Chemistry:Iron(III) chloride

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Short description: Inorganic compound
Iron(III) chloride
Iron(III) chloride anhydrate.jpg
Iron(III) chloride (anhydrous)
Хлорид железа.jpg
Iron(III) chloride (hydrate)
Iron-trichloride-sheet-3D-polyhedra.png
Iron-trichloride-sheets-stacking-3D-polyhedra.png
Names
IUPAC names
Iron(III) chloride
Iron trichloride
Other names
  • Ferric chloride
  • Molysite
  • Flores martis
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 231-729-4
RTECS number
  • LJ9100000
UNII
UN number
  • 1773 (anhydrous)
  • 2582 (aqueous solution)
Properties
FeCl
3
Molar mass
  • 162.204 g/mol (anhydrous)
  • 270.295 g/mol (hexahydrate)[1]
Appearance Green-black by reflected light; purple-red by transmitted light; yellow solid as hexahydrate; brown as aqueous solution
Odor Slight HCl
Density
  • 2.90 g/cm3 (anhydrous)
  • 1.82 g/cm3 (hexahydrate)[1]
Melting point 307.6 °C (585.7 °F; 580.8 K) (anhydrous)
37 °C (99 °F; 310 K) (hexahydrate)[1]
Boiling point
  • 316 °C (601 °F; 589 K) (anhydrous, decomposes)[1]
  • 280 °C (536 °F; 553 K) (hexahydrate, decomposes)
 
912 g/L (anhydrous or hexahydrate, 25 °C)[1]
Solubility in
  •  
  • 630 g/L (18 °C)
  • Highly soluble
  • 830 g/L
  • Highly soluble
+13,450·10−6 cm3/mol[2]
Viscosity 12 cP (40% solution)
Hazards[4][5][Note 1]
GHS pictograms Corr. Met. 1; Skin Corr. 1C; Eye Dam. 1Acute Tox. 4 (oral)
GHS Signal word DANGER
H290, H302, H314
P234, P260, P264, P270, P273, P280, P301+312, P301+330+331, P303+361+353, P363, P304+340, P310, P321, P305+351+338, P390, P405, P406, P501
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no codeNFPA 704 four-colored diamond
0
2
0
Flash point Non-flammable
NIOSH (US health exposure limits):
REL (Recommended)
 TWA 1 mg/m3[3]
Related compounds
Other anions
Other cations
Related coagulants
Structure
Hexagonal, hR24
R3, No. 148[7]
a = 0.6065 nm, b = 0.6065 nm, c = 1.742 nm
α = 90°, β = 90°, γ = 120°
6
Octahedral
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is ☑Y☒N ?)
Infobox references

Iron(III) chloride describes the inorganic compounds with the formula FeCl
3(H2O)x. Also called ferric chloride, these compounds are some of the most important and commonplace compounds of iron. They are available both in anhydrous and in hydrated forms which are both hygroscopic. They feature iron in its +3 oxidation state. The anhydrous derivative is a Lewis acid, while all forms are mild oxidizing agent. It is used as a water cleaner and as an etchant for metals.

Electronic and optical properties

Anhydrous iron(III) chloride evaporates at relatively mild temperatures to give the bitetrahedral dimer.

All forms of ferric chloride are paramagnetic, owing to the presence of unpaired electrons residing in 3d orbitals. Although Fe(III) chloride can be octahedral or tetrahedral (or both, see structure section), all of these forms have five unpaired electrons, one per d-orbial. The high spin d5 electronic configuration requires that d-d electronic transitions are spin forbidden, in addition to violating the Laporte Rule. This double forbidden-ness results in its solutions being only pale colored. Or, stated more technically, the optical transitions are non-intense. Aqueous ferric sulfate and ferric nitrate, which contain [Fe(H
2O)
6]3+, are nearly colorless, whereas the chloride solutions are yellow. Clearly, the chloride ligands significantly influence the optical properties of the iron center.[8][9]

Structure

Iron(III) chloride can exist as an anhydrous material and a series of hydrates. Their structures are quite distinct.

