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Cyanogen
Cyanogen
Cyanogen
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Cyanogen
Skeletal formula of cyanogen
Skeletal formula of cyanogen
Ball and stick model of cyanogen
Ball and stick model of cyanogen
Spacefill model of cyanogen
Spacefill model of cyanogen
Names
Preferred IUPAC name
Oxalonitrile[4]
Systematic IUPAC name
Ethanedinitrile[4]
Other names
  • Cyanogen
  • Bis(nitridocarbon)(C—C)[1]
  • Dicyan[2][3]
  • Carbon nitride[2]
  • Oxalic acid dinitrile[3]
  • Dicyanogen
  • Nitriloacetonitrile
  • CY
Identifiers
3D model (JSmol)
1732464
ChEBI
ChemSpider
ECHA InfoCard 100.006.643 Edit this at Wikidata
EC Number
  • 207-306-5
1090
MeSH cyanogen
RTECS number
  • GT1925000
UNII
UN number 1026
  • InChI=1S/C2N2/c3-1-2-4 checkY
    Key: JMANVNJQNLATNU-UHFFFAOYSA-N checkY
  • N#CC#N
Properties
N≡C−C≡N
Molar mass 52.036 g·mol−1
Appearance Colourless gas
Odor pungent, bitter almond-like
Density 950 mg/mL (at −21 °C)
Melting point −28 °C (−18 °F; 245 K)
Boiling point −21.1 °C; −6.1 °F; 252.0 K
45 g/100 mL (at 20 °C)
Solubility soluble in ethanol, ethyl ether
Vapor pressure 5.1 atm (21 °C)[5]
1.9 μmol/(Pa·kg)
−21.6·10−6 cm3/mol
1.327 (18 °C)
Thermochemistry
241.57 J/(K·mol)
309.07 kJ/mol
−1.0978–−1.0942 MJ/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 forms cyanide in the body; flammable[5]
GHS labelling:
GHS02: FlammableGHS06: ToxicGHS09: Environmental hazard
Danger
H220, H331, H410
P210, P261, P271, P273, P304+P340, P311, P321, P377, P381, P391, P403, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 4: Will rapidly or completely vaporize at normal atmospheric pressure and temperature, or is readily dispersed in air and will burn readily. Flash point below 23 °C (73 °F). E.g. propaneInstability 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g. white phosphorusSpecial hazards (white): no code
4
4
2
Explosive limits 6.6–32%[5]
NIOSH (US health exposure limits):
PEL (Permissible)
 none[5]
REL (Recommended)
 TWA 10 ppm (20 mg/m3)[5]
IDLH (Immediate danger)
 N.D.[5]
Safety data sheet (SDS) inchem.org
Related compounds
Related alkanenitriles
Related compounds
 DBNPA
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Cyanogen is the chemical compound with the formula (CN)2. Its structure is N≡C−C≡N. The simplest stable carbon nitride, it is a colorless and highly toxic gas with a pungent odor. The molecule is a pseudohalogen. Cyanogen molecules are linear, and consist of two CN groups ‒ analogous to diatomic halogen molecules, such as Cl2, but far less oxidizing. The two cyano groups are bonded together at their carbon atoms, though other isomers have been detected.[6] The name is also used for the CN radical,[7] and hence is used for compounds such as cyanogen bromide (Br−C≡N)[8] (but see also Cyano radical). When burned at increased pressure with oxygen, it is possible to get a blue tinted flame, the temperature of which is about 4,800 °C (8,670 °F) (a higher temperature is possible with ozone). It is as such regarded as the gas with the second highest temperature of burning (after dicyanoacetylene).

Cyanogen is the anhydride of oxamide:

H2N−C(=O)−C(=O)−NH2 → N≡C−C≡N + 2 H2O

Oxamide is manufactured from cyanogen by hydration:[9]

N≡C−C≡N + 2 H2O → H2N−C(=O)−C(=O)−NH2

Preparation

[edit]

Cyanogen is typically generated from cyanide compounds. One laboratory method entails thermal decomposition of mercuric cyanide:

2 Hg(CN)2 → (CN)2 + Hg2(CN)2

Or, one can combine solutions of copper(II) salts (such as copper(II) sulfate) with cyanides; an unstable copper(II) cyanide is formed which rapidly decomposes into copper(I) cyanide and cyanogen.[10]

2 CuSO4 + 4 KCN → (CN)2 + 2 CuCN + 2 K2SO4

Industrially, it is created by the oxidation of hydrogen cyanide, usually using chlorine over an activated silicon dioxide catalyst or nitrogen dioxide over a copper salt. It is also formed when nitrogen and acetylene are reacted by an electrical spark or discharge.[11]

Reactions

[edit]

For the two less stable isomers of cyanogen, the order of the atoms differs. Isocyanogen (or cyanogen cyanide) is −C≡N+−C≡N.[12] It has been detected in the interstellar medium.[13]

Addition of disulfur dichloride to cyanogen gives 3,4-dichloro-1,2,5-thiadiazole.

