Thiocyanogen

Thiocyanogen
Names
Preferred IUPAC name
Cyanic dithioperoxyanhydride
Other names
Dicyanodisulfane
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
UNII
  • InChI=1S/C2N2S2/c3-1-5-6-2-4 checkY
    Key: DTMHTVJOHYTUHE-UHFFFAOYSA-N checkY
  • InChI=1/C2N2S2/c3-1-5-6-2-4
    Key: DTMHTVJOHYTUHE-UHFFFAOYAE
  • N#CSSC#N
Properties
C2N2S2
Molar mass 116.16 g mol−1
Appearance Colorless crystal or liquid[1]: 241, 255–256 
Melting point −2.5 °C (27.5 °F; 270.6 K)[1]: 241 
Boiling point ≈20 °C (decomposes)[2]
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Thiocyanogen, (SCN)2, is a pseudohalogen derived from the pseudohalide thiocyanate, [SCN]−, with behavior intermediate between dibromine and diiodine.[2] This hexatomic compound exhibits C2 point group symmetry and has the connectivity NCS-SCN.[3]

In the lungs, lactoperoxidase may oxidize thiocyanate to thiocyanogen[4] or hypothiocyanite.[5]

History

Berzelius first proposed that thiocyanogen ought exist as part of his radical theory, but the compound's isolation proved problematic. Liebig pursued a wide variety of synthetic routes for the better part of a century, but, even with Wöhler's assistance, only succeeded in producing a complex mixture with the proportions of thiocyanic acid. In 1861, Linnemann generated appreciable quantities of thiocyanogen from a silver thiocyanate suspension in diethyl ether and excess iodine, but misidentified the minor product as sulfur iodide cyanide (ISCN).[6] Indeed, that reaction suffers from competing equilibria attributed to the weak oxidizing power of iodine; the major product is sulfur dicyanide.[7] The following year, Schneider produced thiocyangen from silver thiocyanate and disulfur dichloride, but the product disproportionated to sulfur and trisulfur dicyanides.[6]

The subject then lay fallow until the 1910s, when Niels Bjerrum began investigating gold thiocyanate complexes. Some eliminated reductively and reversibly, whereas others appeared to irreversibly generate cyanide and sulfate salt solutions. Understanding the process required reanalyzing the decomposition of thiocyanogen using the then-new techniques of physical chemistry. Bjerrum's work revealed that water catalyzed thiocyanogen's decomposition via hypothiocyanous acid. Moreover, the oxidation potential of thiocyanogen appeared to be 0.769 V, slightly greater than iodine but less than bromine.[6] In 1919, Söderbäck successfully isolated stable thiocyanogen from oxidation of oxidation of plumbous thiocyanate with bromine.[6][7]

Preparation

Modern syntheses typically differ little from Söderbäck's process. Thiocyanogen synthesis begins when aqueous solutions of lead(II) nitrate and sodium thiocyanate, combined, precipitate plumbous thiocyanate. Treating an anhydrous Pb(SCN)2 suspension in glacial acetic acid with bromine then affords a 0.1M solution of thiocyanogen that is stable for days.[8] Alternatively, a solution of bromine in methylene chloride is added to a suspension of Pb(SCN)2 in methylene chloride at 0 °C.[9]

Pb(SCN)2 + Br2 → (SCN)2 + PbBr2

In either case, the oxidation is exothermic.[1]: 255 

An alternative technique is the thermal decomposition of cupric thiocyanate at 35–80 °C:[1]: 253 

2Cu(SCN)2 → CuSCN + (SCN)2

Reactions

In general, thiocyanogen is stored in solution, as the pure compound explodes above 20 °C[2] to a red-orange polymer.[1]: 241  However, the sulfur atoms disproportionate in water:[1]: 241–242 [10]

3(SCN)2 + 4H2O → H2SO4 + HCN + 5HSCN

Thiocyanogen is a weak electrophile, attacking only highly activated (phenolic or anilinic) or polycyclic arenes.[1]: 243–245  It attacks carbonyls at the α position.[1] Heteratoms are attacked more easily, and the compound thiocyanates sulfur, nitrogen, and various poor metals.[1]: 241  Thiocyanogen solutions in nonpolar solvents react almost completely with chlorine to give chlorine thiocyanate; but the corresponding bromine thiocyanate is unstable above −50 °C, forming polymeric thiocyanogen and bromine.[11]