Anhydrous

The anhydrous compound is a hygroscopic crystalline solid with a melting point of 307.6 °C. The colour 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 has the BiI
3 structure, with octahedral Fe(III) centres interconnected by two-coordinate chloride ligands.[7][10]

Iron(III) chloride has a relatively low melting point and boils at around 315 °C. The vapor consists of the dimer Fe
2Cl
6, much like aluminium chloride. This dimer dissociates into the monomeric FeCl
3 (with D3h point group molecular symmetry) at higher temperatures, in competition with its reversible decomposition to give iron(II) chloride and chlorine gas.[11]

Hydrates

Ferric chloride form hydrates upon exposure to water, reflecting its Lewis acidity. All hydrates exhibit deliquescence, meaning that they become liquid by absorbing moisture from the air. Hydration invariably gives derivatives of aquo complexes with the formula [FeCl
2(H
2O)
4]+
. This cation can adopt either trans or cis stereochemistry, reflecting the relative location of the chloride ligands on the octahedral Fe center. Four hydrates have been characterized by X-ray crystallography: the dihydrate FeCl
3 · 2H2O, the disesquihydrate FeCl
3 · 2.5H2O, the trisesquihydrate FeCl
3 · 3.5H2O, and finally the hexahydrate FeCl
3 · 6H2O. These species differ with respect to the stereochemistry of the octahedral iron cation, the identity of the anions, and the presence or absence of water of crystallization.[9] The structural formulas are [trans–FeCl
2(H
2O)
4][FeCl
4], [cis–FeCl
2(H
2O)
4][FeCl
4] · H2O, [cis–FeCl
2(H
2O)
4][FeCl
4] · H2O, and [trans–FeCl
2(H
2O)
4]Cl · 2H2O. The first three members of this series have the tetrahedral tetrachloroferrate ([FeCl
4]−
) anion.[12]

Solution

A brown, acidic solution of iron(III) chloride.

Like the solid hydrates, aqueous solutions of ferric chloride also consist of the octahedral [FeCl
2(H
2O)
4]+
 of unspecified stereochemistry.[9] Detailed speciation of aqueous solutions of ferric chloride is challenging because the individual components do not have distinctive spectroscopic signatures. Iron(III) complexes, with a high spin d5 configuration, is kinetically labile, which means that ligands rapidly dissociate and reassociate. A further complication is that these solutions are strongly acidic, as expected for aquo complexes of a tricationic metal. Iron aquo complexes are prone to olation, the formation of polymeric oxo derivatives. Dilute solutions of ferric chloride produce soluble nanoparticles with molecular weight of 104, which exhibit the property of "aging", i.e., the structure change or evolve over the course of days.[13] The polymeric species formed by the hydrolysis of ferric chlorides are key to the use of ferric chloride for water treatment.

In contrast to the complicated behavior of its aqueous solutions, solutions of iron(III) chloride in diethyl ether and tetrahydrofuran are well-behaved. Both ethers form 1:2 adducts of the general formula FeCl3(ether)2. In these complexes, the iron is pentacoordinate.[14]

Preparation

Several hundred thousand kilograms of anhydrous iron(III) chloride are produced annually. The principal method, called direct chlorination, uses scrap iron as a precursor:[10]

2 Fe + 3 Cl
2 → 2 FeCl
3

The reaction is conducted at several hundred degrees such that the product is gaseous. Using excess chlorine guarantees that the intermediate ferrous chloride is converted to the ferric state.[10] A similar but laboratory-scale process also has been described.[15][16]

Aqueous solutions of iron(III) chloride are also produced industrially from a number of iron precursors, including iron oxides:

Fe
2O
3 + 12 HCl + 9 H
2O → FeCl
3(H
2O)
6

In complementary route, iron metal can be oxidized by hydrochloric acid followed by chlorination:[10]

Fe + 2 HCl → FeCl
2 + H
2
FeCl
2 + 0.5 Cl
2 + 6 H
2O → FeCl
3(H
2O)
6

A number of variables apply to these processes, including the oxidation of iron by ferric chloride and the hydration of intermediates.[10] Hydrates of iron(III) chloride do not readily yield anhydrous ferric chloride. Attempted thermal dehydration yields hydrochloric acid and iron oxychloride. In the laboratory, hydrated iron(III) chloride can be converted to the anhydrous form by treatment with thionyl chloride[17] or trimethylsilyl chloride:[18]