Paracyanogen

[edit]

Paracyanogen is a polymer of cyanogen. It can be best prepared by heating mercury(II) cyanide. It can also be prepared by heating silver cyanide, silver cyanate, cyanogen iodide or cyanuric iodide.[14] It can also be prepared by the polymerization of cyanogen at 300 to 500 °C (572 to 932 °F) in the presence of trace impurities. Paracyanogen can also be converted back to cyanogen by heating to 800 °C (1,470 °F).[9] Based on experimental evidence, the structure of this polymeric material is thought to be rather irregular, with most of the carbon atoms being of sp2 type and localized domains of π conjugation.[15]

History

[edit]

Cyanogen was first synthesized in 1815 by Joseph Louis Gay-Lussac, who determined its empirical formula and named it. Gay-Lussac coined the word "cyanogène" from the Greek words κυανός (kyanos, blue) and γεννάω (gennao, to create), because cyanide was first isolated by Swedish chemist Carl Wilhelm Scheele from the pigment Prussian blue.[16] It attained importance with the growth of the fertilizer industry in the late 19th century and remains an important intermediate in the production of many fertilizers. It is also used as a stabilizer in the production of nitrocellulose.

Cyanogen is commonly found in comets.[17] In 1910 a spectroscopic analysis of Halley's Comet found cyanogen in the comet's tail, which led to public fear that the Earth would be poisoned as it passed through the tail. People in New York wore gas masks, and merchants sold quack "comet pills" claimed to neutralize poisoning.[17] Because of the extremely diffuse nature of the tail, there was no effect when the planet passed through it.[18][19]

Safety

[edit]

Like other cyanides, cyanogen is very toxic, as it readily undergoes reduction to cyanide, which poisons the cytochrome c oxidase complex, thus interrupting the mitochondrial electron transfer chain. Cyanogen gas is an irritant to the eyes and respiratory system. Inhalation can lead to headache, dizziness, rapid pulse, nausea, vomiting, loss of consciousness, convulsions, and death, depending on exposure.[20] Lethal dose through inhalation typically ranges from 100 to 150 milligrams (1.5 to 2.3 grains).

Cyanogen produces the second-hottest-known natural flame (after dicyanoacetylene aka carbon subnitride) with a temperature of over 4,525 °C (8,177 °F) when it burns in oxygen.[21][22]

See also

[edit]

References

[edit]
[edit]
Revisions and contributorsEdit on WikipediaRead on Wikipedia

Cyanogen

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Cyanogen is a highly toxic, colorless gas with the (CN)₂, also known as dicyanogen or oxalonitrile, representing the simplest stable . Its molecular structure is linear, consisting of two groups linked by a (N≡C–C≡N), and it has a molecular weight of 52.03 g/mol. At standard conditions, cyanogen exhibits a pungent, almond-like and boils at −21 °C (−6 °F), making it a volatile substance that forms mixtures with air. It can be prepared industrially by the oxidation of (HCN) using agents such as oxygen or , though this process requires careful control due to its reactivity. Cyanogen's primary applications include its use as a chemical intermediate in the synthesis of other compounds, such as in the of certain polymers and pharmaceuticals, and historically as a fumigant for agricultural and storage purposes. Due to its extreme , exposure to cyanogen primarily occurs through , leading to rapid absorption and conversion to ions in the body, which inhibit by binding to . Symptoms of acute include , , , , and potentially or , with a of 10 ppm (8-hour TWA) in occupational settings to mitigate risks. Combustion of cyanogen produces highly toxic fumes, including , , and nitrogen oxides, further emphasizing the need for stringent safety protocols in handling.

Properties

Physical Properties

Cyanogen (C₂N₂) is a colorless, poisonous gas that exhibits a pungent resembling that of almonds or . It has a of -27.9 °C and a of -21.2 °C. The density of cyanogen gas is 2.32 g/L at 0 °C and 1 atm. Cyanogen shows slight in , approximately 1.5 g/100 mL at 20 °C, while it is more soluble in organic solvents such as (about 23 g/100 g at 20 °C) and (about 5 g/100 g at 20 °C). of cyanogen reveals characteristic absorption bands at 2150 cm⁻¹, corresponding to the C≡N stretching vibration, and around 800 cm⁻¹, associated with bending modes. shows an absorption maximum at 219 nm in the gas phase. The critical temperature of cyanogen is 128.3 °C, and the critical pressure is 60.8 bar (approximately 60 ).

