit]
World War I[edit]
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.[63] As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form hydrochloric acid, an irritant that can be lethal. The damage done by chlorine gas can be prevented by the activated charcoal commonly found in gas masks, or other filtration methods, which makes the overall chance of death by chlorine gas much lower than those of 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, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber’s role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide.[64] After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly phosgene and mustard gas.[65] 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.”[66]

Iraq War[edit]
Main article: 2007 chlorine bombings in Iraq
Chlorine gas has also been used by insurgents against the local population and coalition forces in the Iraq War in the form of chlorine bombs. On March 17, 2007, for example, three chlorine-filled trucks were detonated in the Anbar province killing two and sickening over 350.[67] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.[68] 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. The Iraqi authorities have tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.

Syrian Civil War[edit]
There have been allegations of chlorine gas attacks during the Syrian Civil War such as the 2014 Kafr Zita chemical attack.

Islamic State of Iraq and the Levant (ISIL/ISIS)[edit]
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.[69]

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.[70]

Health effects and hazards[edit]
NFPA 704
“fire diamond”
NFPA 704 four-colored diamond
Chlorine is a toxic gas that irritates the respiratory system. 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.[71][72]

Chlorine is detectable with measuring devices in concentrations of 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.[26] The IDLH (immediately dangerous to life and health) concentration is 10 ppm.[73] Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.[74] The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations above 30 ppm, it begins to react with water and cells, which change it into 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.[75][76]

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.[73]

Chlorine induced cracking in structural materials[edit]

Chlorine “attack” on an acetal resin plumbing joint.
The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to chlorine induced stress corrosion cracking of stainless steel rods used to suspend them.[77] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic supplies, 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 finally triggered. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.[78]

Chlorine-iron fire[edit]
The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.[79][80] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework used to carry chlorine gas is made of steel.[79][80]

Organochlorine compounds as pollutants[edit]
Some organochlorine compounds are serious pollutants. These are produced either as by-products or end products of industrial processes which are persistent in the environment, such as certain chlorinated pesticides and chlorofluorocarbons. Chlorine is added both to pesticides and pharmaceuticals to make the molecules more resistant to enzymatic degradation by bacteria, insects, and mammals, but this property also has the effect of prolonging the residence time of these compounds when they enter the environment. In this respect chlorinated organics have some resemblance to fluorinated organics.

From Wikipedia, the free encyclopedia
Iron(III) chloride
Iron(III) chloride hexahydrate.jpg
IUPAC names
Iron(III) chloride
Iron trichloride
Other names
Ferric chloride
Flores martis
CAS Registry Number
7705-08-0 Yes
10025-77-1 (hexahydrate)
ChEBI CHEBI:30808 Yes
ChemSpider 22792 Yes
EC number 231-729-4
Jmol-3D images Image
PubChem 24380
RTECS number LJ9100000
0I2XIN602U (hexahydrate)
UN number 1773 (anhydrous)
2582 (aq. soln.)
Chemical formula
Molar mass 162.2 g/mol (anhydrous)
270.3 g/mol (hexahydrate)
Appearance green-black by reflected light; purple-red by transmitted light
hexahydrate: yellow solid
aq. solutions: brown
Odor slight HCl
Density 2.898 g/cm3 (anhydrous)
1.82 g/cm3 (hexahydrate)
Melting point 306 °C (583 °F; 579 K) (anhydrous)
37 °C (99 °F; 310 K) (hexahydrate)
Boiling point 315 °C (599 °F; 588 K) (anhydrous, decomposes)
280 °C (536 °F; 553 K) (hexahydrate, decomposes) partial decomposition to FeCl2 + Cl2
Solubility in water
74.4 g/100 mL (0 °C)[1]
92 g/100 mL (hexahydrate, 20 °C)
Solubility in acetone
Diethyl ether 63 g/100 ml (18 °C)
highly soluble
83 g/100 ml
highly soluble
Viscosity 40% solution: 12 cP
Crystal structure
Coordination geometry
Hazards[3][4][Note 1]
Safety data sheet ICSC 1499
GHS pictograms Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
GHS signal word DANGER
GHS hazard statements
H290, H302, H314, H318
GHS precautionary statements
P234, P260, P264, P270, P273, P280, P301+312, P301+330+331, P303+361+353, P363, P304+340, P310, P321, P305+351+338
NFPA 704
NFPA 704 four-colored diamond
Flash point Non-flammable
US health exposure limits (NIOSH):
REL (Recommended)
TWA 1 mg/m3[2]
Related compounds
Other anions
Iron(III) fluoride
Iron(III) bromide
Other cations
Iron(II) chloride
Manganese(II) chloride
Cobalt(II) chloride
Ruthenium(III) chloride
Related coagulants
Iron(II) sulfate
Polyaluminium chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
verify (what is: Yes/?)
Infobox references
Iron(III) chloride, also called ferric chloride, is an industrial scale commodity chemical compound, with the formula FeCl3 and with iron in the +3 oxidation state. The colour of iron(III) chloride crystals depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red. Anhydrous iron(III) chloride is deliquescent, forming hydrated hydrogen chloride mists in moist air. It is rarely observed in its natural form, mineral molysite, known mainly from some fumaroles.