FeCl
3 · 6H2O + 12 (CH
3)
3SiCl → FeCl
3 + 6 ((CH
3)
3Si)
2O + 12 HCl

Reactions

Being high spin d5 electronic configuration iron(III) chlorides are labile, meaning that its Cl- and H2O ligands exchange rapidly with free chloride and water.[9][19] In contrast to their kinetic lability, iron(III) chlorides are thermodynamically robust, as reflected by the vigorous methods applied to their synthesis, as described above.

Anhydrous FeCl3

Aside from lability, which applies to anhydrous and hydrated forms, the reactivity of anhydrous ferric chloride reveals two trends: It is a Lewis acid and an oxidizing agent.[20]

Reactions of anhydrous iron(III) chloride reflect its description as both oxophilic and a hard Lewis acid. Myriad manifestations of the oxophiliicty of iron(III) chloride are available. When heated with iron(III) oxide at 350 °C it reacts to give iron oxychloride:[21]

FeCl
3 + Fe
2O
3 → 3FeOCl

Alkali metal alkoxides react to give the iron(III) alkoxide complexes. These products have more complicated structures than anhydrous iron(III) chloride.[22][23] In the solid phase a variety of multinuclear complexes have been described for the nominal stoichiometric reaction between FeCl
3 and sodium ethoxide:

FeCl
3 + 3 CH
3CH
2ONa → "Fe(OCH
2CH
3)
3" + 3 NaCl

Iron(III) chloride forms a 1:2 adduct with Lewis bases such as triphenylphosphine oxide; e.g., FeCl
3(OP(C
6H
5)
3)
2. The related 1:2 complex FeCl
3(OEt
2)
2, where Et = C
2H
5), has been crystallized from ether solution.[14]

Iron(III) chloride also reacts with tetraethylammonium chloride to give the yellow salt of the tetrachloroferrate ion ((Et
4N)[FeCl
4]). Similarly combining FeCl3 with NaCl and KCl gives Na[FeCl
4] and K[FeCl
4], respectively.[24]

In addition to these simple stoichiometric reactions, the Lewis acidity of ferric chloride enables its use in a variety of acid-catalyzed reactions as described below in the section on organic chemistry.[10]

In terms of its being an oxidant, iron(III) chloride oxidizes iron powder to form iron(II) chloride via a comproportionation reaction:[10]

2 FeCl
3 + Fe → 3 FeCl
2

A traditional synthesis of anhydrous ferrous chloride is the reduction of FeCl3 with chlorobenzene:[25]

2 FeCl
3 + C
6H
5Cl → 2 FeCl
2 + C
6H
4Cl
2 + HCl

iron(III) chloride releases chlorine gas when heated above 160 °C, generating ferrous chloride:[16]

2FeCl
3 → 2FeCl
2 + Cl
2

To suppress this reaction, the preparation of iron(III) chloride requires an excess of chlorinating agent, as discussed above.[16][10]

Hydrated FeCl3

Unlike the anhydrous material, hydrated ferric chloride is not a particularly strong Lewis acid since water ligands have quenched the Lewis acidity by binding to Fe(III).

Like the anhydrous material, hydrated ferric chloride is oxophilic. For example, oxalate salts react rapidly with aqueous iron(III) chloride to give [Fe(C
2O
4)
3]3−, known as ferrioxalate. Other carboxylate sources, e.g., citrate and tartrate, bind as well to give carboxylate complexes. The affinity of iron(III) for oxygen ligands was the basis of qualitative tests for phenols. Although superseded by spectroscopic methods, the ferric chloride test is a traditional colorimetric test.[26] The affinity of iron(III) for phenols is exploited in the Trinder spot test.[27]

Aqueous iron(III) chloride serves as a one-electron oxidant illustrated by its reaction with copper(I) chloride to give copper(II) chloride and iron(II) chloride.

FeCl
3 + CuCl → FeCl
2 + CuCl
2

This fundamental reaction is relevant to the use of ferric chloride solutions in etching copper.

