Lithium cyanide

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Lithium cyanide[1][2][3]
Lithium-cyanide-unit-cell-3D-SF.png
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.017.554
UN number 1935
Properties
LiCN
Molar mass 32.959 g/mol
Appearance White Powder
Density 1.073 g/cm3 (18 °C)
Melting point 160 °C (320 °F; 433 K) Dark colored
Boiling point N/A
Soluble
N/A
Structure
-
Fourfold
Hazards
Safety data sheet 742899
T+, Very Toxic N, Dangerous for the environment
R-phrases (outdated) 26/27/28-32-50/53
S-phrases (outdated) 7-28-29-45-60-61
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 4: Very short exposure could cause death or major residual injury. E.g., VX gas Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point 57 °C (135 °F; 330 K)
N/A
Related compounds
Related compounds
Sodium cyanide, Potassium cyanide, Hydrogen cyanide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Lithium cyanide is an inorganic compound with the chemical formula LiCN. It is a white powder at room temperature. Lithium cyanide is commonly used as a reagent in inorganic/organometallic reactions. Lithium cyanide can be found in the environment from the reaction of lithium and acetonitrile, both found in lithium sulfur oxide batteries. When the compound is exposed to the environment it can produce toxic fumes with weak acids found in nature.

Properties[edit]

Stability and reactivity[edit]

Lithium cyanide as a solid is stable at room temperature. LiCN, when melted at 160 °C, is highly hygroscopic. The compound decomposes to cyanamide and carbon when heated to a temperature close to but below 600°C. When acids, chlorates, and other strong oxidizing agents react with LiCN, HCN is formed. HCN vapors are very toxic and reactive. If LiCN is heated in fire carbon dioxide CO2, nitrogen oxides NOx, and lithium oxide will form.[4]

Reactions[edit]

Synthesis and production[edit]

Li + R-CN → LiCN
Li + HCN(Benzene)[clarification needed] → LiCN

Lithium cyanide can be synthesized in high yields with liquid hydrogen cyanide and n-butyllithium. Other methods exist, in which the lithium cation is added to the cyanide anion.[4]

Cyanation[edit]

RX + LiCN —THF[clarification needed]→ RCN

Lithium cyanide is commonly used as a reagent in syntheses of cyanide compounds, for example halide cyanides. The reagent offers the advantage of allowing non-aqueous methods of cyanation.[5]

Environmental exposure[edit]

Lithium cyanide is an inorganic compound not commonly found in nature without human involvement. Likely, the most obtainable source[clarification needed] of lithium cyanide are lithium sulfur dioxide batteries, where the reaction between the two chemicals found inside the battery, elemental lithium and acetonitrile, has been shown to lead to the formation of lithium cyanide.[6] When lithium cyanide is introduced into the environment, it can react with acids or strong oxidizing agents to produce toxic HCN vapors in the environment or can produce carbon dioxide, nitrogen oxides, and lithium oxide if introduced to fire. Concerns about the hazardous nature of lithium sulfur oxide batteries waste were raised as lithium batteries became more obtainable. The US Environmental Protection Agency and Department of Defense evaluated the lithium sulfur oxide batteries and concluded that LiCN formation was one of the compounds leading to the hazardous waste.[6][7]

References[edit]

  1. ^ J. A. Lely,, J. M. Bijvoet (1942), "The Crystal Structure of Lithium Cyanide", Recueil des Travaux Chimiques des Pays-Bas, 61, London: WILEY-VCH Verlag 
  2. ^ Haynes, W.M (2013), "Bernard Lewis", in Bruno, Thomas., Handbook of Chemistry and Physics (93 ed.), Boca Raton, Florida: Fitzroy Dearborn 
  3. ^ Material Safety Data Sheet: Lithium Cyanide, 0.5M Solution in N,N-Dimethylformamide, Fisher Scientific, 16 June 1999 
  4. ^ a b "Cyanides". E. I. du Pont de Nemours & Co., Inc. Retrieved 2012-11-02. 
  5. ^ "Non-aqueous cyanation of halides using lithium cyanide". Elsevier. Retrieved 2012-10-17. 
  6. ^ a b "Regulatory status of spent and/or discarded lithium-sulfur dioxide (Li/S02) batteries". United States Environmental Protection Agency. 7 March 1984. Retrieved 2017-05-18. 
  7. ^ "Evaluation of Lithium Sulfur Dioxide Batteries" (PDF). U.S. Army Communications - Electronics Command and U.S. Army Electronics Research and Development Command. Retrieved 2012-10-23. 
Salts and covalent derivatives of the cyanide ion
HCN He
LiCN Be(CN)2 B C NH4CN OCN,
-NCO
FCN Ne
NaCN Mg(CN)2 Al(CN)3 SiCN P(CN)3 SCN,
-NCS,
(SCN)2,
S(CN)2
ClCN Ar
KCN Ca(CN)2 Sc(CN)3 Ti(CN)4 VO(CN)3 Cr(CN)3 Mn(CN)2 Fe(CN)3,
Fe(CN)64−,
Fe(CN)63−
Co(CN)2,
Co(CN)3
Ni(CN)2
Ni(CN)42−
CuCN Zn(CN)2 Ga(CN)3 Ge As(CN)3 SeCN
(SeCN)2
Se(CN)2
BrCN Kr
RbCN Sr(CN)2 Y(CN)3 Zr(CN)4 Nb Mo Tc Ru Rh Pd(CN)2 AgCN Cd(CN)2 In(CN)3 Sn Sb Te(CN)2,
Te(CN)4
ICN XeCN
CsCN Ba(CN)2   Hf Ta W Re Os Ir Pt Au Hg2(CN)2,
Hg(CN)2
TlCN Pb(CN)2 Bi(CN)3 Po At Rn
Fr Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
La Ce(CN)3,
Ce(CN)4
Pr Nd Pm Sm Eu Gd(CN)3 Tb Dy Ho Er Tm Yb Lu
Ac Th Pa UO2(CN)2 Np Pu Am Cm Bk Cf Es Fm Md No Lr