When dissolved in water, iron(III) chloride undergoes hydrolysis and gives off heat in an exothermic reaction. The resulting brown, acidic, and corrosive solution is used as a flocculant in sewage treatment and drinking water production, and as an etchant for copper-based metals in printed circuit boards. Anhydrous iron(III) chloride is a fairly strong Lewis acid, and it is used as a catalyst in organic synthesis.

Contents [hide]
1 Nomenclature
2 Structure and properties
3 Preparation
4 Reactions
4.1 Oxidation
5 Uses
5.1 Industrial
5.2 Laboratory use
5.3 Other uses
6 Safety
7 See also
8 Notes and references
8.1 Notes
8.2 References
9 Further reading
The descriptor hydrated or anhydrous is used when referring to iron(III) chloride, to distinguish between the two common forms. The hexahydrate is usually given as the simplified empirical formula FeCl3⋅6H2O. It may also be given as trans-[Fe(H2O)4Cl2]Cl⋅2H2O and the systematic name tetraaquadichloroiron(III) chloride dihydrate, which more clearly represents its structure.

Structure and properties[edit]
Anhydrous iron(III) chloride adopts the BiI3 structure, which features octahedral Fe(III) centres interconnected by two-coordinate chloride ligands. Iron(III) chloride hexahydrate consists of trans-[Fe(H2O)4Cl2]+ cationic complexes and chloride anions, with the remaining two H2O molecules embedded within the monoclinic crystal structure.[6]

Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapour consists of the dimer Fe2Cl6 (c.f. aluminium chloride) which increasingly dissociates into the monomeric FeCl3 (D3h point group molecular symmetry) at higher temperature, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[7]

Anhydrous iron(III) chloride may be prepared by union of the elements:[8]

2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s)
Solutions of iron(III) chloride are produced industrially both from iron and from ore, in a closed-loop process.

Dissolving pure iron in a solution of iron(III) chloride
Fe(s) + 2 FeCl3(aq) → 3 FeCl2(aq)
Dissolving iron ore in hydrochloric acid
Fe3O4(s) + 8 HCl(aq) → FeCl2(aq) + 2 FeCl3(aq) + 4 H2O
Oxidation of iron (II) chloride with chlorine
2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq)
Oxidation of iron (II) chloride with oxygen
4FeCl2(aq) + O2 + 4HCl → 4FeCl3(aq) + 2H2O
Reacting Iron with hydrochloric acid, then with hydrogen peroxide. The hydrogen peroxide is the catalyst in turning iron chloride into ferric chloride
Like many other hydrated metal chlorides, hydrated iron(III) chloride can be converted to the anhydrous salt by refluxing with thionyl chloride.[9] Conversion of the hydrate to anhydrous iron(III) chloride is not accomplished by heating, as HCl and iron oxychlorides are produced.


A brown, acidic solution of iron(III) chloride
Iron(III) chloride undergoes hydrolysis to give an acidic solution. When heated with iron(III) oxide at 350 °C, iron(III) chloride gives iron oxychloride, a layered solid and intercalation host.[citation needed]

FeCl3 + Fe2O3 → 3 FeOCl
It is a moderately strong Lewis acid, forming adducts with Lewis bases such as triphenylphosphine oxide, e.g. FeCl3(OPPh3)2 where Ph = phenyl. It also reacts with other chloride salts to give the yellow tetrahedral FeCl4− ion. Salts of FeCl4− in hydrochloric acid can be extracted into diethyl ether.