Organometallic chemistry

The interaction of anhydrous iron(III) chloride with organolithium and organomagnesium compounds has been examined often. These studies are enabled because of the solubility of FeCl3 in ethereal solvents, which avoids the possibility of hydrolysis of the nucleophilic alkylating agents. Such studies may be relevant to the mechanism of FeCl3-catalyzed cross-coupling reactions.[28] The isolation of organoiron(III) intermediates requires low-temperature reactions, lest the [FeR4]− intermediates degrade. Using methylmagnesium bromide as the alkylation agent, salts of Fe(CH3)4]− have been isolated.[29] Illustrating the sensitivity of these reactions, methyl lithium LiCH
3 reacts with iron(III) chloride to give lithium tetrachloroferrate(II) Li
2[FeCl
4]:[30]

2 FeCl
3 + LiCH
3 → FeCl
2 + Li[FeCl
4] + 0.5 CH
3CH
3
Li[FeCl
4] + LiCH
3 → Li
2[FeCl
4] + 0.5 CH
3CH
3

To a significant extent, iron(III) acetylacetonate and related beta-diketonate complexes are more widely used than FeCl3 as ether-soluble sources of ferric ion.[20] These diketonate complexes have the advantages that they do not form hydrates, unlike iron(III) chloride, and they are more soluble in relevant solvents.[28] Cyclopentadienyl magnesium bromide undergoes a complex reaction with iron(III) chloride, resulting in ferrocene:[31]

3 C
5H
5MgBr + FeCl
3 → Fe(C
5H
5)
2 + 1/n (C
5H
5)
n + 3 MgBrCl

This conversion, although not of practical value, was important in the history of organometallic chemistry where ferrocene is emblematic of the field.[32]

Uses

Water treatment

In the largest application iron(III) chloride is in sewage treatment and drinking water production. By forming highly dispersed networks of Fe-O-Fe containing materials, ferric chlorides serve as coagulant and flocculants.[33] In this application, an aqueous solution of FeCl
3 is treated with base to form a floc of iron(III) hydroxide (Fe(OH)
3), also formulated as FeO(OH) (ferrihydrite). This floc facilitates the separation of suspended materials, clarifying the water.[10]

Iron(III) chloride is also used to remove soluble phosphate from wastewater. Iron(III) phosphate is insoluble and thus precipitates as a solid.[34] One potential advantage to its use in water treatment, ferric ion oxidizes (deodorizes) hydrogen sulfide.[35]

Etching and metal cleaning

It is also used as a leaching agent in chloride hydrometallurgy,[36] for example in the production of Si from FeSi (Silgrain process by Elkem).[37]

In another commercial application, a solution of iron(III) chloride is useful for etching copper according to the following equation:

2 FeCl
3 + Cu → 2 FeCl
2 + CuCl
2

The soluble copper(II) chloride is rinsed away, leaving a copper pattern. This chemistry is used in the production of printed circuit boards (PCB).[19]

Iron(III) chloride is used in many other hobbies involving metallic objects.[38][39][40][41][42]

Organic chemistry

Structure of FeCl3(diethylether)2. Color code: Cl=green,Fe = blue, O = red.

In industry, iron(III) chloride is used as a catalyst for the reaction of ethylene with chlorine, forming ethylene dichloride (1,2-dichloroethane):[43]

H
2C=CH
2 + Cl
2 → ClCH
2CH
2Cl

Ethylene dichloride is a commodity chemical, which is mainly used for the industrial production of vinyl chloride, the monomer for making PVC.[44]

Illustrating it use as a Lewis acid, iron(III) chloride catalyses electrophilic aromatic substitution and chlorinations. In this role, its function is similar to that of aluminium chloride. In some cases, mixtures of the two are used.[45]

Organic synthesis research

Although iron(III) chlorides are seldom used in practical organic synthesis, they have received considerable attention as reagents because they are inexpensive, earth abundant. relatively nontoxic. Many experiments probe both its redox activity and its Lewis acidity.[20] For example, iron(III) chloride oxidizes naphthols to naphthoquinones:[20][46] 3-Alkylthiophenes are polymerized to polythiophenes upon treatment with ferric chloride.[47] Iron(III) chloride has been shown to promote C-C coupling reaction.[48]