Molecular Structure

Cyanogen exhibits a with the connectivity N≡C–C≡N, consisting of a central C–C linking two cyano (–CN) groups. This structure arises from the sp hybridization of the carbon atoms, resulting in a bond angle of 180° and belonging to the D∞h point group. The linearity is corroborated by , which reveals symmetric top rotational constants consistent with a centrosymmetric arrangement. Experimental bond lengths, determined via rotational , measure the C≡N triple bonds at approximately 1.16 and the central C–C bond at 1.38 . The C–C bond length is longer than that of a standard aliphatic C–C (1.54 ) yet reflects the influence of sp hybridization, which increases bond strength and shortens distances relative to sp³ systems while remaining extended compared to multiple bonds due to the single-bond order. The electronic configuration of cyanogen is analyzed through , accommodating 18 valence electrons in a framework of σ and π orbitals formed from the atomic p orbitals of carbon and . The highest occupied (HOMO) corresponds to a bonding π orbital (1πu), which delocalizes along the molecular axis and contributes to the overall bonding stability. Due to its centrosymmetric linear structure, cyanogen possesses a dipole moment of essentially zero, as confirmed by the absence of a permanent electric in spectroscopic observations, further validating the symmetric N≡C–C≡N arrangement. The bonding in cyanogen is best represented by resonance structures, with the primary form being N≡C–C≡N and significant contributions from charge-separated zwitterionic forms such as ⁻N≡C–C≡N⁺ (and its symmetric equivalent ⁺N≡C–C≡N⁻). These resonance hybrids account for the partial double-bond character in the C–C linkage and the observed bond lengths, as electron delocalization weakens the central bond relative to a pure .

Synthesis

Laboratory Preparation

Cyanogen is commonly prepared in settings using small-scale methods that prioritize and control, given its and tendency to polymerize. The classic approach involves the of oxamide with as the dehydrating agent. In this reaction, oxamide (\ce(CONH2)2)( \ce{(CONH2)2} ) is heated with \ceP4O10\ce{P4O10} at temperatures of 200–300 °C, yielding cyanogen gas \ce(CN)2\ce{(CN)2} and water according to the equation \ce(CONH2)2−>(CN)2+2H2O\ce{(CONH2)2 -> (CN)2 + 2 H2O}. Typical yields range from 70% to 80%, with the elevated temperature range selected to facilitate while minimizing unwanted of the product. An alternative traditional method employs the of mercuric cyanide. Heating dry \ceHg(CN)2\ce{Hg(CN)2} to approximately 300–400 °C decomposes it into cyanogen gas and mercury, as shown by \ceHg(CN)2−>(CN)2+Hg\ce{Hg(CN)2 -> (CN)2 + Hg}. This process, first described in the early , generates the gas directly but requires careful handling due to the of mercury compounds and the need for inert atmospheres to prevent side reactions. Yields can approach 80% under optimized conditions, though the method is less favored today owing to environmental concerns with mercury. More contemporary laboratory synthesis utilizes electrochemical oxidation of cyanide ions in aqueous media. At a platinum electrode, cyanide anions are oxidized to cyanogen via the half-reaction \ce2CN−−>(CN)2+2e−\ce{2CN^- -> (CN)2 + 2e^-}, typically conducted in a divided cell with controlled potential (around 0.8–1.0 V vs. standard hydrogen electrode) to suppress hydrogen evolution. This method offers clean generation without solid byproducts and yields exceeding 70%, making it suitable for analytical or small-scale applications. The reaction is diffusion-controlled and proceeds efficiently in neutral to alkaline solutions. Regardless of the synthesis route, purification of cyanogen is essential to remove impurities such as (HCN), which forms as a common byproduct. The gas is typically collected over mercury or in a cold trap and then subjected to under reduced pressure (e.g., 50–100 mmHg) at low temperatures (below 0 °C) to exploit differences in volatility—cyanogen boils at -21 °C, while HCN boils at 26 °C. This step ensures high purity (>95%) for subsequent use, with traps often employed to capture residual HCN safely.

Industrial Production

Cyanogen is produced industrially primarily through the oxidation of (HCN). One common method involves the gas-phase oxidation of HCN with oxygen over a silver catalyst at temperatures around 500–600 °C, yielding (CN)₂ along with and oxides as byproducts, per the reaction 2 HCN + ½ O₂ → (CN)₂ + H₂O. This process operates continuously in specialized reactors with precise control to manage exothermic reactions and prevent explosions. Alternatively, chlorination of HCN using gas at elevated temperatures (approximately 400–500 °C) can produce cyanogen and , though it requires subsequent purification to remove chlorine residues. These methods are conducted on a large scale for use as intermediates in chemical manufacturing, with safety measures including inert diluents to avoid ignition.