The compound adds trans to alkenes to give 1,2-bis(thiocyanato) compounds; the intermediate thiiranium ion can be trapped with many nucleophiles.[2] Radical polymerization is the most likely side-reaction, and yields improve when cold and dark.[2][1]: 247  However, the addition reaction is slow, and light may be necessary to accelerate the process.[1]: 247  Titanacyclopentadienes give (Z,Z)-1,4-bis(thiocyanato)-1,3-butadienes, which in turn can be converted to 1,2-dithiins.[9] Thiocyanogen only adds once to alkynes; the resulting dithioacyloin dicyanate is not particularly olefinic.[1]: 247 

Selenocyanogen, (SeCN)2, prepared from reaction of silver selenocyanate with iodine in tetrahydrofuran at 0 °C,[12] reacts in a similar manner to thiocyanogen.[9]

Applications

Thiocyanogen has been used to estimate the degree of unsaturation in fatty acids, similar to the iodine value.[6][1]: 247 

References

  1. ^ a b c d e f g h i j k l m Wood, John L. (August 1947) [1946]. "Substitution and addition reactions of thiocyanogen". In Adams, Roger (ed.). Organic Reactions (PDF). Vol. 3 (3rd reprint ed.). New York / London: Wiley / Chapman Hall. pp. 241–266.
  2. ^ a b c d e Schwan, Adrian L. (2001-04-15), "Thiocyanogen", Encyclopedia of Reagents for Organic Synthesis, Chichester, UK: John Wiley & Sons, Ltd, doi:10.1002/047084289x.rt095, ISBN 978-0-471-93623-7, retrieved 2024-03-30
  3. ^ Jensen, James (2005). "Vibrational frequencies and structural determination of thiocyanogen". Journal of Molecular Structure: THEOCHEM. 714 (2–3): 137–141. doi:10.1016/j.theochem.2004.09.046.
  4. ^ Aune, Thomas M.; Thomas, Edwin L. (1977) [2 May 1977]. "Accumulation of hypothiocyanite ion during peroxidase-catalyzed oxidation of thiocyanate ion". European Journal of Biochemistry. 80: 209–214. doi:10.1111/j.1432-1033.1977.tb11873.x.
  5. ^ Lemma, Kelemu; Ashby, Michael T. (2009-09-21). "Reactive Sulfur Species: Kinetics and Mechanism of the Reaction of Hypothiocyanous Acid with Cyanide To Give Dicyanosulfide in Aqueous Solution". Chemical Research in Toxicology. 22 (9): 1622–1628. doi:10.1021/tx900212r. ISSN 0893-228X.
  6. ^ a b c d e Kaufmann, H. P. (1925). "Das freie Rhodan und seine Anwendung in der Maßanalyse. Eine neue Kennzahl der Fette" [Unbound rhodanium and its application to elemental analysis: A new measurement technique for fats]. Archiv der Pharmazie und Berichte der Deutschen Pharmazeutischen Gesellschaft (in German). 263: 675–721 – via HathiTrust.
  7. ^ a b Söderbäck, Erik (1919). "Studien über das freie Rhodan". Justus Liebig's Annalen der Chemie. 419 (3): 217–322. doi:10.1002/jlac.19194190302. hdl:2027/uc1.$b133351.
  8. ^ Gardner, William Howlett; Weinberger, Harold (1939). "Thiocyanogen Solution". Inorganic Syntheses. Inorganic Syntheses. Vol. 1. pp. 84–86. doi:10.1002/9780470132326.ch29. ISBN 978-0-470-13232-6.
  9. ^ a b c Block, E; Birringer, M; DeOrazio, R; Fabian, J; Glass, RS; Guo, C; He, C; Lorance, E; Qian, Q; Schroeder, TB; Shan, Z; Thiruvazhi, M; Wilson, GC; Zhang, Z (2000). "Synthesis, Properties, Oxidation, and Electrochemistry of 1,2-Dichalcogenins". J. Am. Chem. Soc. 122 (21): 5052–5064. doi:10.1021/ja994134s.
  10. ^ Stedman, G.; Whincup, P. A. E. (1969). "Oxidation of metal thiocyanates by nitric and nitrous acids. Part I. Products". Journal of the Chemical Society A: Inorganic, Physical, Theoretical: 1145. doi:10.1039/j19690001145. ISSN 0022-4944.
  11. ^ Magee, Philip S. (1971). "The Sulfur–Bromine Bond". In Senning, Alexander (ed.). Sulfur in Organic and Inorganic Chemistry. Vol. 1. New York: Marcel Dekker. pp. 269–270. ISBN 0-8247-1615-9. LCCN 70-154612.
  12. ^ Meinke, PT; Krafft, GA; Guram, A (1988). "Synthesis of selenocyanates via cyanoselenation of organocopper reagents". J. Org. Chem. 53 (15): 3632–3634. doi:10.1021/jo00250a047.