Several reagents for have been developed based on supported iron(III) chloride. On silica gel, the anhydrous salt has been applied to certain dehydration and pinacol-type rearrangement reactions. A similar reagent but moistened induces hydrolysis or epimerization reactions.[49] On alumina ferric chloride has been shown to accelerate ene reactions.[50]

When pretreated with sodium hydride, iron(III) chloride gives a hydride reducing agent that convert alkenes and ketones into alkanes and alcohols, respectfully.[51]

FeCl3oxidation.svg

Histology

Iron(III) chloride is a component of useful stains, such as Carnoy's solution, a histological fixative with many applications. Also, it is used to prepare Verhoeff's stain.[52]

Natural occurrence

Like many metal halides, FeCl
3 naturally occurs as a trace mineral. The rare mineral molysite is usually associated with volcanoes and fumaroles.[53][54]

FeCl
3-based aerosol are produced by a reaction between iron-rich dust and hydrochloric acid from sea salt. This iron salt aerosol causes about 1-5% of naturally-occurring oxidization of methane and is thought to have a range of cooling effects.[55] Therefore it has been proposed as a catalyst for Atmospheric Methane Removal.

The clouds of Venus are hypothesized to contain approximately 1% FeCl
3 dissolved in sulfuric acid.[56][57]

Safety

Iron(III) chlorides are widely used in production of drinking water,[10] so they pose few problems as poisons, at low concentrations. Nonetheless, anhydrous iron(III) chloride, as well as concentrated FeCl
3 aqueous solution, is highly corrosive, and must be handled using proper protective equipment.[20]

Notes

  1. ↑ An alternative GHS classification from the Japanese GHS Inter-ministerial Committee (2006)[6] notes the possibility of respiratory tract irritation from FeCl
    3     and differs slightly in other respects from the classification used here.