Chemical Reactivity

Hydrolysis and Nucleophilic Reactions

Cyanogen exhibits significant reactivity toward nucleophiles due to the electron-deficient carbon atoms in its C≡N bonds, facilitating addition and subsequent transformation. Hydrolysis of cyanogen proceeds differently under acidic and basic conditions, reflecting the compound's sensitivity to pH. Under acidic conditions, cyanogen undergoes hydrolysis to form oxamide as the primary product. The reaction is (CN)_2 + 2 H_2O \to (NH_2)_2C_2O_2, often catalyzed by hydrohalic acids or organic Lewis acids such as acetaldehyde. This process is industrially relevant for oxamide production, a nitrogen fertilizer, and typically occurs in the liquid phase at moderate temperatures (<70 °C) and pressures (~3 bar). The mechanism begins with nucleophilic attack by water on one cyano carbon, forming an intermediate like cyanoformamide (NC-CH(OH)NH_2), followed by proton transfer, tautomerization, and addition to the second cyano group to yield oxamide. Without catalysts, the reaction yields complex mixtures including HCN, CO_2, and NH_3. In basic media, hydrolysis follows a distinct pathway involving hydroxide ion addition. The overall reaction is (CN)_2 + 2 OH^- \to CN^- + OCN^- + H_2O, proceeding quantitatively to cyanide and cyanate ions. The mechanism involves initial nucleophilic attack by OH^- on one cyano carbon, forming the adduct ^-N=C-C(OH)=NH (or resonance form N≡C-C(=O)NH^-), which then undergoes C-C bond cleavage. This cleavage occurs via two competitive paths: direct fragmentation (29% at 25 °C) to yield CN^- and HNCO (which hydrolyzes to OCN^-), or rearrangement to 1-cyanoformamide (NC-CH=O NH_2, 71% pathway), followed by deprotonation (pK_a = 10.8) and slower cleavage to the same products. The reaction is second-order overall, with rate constants for the initial addition step of k_1 = 8.9 \times 10^2 M^{-1} s^{-1} and k_2 = 2.17 \times 10^3 M^{-1} s^{-1} at 25 °C; activation energies are 66 kJ/mol (k_1) and 58 kJ/mol (k_2) for the addition, and 94 kJ/mol for the 1-cyanoformamide decomposition. Beyond hydrolysis, cyanogen participates in nucleophilic additions with amines, targeting the cyano groups to form imines, amidines, or related derivatives. For instance, primary amines (RNH_2) add across the molecule, with the general reaction (CN)_2 + 2 RNH_2 \to (RNH)_2C=NH + HCN, where the first amine adds to one C≡N bond forming an imine intermediate, followed by a second addition and elimination of HCN. This proceeds via nucleophilic attack on the carbon, proton transfer, and tautomerization to the amidine. Such reactions highlight cyanogen's utility in synthesizing nitrogen-rich heterocycles, though they require controlled conditions due to competing hydrolysis. The mechanism aligns with general nitrile nucleophilic additions, emphasizing the linear structure's role in enabling sequential attacks.

Polymerization and Thermal Decomposition

Cyanogen undergoes thermal at temperatures of 300–400 °C under atmospheric pressure, forming paracyanogen, a black polymer represented as [–C≡N–]n. This process can occur at lower temperatures, such as 170 °C, when elevated pressures around 300 atmospheres are applied. The polymerization is pressure-dependent, with higher pressures promoting the formation of the polymer. The mechanism involves a free-radical chain process initiated by homolysis of the central C–C bond in cyanogen, generating radicals. These radicals propagate the chain through addition to additional cyanogen molecules, forming unstable intermediates like (CN)3 that further polymerize. The C–C bond dissociation energy in cyanogen is approximately 680 kJ/mol at 0 K, reflecting the energy barrier for radical initiation. At much higher temperatures (e.g., above ~1200 °C), particularly under low-pressure or conditions, cyanogen undergoes over , dissociating into two CN• radicals via C–C bond homolysis: (CN)2 → 2 CN•. The resulting CN• radicals may recombine to reform cyanogen or undergo further fragmentation and reactions. Detailed studies indicate significant dissociation occurs around 1200 °C, with rates increasing markedly in shock-heated environments above 2500 K. Yields of paracyanogen from thermal can reach 30–45% under controlled conditions, though optimization varies with and duration.

