Iron (III) chloride
|__ Fe 3+ __ Cl -|
|Surname||Iron (III) chloride|
|Ratio formula||FeCl 3|
|External identifiers / databases|
2.89 g cm −3
304 ° C
319 ° C , from 120 ° C sublimation
|As far as possible and customary, SI units are used. Unless otherwise noted, the data given apply to standard conditions .|
Iron (III) chloride is a chemical compound of iron (III) and chloride ions . The Roman number III indicates the oxidation number of the iron ion (in this case +3). Iron (III) chloride belongs to the group of iron halides .
The term iron chloride also includes the compound iron (II) chloride (FeCl 2 ).
Iron (III) chloride occurs naturally in the form of the minerals molysite ( anhydrate ) and hydromolysite (hexahydrate).
Extraction and presentation
Anhydrous iron (III) chloride can be obtained in the laboratory by conducting chlorine at temperatures of 250 to 400 ° C over iron wire, wool or the like. Subsequently, the product is purified in a stream of chlorine at 220 to a maximum of 300 ° C sublimes . It is important to ensure that devices and chemicals are as free of water as possible.
Alternatively, the anhydrous connection can be made by reacting the hexahydrate with thionyl chloride :
Iron chloride containing water of crystallization can also be obtained by dissolving iron powder in hydrochloric acid
and the subsequent introduction of chlorine, the iron (II) chloride initially formed being converted into iron (III) chloride:
This can then be obtained by evaporating the solution. For technical production, chlorine is passed over scrap iron at around 650 ° C.
It is also possible to oxidize the iron (II) ion to the iron (III) ion with hydrogen peroxide :
This reaction can be used to detect iron in water, since a blood-red solution of iron (III) rhodanide is formed when a potassium or ammonium rhodanide solution is subsequently added in the presence of iron .
Anhydrous iron (III) chloride is stored under protective gas (e.g. nitrogen ) in the absence of air to protect it from water .
Anhydrous iron (III) chloride is a black substance with a slightly pungent odor of hydrochloric acid . As an anhydrous compound, it is extremely hygroscopic , i.e. it removes water from the air . With increasing water content the hygroscopic nature decreases and the color changes from red-brownish to yellowish, iron (III) chloride hexahydrate (FeCl 3 · 6 H 2 O) is formed. This reacts strongly acidic through hydrolysis . When heated, the hydrate decomposes with elimination of water and hydrogen chloride; It is therefore not possible in this way to obtain anhydrous iron (III) chloride from it again. However, anhydrous iron (III) chloride is accessible via the hexahydrate if it is mixed with an excess of thionyl chloride (10 eq.) Under a protective gas atmosphere . The resulting gases are boiled off and neutralized by introducing them into a cooled sodium hydroxide solution (approx. 0.5 - 1.0 M).
Iron (III) chloride is a predominantly covalent compound with a layered structure. Above the sublimation point it is mainly in the form of gaseous Fe 2 Cl 6 , which increasingly dissociates to FeCl 3 as the temperature rises . Anhydrous iron (III) chloride behaves chemically in a similar way to anhydrous aluminum chloride . Just like this, it is a moderately strong Lewis acid .
Iron (III) chloride and iron (II) chloride form a eutectic that melts at 297.5 ° C and contains 13.4 mol% of iron (II) chloride.
Iron (III) chloride can oxidize and dissolve copper ; Therefore, you can use aqueous iron (III) chloride solutions for the gentle etching of circuit boards and printing plates :
Iron (III) chloride is used to bind hydrogen sulfide , to precipitate phosphate and also as a precipitant in simultaneous precipitation and generally as a flocculant in biological wastewater treatment . In the chemical industry it is used as a selective catalyst in many Friedel-Crafts reactions . Many phenols produce green or blue colored complexes with iron (III) chloride and can be detected in this way ( iron chloride test ). The Berlin blue pigment can be produced by adding potassium hexacyanoferrate (II) (see below).
In aqueous solution it is used in textile printing as an oxidizing agent and stain, in medicine for intravenous substitution in severe deficiency conditions and as a hemostatic agent ( hemostyptic or astringent, no longer available in Germany), for corrosion tests (according to ASTM G48A) and for etching of metals (e.g. in copper gravure printing ) and of circuit boards in printed circuits and in the production of dyes (e.g. aniline black ).
Iron (III) chloride is harmful if swallowed and irritates the skin. There is a risk of serious eye damage. In combination with alkali metals , allyl chloride and ethylene oxide there is a risk of explosion.
About Fe 3+ ions
If you add potassium hexacyanoferrate (II) to iron (III) chloride solution , a deep blue precipitate of the pigment Berlin blue is produced :
Another very sensitive detection is carried out using thiocyanate ions (SCN - ):
The complex pentaquathiocyanato iron (III) ions formed appear intensely red.
Another proof would be the red-brown precipitate of iron (III) oxide hydrate (" iron (III) hydroxide "), which is formed when reacting with OH - ions.
The reaction with 3-chlorosalicylic acid gives an intense purple color.
- Gerhart Jander , Ewald Blasius et al .: Introduction to the inorganic-chemical internship . 14. re-edit Ed. Hirzel, Stuttgart, 1995. ISBN 3-7776-0672-3 .
- Michael Binnewies , Manfred Jäckel et al .: General and Inorganic Chemistry . Spectrum Academic Publishing House, 2003. ISBN 3-8274-0208-5 .
- ↑ a b c Entry on iron chlorides. In: Römpp Online . Georg Thieme Verlag, accessed on May 30, 2014.
- ↑ a b c Arnold Willmes, Taschenbuch Chemical Substances , Harri Deutsch, Frankfurt (M.), 2007.
- ↑ a b Georg Brauer: Handbook of preparative inorganic chemistry . 3., reworked. Edition. tape III . Enke, Stuttgart 1981, ISBN 3-432-87823-0 , pp. 1641 .
- ^ Alfred R. Pray: Anhydrous metal chlorides . In: Therald Moeller (Ed.): Inorganic Syntheses . tape 5 . McGraw-Hill, Inc., 1957, pp. 153-156 (English).
- ↑ a b c Wilhelm Uphoff: Chemical-technical internship: With 36 pictures . Vieweg, 1966, p. 21 ( limited preview in Google Book search).
- ↑ N. Walker, GH Wiltshire: The decomposition of 1-chloro- and 1-bromonaphthalene by soil bacteria. In: Journal of General Microbiology. 12 (3), 1955, pp. 478-483, doi: 10.1099 / 00221287-12-3-478 , PMID 14392303 .