References

  1. ↑ 1.0 1.1 1.2 1.3 1.4 1.5 Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.69. ISBN 1439855110. 
  2. ↑ Haynes, William M., ed (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.133. ISBN 1439855110. 
  3. ↑ NIOSH Pocket Guide to Chemical Hazards. "#0346". National Institute for Occupational Safety and Health (NIOSH). https://www.cdc.gov/niosh/npg/npgd0346.html. 
  4. ↑ HSNO Chemical Classification Information Database, New Zealand Environmental Risk Management Authority, http://www.epa.govt.nz/search-databases/Pages/ccid-details.aspx?SubstanceID=10764, retrieved 19 Sep 2010 
  5. ↑ Various suppliers, collated by the Baylor College of Dentistry, Texas A&M University. (accessed 2010-09-19)
  6. ↑ Template:GHS class JP
  7. ↑ 7.0 7.1 "Structure refinement of an FeCl3 crystal using a thin plate sample". J. Appl. Crystallogr. 22 (2): 173–180. 1989. doi:10.1107/S0021889888013913. 
  8. ↑ Housecroft, C. E.; Sharpe, A. G. (2012). Inorganic Chemistry (4th ed.). Prentice Hall. pp. 747. ISBN 978-0-273-74275-3. 
  9. ↑ 9.0 9.1 9.2 9.3 Simon A. Cotton (2018). "Iron(III) Chloride and Its Coordination Chemistry". Journal of Coordination Chemistry 71 (21): 3415–3443. doi:10.1080/00958972.2018.1519188. 
  10. ↑ 10.0 10.1 10.2 10.3 10.4 10.5 10.6 10.7 10.8 10.9 Wildermuth, Egon; Stark, Hans; Friedrich, Gabriele; Ebenhöch, Franz Ludwig; Kühborth, Brigitte; Silver, Jack; Rituper, Rafael (2000). "Iron Compounds". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a14_591. ISBN 3527306730. 
  11. ↑ Inorganic Chemistry. San Diego: Academic Press. 2001. ISBN 978-0-12-352651-9. 
  12. ↑ Lind, M. D. (1967). "Crystal Structure of Ferric Chloride Hexahydrate". The Journal of Chemical Physics 47 (3): 990–993. doi:10.1063/1.1712067. Bibcode: 1967JChPh..47..990L. 
  13. ↑ Flynn, Charles M. (1984). "Hydrolysis of Inorganic Iron(III) Salts". Chemical Reviews 84: 31–41. doi:10.1021/cr00059a003. 
  14. ↑ 14.0 14.1 Spandl, Johann; Kusserow, M.; Brüdgam, I. (2003). "Alkoxo-Verbindungen des dreiwertigen Eisen: Synthese und Charakterisierung von [Fe2(Ot Bu)6], [Fe2Cl2(Ot Bu)4], [Fe2Cl4(Ot Bu)2] und [N(n Bu)4]2[Fe6OCl6(OMe)12]". Zeitschrift für anorganische und allgemeine Chemie 629 (6): 968–974. doi:10.1002/zaac.200300008. 
  15. ↑ "Anhydrous Iron(III) Chloride (Ferric Chloride)". Inorganic Syntheses. 3. 1950. pp. 191–194. doi:10.1002/9780470132340.ch51. ISBN 9780470131626. 
  16. ↑ 16.0 16.1 16.2 H. Lux (1963). "Iron (III) Chloride". in G. Brauer. Handbook of Preparative Inorganic Chemistry, 2nd Ed.. 2. NY, NY: Academic Press. pp. 1492. 
  17. ↑ "Anhydrous Metal Chlorides". Inorganic Syntheses. 28. 1990. pp. 321–323. doi:10.1002/9780470132593.ch80. ISBN 9780470132593. 
  18. ↑ "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorganic Syntheses. 29. 1992. pp. 108–111. doi:10.1002/9780470132609.ch26. ISBN 9780470132609. 
  19. ↑ 19.0 19.1 Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. 1997. pp. 1084. ISBN 9780750633659. 
  20. ↑ 20.0 20.1 20.2 20.3 20.4 White, Andrew D.; Gallou, Fabrice (2006). "Iron(III) Chloride". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.ri054.pub2. ISBN 0471936235. 
  21. ↑ "Layered Intercalation Compounds". Inorganic Syntheses. 22. John Wiley & Sons, Inc.. 1984. pp. 86–89. doi:10.1002/9780470132531.ch17. ISBN 9780470132531. 
  22. ↑ "12.22.1 Synthesis". The Chemistry of Metal Alkoxides. Springer Science. 2002. pp. 481. ISBN 0306476576. https://books.google.com/books?id=rPzaMRjK8pQC. 
  23. ↑ "3.2.10. Alkoxides of later 3d metals". Alkoxo and aryloxo derivatives of metals. San Diego: Academic Press. 2001. pp. 69. ISBN 9780121241407. OCLC 162129468. https://books.google.com/books?id=ixNEpx8aDHUC&pg=PA69. 