Paracyanogen

Paracyanogen is a polymeric derivative of cyanogen, appearing as an insoluble black or brown powder with the (CN)_n. It forms via the of cyanogen gas, typically under thermal or photochemical conditions. This material has been recognized since the early as a stable, composed of and in a 1:1 ratio. The structure of paracyanogen is complex and not fully resolved, but it is generally described as a of linear and cyclic poly(cyano) chains incorporating conjugated double-bond systems. reveals characteristic bands around 1570 cm^{-1} indicative of these conjugated bonds, while theoretical models suggest possible layer-like lattices resembling , with rings where every second carbon atom is replaced by . Experimental evidence from and computational simulations supports the presence of both chain-like and ring-based motifs, contributing to its polymeric nature. Key physical properties include a of approximately 2.0 g/cm³, rendering it a relatively dense comparable to some carbon-based polymers. It decomposes thermally above 500 °C (773 ) without undergoing melting, releasing cyanogen gas and eventually yielding carbon and nitrogen at higher temperatures around 860 °C (1133 ). Paracyanogen is insoluble in , most organic solvents, and even liquid , though it shows partial in concentrated strong acids such as perchloric, sulfuric, or . Electrically, paracyanogen behaves as a , with conductivity values ranging from 10^{-3} to 10^{-11} Ω^{-1} cm^{-1} depending on preparation and form. Thin films exhibit semimetal-like properties with near-zero for conduction, and optical studies estimate a bandgap of approximately 2.8 eV, positioning it as an early subject in research on organic and conjugated semiconductors. Its conjugated facilitates delocalization, akin to polyacenes, making it relevant for understanding charge transport in nitrogen-containing polymers. In terms of stability, paracyanogen is resistant to dilute acids and bases at ambient conditions but reacts with strong bases like KOH or Na₂CO₃ upon fusion, evolving . It demonstrates good and stability, withstanding γ-ray doses up to 5 × 10^7 J kg^{-1}, yet it slowly degrades in air over periods of years through oxidation, forming colored degradation products.

Cyanogen Halides

Cyanogen halides are a class of compounds with the general XCN, where X represents a atom (Cl, Br, or I). These pseudohalides exhibit properties analogous to cyanogen (N≡C–C≡N) but incorporate a halogen in place of one cyano group, rendering them more stable and versatile in synthetic applications. Cyanogen chloride (ClCN) is a colorless, highly volatile gas or liquid with a boiling point of approximately 13 °C and a pungent, irritating odor. It is synthesized industrially by the direct of : HCN+Cl2→ClCN+HCl\text{HCN} + \text{Cl}_2 \rightarrow \text{ClCN} + \text{HCl} This reaction proceeds exothermically and requires careful control to manage the corrosive byproducts. ClCN has been employed as a due to its potent lachrymatory effects, causing severe irritation to the eyes, , and mucous membranes, though its use in this capacity is limited by its high toxicity. Cyanogen bromide (BrCN) exists as colorless, deliquescent crystals with a of 52 °C. It is prepared similarly through of HCN: HCN+Br2→BrCN+HBr\text{HCN} + \text{Br}_2 \rightarrow \text{BrCN} + \text{HBr} In biochemical applications, BrCN is widely used for specific cleavage of polypeptide chains at residues, facilitating and structural analysis by generating defined fragments. This cleavage occurs via an SN2 mechanism where the sulfur of attacks the carbon of the cyano group, leading to homoserine formation. Cyanogen iodide (ICN), the iodine analog, forms white needles with a strong pungent odor and is less volatile than its lighter counterparts. It shares the general preparation method via HCN but is more light-sensitive and typically handled with stabilizers. ICN is employed in niche applications such as preservatives in , though its reactivity limits broader use. These compounds behave as pseudohalides, mimicking the reactivity of ions in many contexts due to the -CN group's ability to form strong bonds. For instance, ClCN undergoes in : ClCN+H2O→HCN+HOCl\text{ClCN} + \text{H}_2\text{O} \rightarrow \text{HCN} + \text{HOCl} This reaction highlights their tendency to release species under mild conditions, contributing to their utility in for introducing cyano or halo functionalities. Cyanogen halides are more thermally stable than cyanogen itself, which decomposes readily above 0 °C, but they remain highly toxic, primarily through liberation of upon or . Exposure leads to symptoms akin to , including respiratory distress, cardiovascular collapse, and central nervous system depression, with a probable human oral for BrCN of less than 5 mg/kg. Their handling requires stringent safety protocols, including fume hoods and cyanide antidotes.