Alkali metal alkoxides react to give the metal alkoxide complexes of varying complexity.[10] The compounds can be dimeric or trimeric.[11] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl3 and sodium ethoxide:[12][13]

FeCl3 + 3 [C2H5O]−Na+ → Fe(OC2H5)3 + 3 NaCl
Oxalates react rapidly with aqueous iron(III) chloride to give [Fe(C2O4)3]3−. Other carboxylate salts form complexes, e.g. citrate and tartrate.

Iron(III) chloride is a mild oxidising agent, for example, it is capable of oxidising copper(I) chloride to copper(II) chloride.

FeCl3 + CuCl → FeCl2 + CuCl2
It also reacts with iron to form iron(II) chloride:

2 FeCl3 + Fe → 3 FeCl2
Reducing agents such as hydrazine convert iron(III) chloride to complexes of iron(II).

Granulated iron(III) chloride hexahydrate
In industrial application, iron(III) chloride is used in sewage treatment and drinking water production.[14] In this application, FeCl3 in slightly basic water reacts with the hydroxide ion to form a floc of iron(III) hydroxide, or more precisely formulated as FeO(OH)−, that can remove suspended materials.

[Fe(H2O)6]3+ + 4 HO− → [Fe(H2O)2(HO)4]− + 4 H2O → [Fe(H2O)O(HO)2]− + 6 H2O
It is also used as a leaching agent in chloride hydrometallurgy,[15] for example in the production of Si from FeSi. (Silgrain process)[16]

Another important application of iron(III) chloride is etching copper in two-step redox reaction to copper(I) chloride and then to copper(II) chloride in the production of printed circuit boards.[17]

FeCl3 + Cu → FeCl2 + CuCl
FeCl3 + CuCl → FeCl2 + CuCl2
Iron(III) chloride is used as catalyst for the reaction of ethylene with chlorine, forming ethylene dichloride (1,2-dichloroethane), an important commodity chemical, which is mainly used for the industrial production of vinyl chloride, the monomer for making PVC.

H2C=CH2 + Cl2 → ClCH2CH2Cl
Laboratory use[edit]
In the laboratory iron(III) chloride is commonly employed as a Lewis acid for catalysing reactions such as chlorination of aromatic compounds and Friedel-Crafts reaction of aromatics. It is less powerful than aluminium chloride, but in some cases this mildness leads to higher yields, for example in the alkylation of benzene:

Iron(III) chloride as a catalyst
The ferric chloride test is a traditional colorimetric test for phenols, which uses a 1% iron(III) chloride solution that has been neutralised with sodium hydroxide until a slight precipitate of FeO(OH) is formed.[18] The mixture is filtered before use. The organic substance is dissolved in water, methanol or ethanol, then the neutralised iron(III) chloride solution is added—a transient or permanent coloration (usually purple, green or blue) indicates the presence of a phenol or enol.

This reaction is exploited in the Trinder spot test, which is used to indicate the presence of salicylates, particularly salicylic acid, which contains a phenolic OH group.

This test can be used to detect the presence of gamma-Hydroxybutyric acid and gamma-butyrolactone,[19] which cause it to turn red-brown.

Other uses[edit]
Used in anhydrous form as a drying reagent in certain reactions.
Used to detect the presence of phenol compounds in organic synthesis e.g.: examining purity of synthesised Aspirin.
Used in water and wastewater treatment to precipitate phosphate as iron(III) phosphate.
Used by American coin collectors to identify the dates of Buffalo nickels that are so badly worn that the date is no longer visible.
Used by Blade-smiths and Artisans in Pattern welding to etch the metal, giving it a contrasting effect, to view metal layering or imperfections.
Used to etch the widmanstatten pattern in iron meteorites.
Necessary for the etching of photogravure plates for printing photographic and fine art images in intaglio and for etching rotogravure cylinders used in the printing industry.
Used to make printed circuit boards (PCBs).
Used in veterinary practice to treat overcropping of an animal’s claws, particularly when the overcropping results in bleeding.
Reacts with cyclopentadienylmagnesium bromide in one preparation of ferrocene, a metal-sandwich complex.[20]
Sometimes used in a technique of Raku ware firing, the iron coloring a pottery piece shades of pink, brown, and orange.
Used to test the pitting and crevice corrosion resistance of stainless steels and other alloys.
Used in conjunction with NaI in acetonitrile to mildly reduce organic azides to primary amines.[21]
Used in an animal thrombosis

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