  24. ↑ Cook, Charles M. Jr.; Dunn, Wendell E. Jr. (1961). "The Reaction of Ferric Chloride with Sodium and Potassium Chlorides". J. Phys. Chem. 65 (9): 1505–1511. doi:10.1021/j100905a008. 
  25. ↑ P. Kovacic and N. O. Brace (1960). "Iron(II) Chloride". Inorganic Syntheses. 6. pp. 172–173. doi:10.1002/9780470132371.ch54. ISBN 9780470132371. 
  26. ↑ Vogel's Textbook of Practical Organic Chemistry (5th ed.). New York: Longman/John Wiley & Sons. 1989. ISBN 9780582462366. https://archive.org/details/Vogels_Textbook_of_Practical_Organic_Chemistry_5ed_1989_Longman_WW. 
  27. ↑ James A. King; Alan B. Storrow; Jeff A. Finkelstein (1995). "Urine Trinder Spot Test: A Rapid Salicylate Screen for the Emergency Department" (in en). Annals of Emergency Medicine 26 (3): 330-333. doi:10.1016/S0196-0644(95)70082-X. 
  28. ↑ 28.0 28.1 Mako, T. L.; Byers, J. A. (2016). "Recent Advances in Iron-Catalysed Cross Coupling Reactions and Their Mechanistic Underpinning". Inorganic Chemistry Frontiers 3 (6): 766–790. doi:10.1039/C5QI00295H. 
  29. ↑ Sears, Jeffrey D.; Muñoz, Salvador B.; Cuenca, Maria Camila Aguilera; Brennessel, William W.; Neidig, Michael L. (2019). "Synthesis and Characterization of a Sterically Encumbered Homoleptic Tetraalkyliron(III) Ferrate Complex". Polyhedron 158: 91–96. doi:10.1016/j.poly.2018.10.041. PMID 31031511.  and references therein.
  30. ↑ "Über die Bildung von Lithiumtetrachloroferrat(II) Li2FeCl4 bei der Umsetzung von Eisen(III)-chlorid mit Lithiummethyl (1:1) in ätherischer Lösung" (in de). Z. Anorg. Allg. Chem. 391 (3): 193–202. 1972. doi:10.1002/zaac.19723910302. 
  31. ↑ "A New Type of Organo-Iron Compound". Nature 168 (4285): 1040. 1951. doi:10.1038/1681039b0. Bibcode: 1951Natur.168.1039K. 
  32. ↑ "Ferrocene—how it all began". Journal of Organometallic Chemistry 637–639: 3–6. 2001. doi:10.1016/S0022-328X(01)01126-3. 
  33. ↑ Water Treatment Chemicals. Akzo Nobel Base Chemicals. 2007. http://www.akzonobel.com/ic/system/images/AkzoNobel_WTCBrochureENG_tcm18-9982.pdf. Retrieved 26 Oct 2007. 
  34. ↑ "Phosphorus Treatment and Removal Technologies". Minnesota Pollution Control Agency. June 2006. https://www.pca.state.mn.us/sites/default/files/wq-wwtp9-02.pdf. 
  35. ↑ Prathna, T. C.; Srivastava, Ankit (2021). "Ferric chloride for odour control: studies from wastewater treatment plants in India". Water Practice and Technology 16 (1): 35–41. doi:10.2166/wpt.2020.111. 
  36. ↑ "A study on the acidified ferric chloride leaching of a complex (Cu–Ni–Co–Fe) matte". Separation and Purification Technology 51 (3): 332–337. 2006. doi:10.1016/j.seppur.2006.02.013. 
  37. ↑ "Validation of a compartmental population balance model of an industrial leaching process: The Silgrain process". Chem. Eng. Sci. 61 (1): 229–245. 2006. doi:10.1016/j.ces.2005.01.047. Bibcode: 2006ChEnS..61..229D. 
  38. ↑ "Safer Printmaking—Intaglio". https://research-groups.usask.ca/saferprintmaking/intaglio.php. 
  39. ↑ Harris, Paul; Hartman, Ron; Hartman, James (November 1, 2002). "Etching Iron Meteorites". Meteorite Times. http://www.meteorite-times.com/articles/etching-iron-meteorites/. 
  40. ↑ Lockwood, Mike. "A message about mirror coating and recoating". http://www.loptics.com/articles/coatingrisk/coatingrisk.html. 
  41. ↑ "Buffalo Nickel No Date Value Guides (Year Reveal Tutorial)". 4 January 2023. https://robpaulsenlive.com/indian-head-nickel-no-date-value/. 
  42. ↑ Scott, David; Schwab, Roland (2019). "3.1.4. Etching". Metallography in Archaeology and Art. Cultural Heritage Science. Springer. doi:10.1007/978-3-030-11265-3. ISBN 978-3-030-11265-3. 