History and Discovery

Early Isolation

Cyanogen was first synthesized and isolated in 1815 by French chemist through the of mercury(II) cyanide (Hg(CN)2) or (AgCN) at temperatures around 400 °C, yielding the metal and the gaseous product (CN)2. Gay-Lussac's experiments demonstrated that this gas combined with oxygen to form cyanic acid and related salts, establishing its role as a distinct chemical entity related to prussic acid (). The name "cyanogen" was coined by Gay-Lussac from the Greek roots "kyanos" (referring to the blue pigment in Prussian blue, an iron cyanide complex) and "gennan" (to produce), reflecting its ability to generate cyanide-containing compounds that produce the characteristic blue coloration. Early chemical analysis involved combustion of the gas, revealing an elemental composition of equal parts carbon and nitrogen by weight, consistent with the empirical formula CN. Vapor density measurements, relative to hydrogen or air, yielded a molecular weight of approximately 52 g/mol, supporting the diatomic structure C2N2 and distinguishing it from hydrogen cyanide (HCN, molecular weight 27 g/mol). The compound's high instability posed significant challenges during early isolations; it readily polymerized to paracyanogen or decomposed under moisture or heat, resulting in impure samples often contaminated with HCN, which led to initial confusions in identification and characterization.

Key Developments

In the early 1930s, the molecular structure of cyanogen was elucidated through studies, confirming its linear configuration as N≡C–C≡N with bond lengths of approximately 1.16 for C≡N and 1.38 for C–C, influenced by Pauling's work on . During , cyanogen was investigated as a precursor for agents, particularly (CK), which the U.S. Army tested for its properties and ability to penetrate gas masks; however, these efforts were not pursued to deployment. In the , spectroscopic investigations using revealed vibrationally excited cyanogen radicals (CN) as key intermediates in the decomposition of cyanogen halides, with absorption spectra showing sequences up to v'' = 6, providing insights into the molecule's photochemical reactivity. Synthetic methods for cyanogen advanced in the mid-20th century, enabling larger-scale production, though details are covered in the synthesis section. Post-2000 research has employed (DFT) to model cyanogen's electronic bonding, highlighting the σ-donation and π-backbonding in its metal complexes and predicting bond dissociation energies around 133 kcal/mol for the central C–C linkage. Additionally, cyanogen has played a role in ; the CN radical, derived from cyanogen, was detected in in the –1940s, and cyanogen itself was detected in the coma of comet 67P/Churyumov-Gerasimenko by the mission in 2015 at abundances of about 0.02% relative to water.

Applications and Uses

Chemical Synthesis

Cyanogen ((CN)2) reacts with organometallic compounds, such as two equivalents of RM (where R is an alkyl or and M is a metal like ), to yield a mixture of glycinonitriles RR′C(CN)NH₂, ketones RCOR′, and tertiary alcohols R₃COH. Organolithium compounds favor elimination pathways leading to ketones. This approach utilizes the dual CN units in cyanogen, though it is less common than other methods for introducing cyano functionality due to the complexity of products. In heterocycle synthesis, cyanogen reacts with conjugated dienes like 1,3-butadiene in a cycloaddition manner, forming 2-cyanopyridine after dehydrogenation, typically at high temperatures (400–600 °C) in the vapor phase. Conversions are low, around 1% or less, offering a route to nitrogen-containing aromatics but limited by efficiency compared to other syntheses. Cyanogen also plays a role in coordination chemistry, reacting with metal cyanides to form mixed-ligand complexes that incorporate both dicyano and cyanogen units. For example, cyanogen interacts with copper(I) cyanide solutions to generate the [Cu(CN)2(NCCN)]– anion, where the cyanogen ligand binds end-on via nitrogen, bridging to produce structures akin to metal dicyanides [M(CN)2]. These complexes serve as precursors for extended metal-cyanide networks, with the linear NCCN unit facilitating assembly into polymeric or cluster motifs.

Industrial and Other Applications

Cyanogen has been employed historically as a fumigant for storage and treatment due to its broad-spectrum efficacy against , nematodes, fungi, and seeds, though its application has been limited by high concerns leading to phased-out or restricted use in favor of alternatives like phosphine. The oxy-cyanogen flame has been studied for producing one of the hottest known chemical flames at approximately 4,525 °C (4800 K), which theoretically enables processing of heat-resistant metals beyond oxy-acetylene capabilities, but it is not utilized in practical or metal cutting due to cyanogen's . Within , cyanogen serves as a reference compound and intermediate in assays for detection, such as methods for quantifying cyanogen and related species like in environmental samples. Cyanogen holds potential in space applications as a high-energy-density component in rocket propellants, where its oxygen combustion products have been studied for propulsion efficiency in missiles and upper-stage engines. As of 2020, cyanogen production remains niche, derived primarily from the oxidation or chlorination of (HCN) and accounting for a small fraction—estimated at less than 5%—of global HCN output, with use curtailed by environmental and toxicity regulations implemented since the 1990s under frameworks like the U.S. Amendments and Reauthorization Act.

