Ammonium thiocyanate

Ammonium thiocyanate is an inorganic compound with the formula NH4SCN. It is the salt of the ammonium cation and the thiocyanate anion.

Ammonium thiocyanate
Space-filling model of the ammonium cation
Space-filling model of the thiocyanate anion
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.015.614
RTECS number
  • XN6465000
UNII
Properties
NH4SCN
Molar mass 76.122 g/mol
Appearance Colorless hygroscopic crystalline solid
Density 1.305 g/cm3
Melting point 149.5 °C (301.1 °F; 422.6 K)
Boiling point 170 °C (338 °F; 443 K) (decomposes)
128 g/100 mL (0 °C)
Solubility soluble in liquid ammonia, alcohol, acetone
-48.1·10−6 cm3/mol
Hazards
Safety data sheet External MSDS
NFPA 704 (fire diamond)
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
1
2
1
Related compounds
Other anions
 Ammonium cyanate
Ammonium cyanide
Other cations
 Sodium thiocyanate
Potassium thiocyanate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

Uses

Ammonium thiocyanate is used in the manufacture of herbicides, thiourea, and transparent artificial resins; in matches; as a stabilizing agent in photography; in various rustproofing compositions; as an adjuvant in textile dyeing and printing; as a tracer in oil fields; in the separation of hafnium from zirconium, and in titrimetric analyses.

In May 1945, USAAF General Victor E. Betrandias advanced a proposal to his superior General Arnold to use of ammonium thiocyanate to reduce rice crops in Japan as part of the bombing raids on their country.[1]

Ammonium thiocyanate can also be used to determine the iron content in soft drinks by colorimetry.

Ammonium thiocyanate may also be used to separate quinidine, from liquors, after the isolation of quinine from the neutral, aqueous, sulphate solution. The salt is added to the hot solution and the gummy solid that forms is strained off from the liquid. The solid is then refluxed with methanol, which dissolves most of the impurities, leaving the quinidine thiocyanate as a crystalline solid of 90 - 95% purity. Following separation, (usually by centrifuge) the solid may then be further purified to pharmaceutical quality. (Quinidine is used for the treatment of heart arrhythmia and therefore has considerable value.)


Preparation

Ammonium thiocyanate is made by the reaction of carbon disulfide with aqueous ammonia. Ammonium dithiocarbamate is formed as an intermediate in this reaction, which upon heating, decomposes to ammonium thiocyanate and hydrogen sulfide:

CS2 + 2 NH3(aq) → NH2C(=S)SNH4 → NH4SCN + H2S

Reactions

Ammonium thiocyanate is stable in air; however, upon heating it isomerizes to thiourea:

The equilibrium mixtures at 150 °C and 180 °C contain 30.3% and 25.3% (by weight) thiourea, respectively. When heated at 200 °C, the dry powder decomposes to ammonia, hydrogen sulfide, and carbon disulfide, leaving a residue of guanidinium thiocyanate.

NH4SCN is weakly acidic due to the ammonium ion; it reacts with alkali hydroxides, such as sodium hydroxide or potassium hydroxide to form sodium thiocyanate or potassium thiocyanate, along with water and ammonia. The thiocyanate anion, specifically, reacts with ferric salts to form a deep-red ferric thiocyanate complex:

6 SCN− + Fe3+ → [Fe(SCN)6]3−

Ammonium thiocyanate reacts with several metal ions including copper, silver, zinc, lead, and mercury, forming their thiocyanate precipitates, which may be extracted into organic solvents.

References
  1. John David Chappell (1997). Before the Bomb: How America Approached the End of the Pacific War. University Press of Kentucky. pp. 91–92. ISBN 978-0-8131-7052-7.
  1. A. F. Wells, Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984. ISBN 978-0198553700
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