  43. ↑ Dreher, Eberhard-Ludwig; Beutel, Klaus K.; Myers, John D.; Lübbe, Thomas; Krieger, Shannon; Pottenger, Lynn H. (2014). "Chloroethanes and Chloroethylenes". Ullmann's Encyclopedia of Industrial Chemistry. pp. 1–81. doi:10.1002/14356007.o06_o01.pub2. ISBN 9783527306732. 
  44. ↑ "Toxic Substances – 1,2-Dichloroethane". https://wwwn.cdc.gov/TSP/index.aspx?toxid=110. 
  45. ↑ Riddell, W. A.; Noller, C. R. (1932). "Mixed Catalysis in the Friedel and Crafts Reaction. The Yields in Typical Reactions using Ferric Chloride–Aluminum Chloride Mixtures as Catalysts". J. Am. Chem. Soc. 54 (1): 290–294. doi:10.1021/ja01340a043. 
  46. ↑ Louis F. Fieser (1937). "1,2-Naphthoquinone". Organic Syntheses 17: 68. doi:10.15227/orgsyn.017.0068. 
  47. ↑ So, Regina C.; Carreon-Asok, Analyn C. (2019). "Molecular Design, Synthetic Strategies, and Applications of Cationic Polythiophenes". Chemical Reviews 119 (21): 11442–11509. doi:10.1021/acs.chemrev.8b00773. PMID 31580649. 
  48. ↑ Albright, Haley; Davis, Ashlee J.; Gomez-Lopez, Jessica L.; Vonesh, Hannah L.; Quach, Phong K.; Lambert, Tristan H.; Schindler, Corinna S. (2021). "Carbonyl–Olefin Metathesis". Chemical Reviews 121 (15): 9359–9406. doi:10.1021/acs.chemrev.0c01096. PMID 34133136. 
  49. ↑ White, Andrew D. (2001). "Iron(III) Chloride-Silica Gel". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.ri059. ISBN 0471936235. 
  50. ↑ White, Andrew D. (2001). "Iron(III) Chloride-Alumina". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.ri057. ISBN 0471936235. 
  51. ↑ White, Andrew D. (2001). "Iron(III) Chloride-Sodium Hydride". Encyclopedia of Reagents for Organic Synthesis. doi:10.1002/047084289X.ri060. ISBN 0471936235. 
  52. ↑ Mallory; Sheehan; Hrapchak (1990). "Histotechnology – A Self-Instructional Text". in Carson, Freida; Cappellano, Christa Hladik. Histotechnology – A Self-Instructional Text. Chicago: ASCP Press. http://tvmouse.ucdavis.edu/bcancercd/52/prcl_elastic.html. Retrieved 2 January 2013. 
  53. ↑ "Molysite". https://www.mindat.org/min-2749.html. 
  54. ↑ "List of Minerals". March 21, 2011. https://www.ima-mineralogy.org/Minlist.htm. 
  55. ↑ Oeste, Franz Dietrich; de Richter, Renaud; Ming, Tingzhen; Caillol, Sylvain (January 13, 2017). "Climate engineering by mimicking natural dust climate control: the iron salt aerosol method". Earth System Dynamics 8 (1): 1–54. doi:10.5194/esd-8-1-2017. Bibcode: 2017ESD.....8....1O. https://esd.copernicus.org/articles/8/1/2017/. 
  56. ↑ Krasnopolsky, V. A.; Parshev, V. A. (1981). "Chemical composition of the atmosphere of Venus". Nature 292 (5824): 610–613. doi:10.1038/292610a0. Bibcode: 1981Natur.292..610K. 
  57. ↑ Krasnopolsky, Vladimir A. (2006). "Chemical composition of Venus atmosphere and clouds: Some unsolved problems". Planetary and Space Science 54 (13–14): 1352–1359. doi:10.1016/j.pss.2006.04.019. Bibcode: 2006P&SS...54.1352K. 

Further reading

  1. CRC Handbook of Chemistry and Physics (71st ed.). Ann Arbor, Michigan, US: CRC Press. 1990. ISBN 9780849304712. 
  2. The Merck Index of Chemicals and Drugs (7th ed.). Rahway, New Jersey, US: Merck & Co. 1960. 
  3. Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.. A Macmillan chemistry text. London: Macmillan Press. 1974. ISBN 9780333170885. 
  4. Structural Inorganic Chemistry. Oxford science publications (5th ed.). Oxford, UK: Oxford University Press. 1984. ISBN 9780198553700. 
  5. Acidic and Basic Reagents. Handbook of Reagents for Organic Synthesis. New York: John Wiley & Sons, Inc.. 1999. ISBN 9780471979258. 



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