Safety and Toxicology

Health Hazards

Cyanogen exerts its primary toxicity through hydrolysis in biological tissues to (HCN) and ions, with HCN subsequently binding to the ferric iron in , thereby inhibiting the terminal enzyme of the mitochondrial and blocking . This disruption prevents (ATP) production, leading to rapid onset of cellular hypoxia, particularly in oxygen-dependent tissues such as the brain and heart. Upon absorption, cyanogen is metabolized primarily in the liver, where it undergoes to HCN and ; the HCN is then detoxified by the enzyme rhodanese (thiosulfate sulfurtransferase) via sulfur donation from , forming the less toxic ion, which is excreted in the . The biological of from such sources is approximately 20 minutes to 1 hour in humans, though efficiency depends on sulfur substrate availability. Acute exposure to in humans can cause of the eyes, , and throat, along with systemic symptoms including headache, giddiness, nausea, fatigue, and rapid breathing due to the released . In animal studies, rats exposed to cyanogen via exhibit gasping, tremors, convulsions, and , with an LC50 of 350 ppm for 1 hour. No mortality was observed in rats at 500 ppm for 30 minutes, indicating cyanogen's acute lethality is somewhat less potent than HCN on a concentration-time basis. Chronic low-level exposure to cyanogen may lead to thyroid disruption through accumulation of thiocyanate, which competitively inhibits iodine uptake by the gland, potentially causing goiter or . Symptoms such as mild and breathing difficulties have been reported in workers with prolonged exposure averaging 8 ppm over years. The odor of cyanogen is not detectable below approximately 235 ppm, though irritation of the and eyes can occur at concentrations as low as 16 ppm for 6-8 minutes; may limit ongoing detection at higher levels, reducing its utility as a warning signal.

Handling and Regulations

Cyanogen is typically stored in cylinders as a under its own , in cool, well-ventilated areas at temperatures not exceeding 52 °C to minimize the risk of and . Safe handling of cyanogen requires operations to be performed in a fume hood with self-contained breathing apparatus (SCBA) to protect against inhalation exposure, given its high toxicity. As a flammable gas with explosive limits ranging from 6.6% to 43% in air, all potential ignition sources such as sparks, open flames, or hot surfaces must be strictly avoided during transfer and use. In emergencies involving exposure, cyanogen is managed as a cyanide precursor, with antidotes including hydroxocobalamin (preferred first-line treatment) and the cyanide antidote kit containing sodium nitrite and sodium thiosulfate, administered intravenously to bind and detoxify cyanide. Decontamination of skin, equipment, or minor spills involves using a dilute bleach (sodium hypochlorite) solution to oxidize any released cyanide ions into less toxic forms. Regulatory standards for occupational exposure to cyanogen include a NIOSH of 10 ppm (20 mg/m³) as an 8-hour time-weighted average, while OSHA has no established . In the , cyanogen is classified as very toxic under the previous system (T+ symbol, with risk phrases R26/27/28 for toxicity by , contact, and ). Under the current , it is classified as a flammable gas (category 1; H220), acutely toxic by (category 3; H331), and hazardous to the aquatic environment (acute category 1, H400; chronic category 1, H410). Indirect regulatory connections to the arise from cyanogen's role as a non--depleting alternative fumigant to controlled substances like methyl bromide. Spill or leak response protocols emphasize immediate evacuation of the affected area for at least 100 meters in all directions, followed by enhanced ventilation to disperse the heavier-than-air gas. Any accumulated liquid should be neutralized cautiously with a (NaOH) solution to form stable, less volatile compounds before disposal.

References

  1. https://pubchem.ncbi.nlm.nih.gov/compound/Cyanogen
  2. https://webbook.nist.gov/cgi/cbook.cgi?ID=460-19-5
  3. https://www.osha.gov/chemicaldata/372
  4. https://cameochemicals.noaa.gov/chemical/490
  5. https://cloudsds.com/chemical-hazard-and-safety/comprehensive-guide-to-cyanogen-c%25E2%2582%2582n%25E2%2582%2582-usage-safety-measures-and-potential-hazards/
  6. https://nj.gov/health/eoh/rtkweb/documents/fs/0554.pdf
  7. https://www.cdc.gov/niosh/npg/npgd0161.html
  8. https://www.inchem.org/documents/icsc/icsc/eics1390.htm
  9. https://nepis.epa.gov/Exe/ZyPURL.cgi?Dockey=9100SIHC.TXT
  10. https://www.chemicalbook.com/ProductChemicalPropertiesCB0809207_EN.htm
  11. https://webbook.nist.gov/cgi/cbook.cgi?ID=C460195&Mask=1EFF
  12. https://www.sciencedirect.com/science/article/abs/pii/0022286087800625
  13. https://pubs.aip.org/aip/jcp/article-pdf/102/2/1100/19198355/1100_1_online.pdf
  14. https://www.smolecule.com/products/s594424
  15. https://www.aanda.org/articles/aa/pdf/2019/05/aa35010-19.pdf
  16. https://analyticalsciencejournals.onlinelibrary.wiley.com/doi/abs/10.1002/jlcr.2580311211
  17. https://pubs.acs.org/doi/10.1021/cr50029a003
  18. https://www.chemicalbook.com/ChemicalProductProperty_EN_CB0809207.htm
  19. https://patents.google.com/patent/US2399361A/en
  20. https://www.sciencedirect.com/science/article/pii/0022072863800054
  21. https://patents.google.com/patent/US2793991A/en
  22. https://www.thieme-connect.com/products/ejournals/pdf/10.1055/s-0039-1691095.pdf
  23. https://patents.google.com/patent/US3037055A/en
  24. https://pubs.acs.org/doi/10.1021/ic00262a022
  25. https://patents.google.com/patent/US3057808A/en
  26. https://pubs.acs.org/doi/10.1021/j100283a029
  27. https://www.russchemrev.org/RCR2608pdf
  28. https://pubs.aip.org/aip/jcp/article/102/2/1100/479731/Determination-of-the-CC-bond-dissociation-energy
  29. https://ui.adsabs.harvard.edu/abs/1984JChPh..80.4982S/abstract
  30. https://www.freepatentsonline.com/4649038.html
  31. https://pubchem.ncbi.nlm.nih.gov/compound/Cyanogen-Chloride
  32. https://www.cdc.gov/niosh/ershdb/emergencyresponsecard_29750039.html
  33. https://cdn.who.int/media/docs/default-source/wash-documents/wash-chemicals/2007-cyanogen_background_document.pdf?sfvrsn=6321b6d3_3
  34. https://pubchem.ncbi.nlm.nih.gov/compound/Cyanogen-Bromide
  35. https://www.sciencedirect.com/topics/pharmacology-toxicology-and-pharmaceutical-science/cyanogen-bromide
  36. https://pubchem.ncbi.nlm.nih.gov/compound/Cyanogen-iodide
  37. https://nj.gov/health/eoh/rtkweb/documents/fs/2302.pdf
  38. https://cameochemicals.noaa.gov/chemical/488
  39. https://en.wikisource.org/wiki/1911_Encyclop%C3%A6dia_Britannica/Cyanogen
  40. https://pubs.acs.org/doi/10.1021/acs.inorgchem.5c01041
  41. https://www.pnas.org/doi/10.1073/pnas.19.9.868
  42. https://www.fireengineering.com/technical-rescue/blood-agent-cyanogen-chloride/
  43. https://www.aanda.org/articles/aa/full_html/2021/03/aa39580-20/aa39580-20.html
  44. https://www.sciencedirect.com/science/article/abs/pii/S0022328X00943158
  45. https://pubs.rsc.org/en/content/articlelanding/1949/jr/jr9490001479
  46. https://pubs.rsc.org/en/content/articlehtml/1968/j1/j19680002444
  47. https://www.epa.gov/sites/default/files/2014-11/documents/cyanogen_final_v17_jun2014.pdf
  48. https://pubs.aip.org/aip/jcp/article-pdf/20/3/369/18800137/369_1_online.pdf
  49. https://pubs.acs.org/doi/10.1021/ac60195a039
  50. https://www.epa.gov/sites/default/files/2015-06/documents/epa-335.4.pdf
  51. https://ntrs.nasa.gov/api/citations/19640018010/downloads/19640018010.pdf
  52. https://www.epa.gov/sites/default/files/2020-11/documents/cyanide.pdf
  53. https://www.ncbi.nlm.nih.gov/books/NBK600901/
  54. https://www.airgas.com/msds/012226.pdf
  55. https://emedicine.medscape.com/article/814287-treatment
  56. https://wwwn.cdc.gov/TSP/MMG/MMGDetails.aspx?mmgid=1141&toxid=249
  57. https://echa.europa.eu/information-on-chemicals/cl-inventory-database/-/discli/details/100.006.643
  58. https://ozone.unep.org/sites/default/files/2019-05/TEAP_Assessment_report_2014.pdf
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