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Acetonitrile
Acetonitrile
Acetonitrile
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Acetonitrile
Skeletal formula of acetonitrile
Skeletal formula of acetonitrile
Skeletal formula of acetonitrile with all explicit hydrogens added
Skeletal formula of acetonitrile with all explicit hydrogens added
Ball and stick model of acetonitrile
Ball and stick model of acetonitrile
Spacefill model of acetonitrile
Spacefill model of acetonitrile
Names
Preferred IUPAC name
Acetonitrile[2]
Systematic IUPAC name
Ethanenitrile[2]
Other names
  • Cyanomethane[1]
  • Ethyl nitrile[1]
  • Methanecarbonitrile[1]
  • Methyl cyanide[1]
  • MeCN
  • ACN
Identifiers
3D model (JSmol)
741857
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.000.760 Edit this at Wikidata
EC Number
  • 200-835-2
895
MeSH acetonitrile
RTECS number
  • AL7700000
UNII
UN number 1648
  • InChI=1S/C2H3N/c1-2-3/h1H3 checkY
    Key: WEVYAHXRMPXWCK-UHFFFAOYSA-N checkY
  • CC#N
Properties
C2H3N
Molar mass 41.053 g·mol−1
Appearance Colorless liquid
Odor Faint, distinct, fruity
Density 0.786 g/cm3 at 25°C
Melting point −46 to −44 °C; −51 to −47 °F; 227 to 229 K
Boiling point 81.3 to 82.1 °C; 178.2 to 179.7 °F; 354.4 to 355.2 K
Miscible
log P −0.334
Vapor pressure 9.71 kPa (at 20.0 °C)
530 μmol/(Pa·kg)
Acidity (pKa) 25
UV-vis (λmax) 195 nm
Absorbance ≤0.10
−28.0×10−6 cm3/mol
1.344
3.92 D
Thermochemistry
91.69 J/(K·mol)
149.62 J/(K·mol)
40.16–40.96 kJ/mol
−1256.03 – −1256.63 kJ/mol
Hazards
GHS labelling:
GHS02: Flammable GHS07: Exclamation mark
Danger
H225, H302, H312, H319, H332
P210, P280, P305+P351+P338
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g. gasolineInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
2
3
0
Flash point 2.0 °C (35.6 °F; 275.1 K)
523.0 °C (973.4 °F; 796.1 K)
Explosive limits 4.4–16.0%
Lethal dose or concentration (LD, LC):
  • 2 g/kg (dermal, rabbit)
  • 2.46 g/kg (oral, rat)
5655 ppm (guinea pig, 4 hr)
2828 ppm (rabbit, 4 hr)
53,000 ppm (rat, 30 min)
7500 ppm (rat, 8 hr)
2693 ppm (mouse, 1 hr)[4]
16,000 ppm (dog, 4 hr)[4]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 40 ppm (70 mg/m3)[3]
REL (Recommended)
 TWA 20 ppm (34 mg/m3)[3]
IDLH (Immediate danger)
 500 ppm[3]
Related compounds
Related alkanenitriles
Related compounds
 DBNPA
Supplementary data page
Acetonitrile (data page)
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 ?)

Acetonitrile, often abbreviated MeCN (methyl cyanide), is the chemical compound with the formula CH3CN and structure H3C−C≡N. This colourless liquid is the simplest organic nitrile (hydrogen cyanide is a simpler nitrile, but the cyanide anion is not classed as organic). It is produced mainly as a byproduct of acrylonitrile manufacture. It is used as a polar aprotic solvent in organic synthesis and in the purification of butadiene.[5] The N≡C−C skeleton is linear with a short C≡N distance of 1.16 Å.[6]

Acetonitrile was first prepared in 1847 by the French chemist Jean-Baptiste Dumas.[7]

Applications

[edit]

Acetonitrile is used mainly as a solvent in the purification of butadiene in refineries. Specifically, acetonitrile is fed into the top of a distillation column filled with hydrocarbons including butadiene, and as the acetonitrile falls down through the column, it absorbs the butadiene which is then sent from the bottom of the tower to a second separating tower. Heat is then employed in the separating tower to separate the butadiene.

In the laboratory, it is used as a medium-polarity non-protic solvent that is miscible with water and a range of organic solvents, but not saturated hydrocarbons. It has a convenient range of temperatures at which it is a liquid, and a high dielectric constant of 38.8. With a dipole moment of 3.92 D,[8] acetonitrile dissolves a wide range of ionic and nonpolar compounds and is useful as a mobile phase in HPLC and LC–MS.

It is widely used in battery applications because of its relatively high dielectric constant and ability to dissolve electrolytes. For similar reasons, it is a popular solvent in cyclic voltammetry.

Its ultraviolet transparency UV cutoff, low viscosity and low chemical reactivity make it a popular choice for high-performance liquid chromatography (HPLC).

Acetonitrile plays a significant role as the dominant solvent used in oligonucleotide synthesis from nucleoside phosphoramidites.

Industrially, it is used as a solvent for the manufacture of pharmaceuticals and photographic film.[9]

Organic synthesis

[edit]

Acetonitrile is a common two-carbon building block in organic synthesis[10] of many useful chemicals, including acetamidine hydrochloride, thiamine, and 1-naphthaleneacetic acid.[11] Its reaction with cyanogen chloride affords malononitrile.[5]

As an electron pair donor

[edit]

Acetonitrile has a free electron pair at the nitrogen atom, which can form many transition metal nitrile complexes. For example, bis(acetonitrile)palladium dichloride is prepared by heating a suspension of palladium chloride in acetonitrile:[12]

PdCl2 + 2 CH3CN → PdCl2(CH3CN)2

A related complex is tetrakis(acetonitrile)copper(I) hexafluorophosphate [Cu(CH3CN)4]+. The CH3CN ligands in these complexes are rapidly displaced.

It also forms Lewis adducts with group 13 Lewis acids like boron trifluoride.[13] In superacids, it is possible to protonate acetonitrile.[14]

Production

[edit]

Acetonitrile is a byproduct from the manufacture of acrylonitrile by catalytic ammoxidation of propylene. Most is combusted to support the intended process but an estimated several thousand tons are retained for the above-mentioned applications.[15] Production trends for acetonitrile thus generally follow those of acrylonitrile. Acetonitrile can also be produced by many other methods, but these are of no commercial importance as of 2002. Illustrative routes are by dehydration of acetamide or by hydrogenation of mixtures of carbon monoxide and ammonia.[16] In 1992, 14,700 tonnes (16,200 short tons) of acetonitrile were produced in the US.

Acetonitrile shortage in 2008–2009

[edit]

Starting in October 2008, the worldwide supply of acetonitrile was low because Chinese production was shut down for the Olympics. Furthermore, a U.S. factory was damaged in Texas during Hurricane Ike.[17] Due to the global economic slowdown, the production of acrylonitrile used in acrylic fibers and acrylonitrile butadiene styrene (ABS) resins decreased. Acetonitrile is a byproduct in the production of acrylonitrile and its production also decreased, further compounding the acetonitrile shortage.[18] The global shortage of acetonitrile continued through early 2009.[needs update]

Safety

[edit]

Toxicity

[edit]

Acetonitrile has only modest toxicity in small doses.[11][19] It can be metabolised to produce hydrogen cyanide, which is the source of the observed toxic effects.[9][20][21] Generally the onset of toxic effects is delayed, due to the time required for the body to metabolize acetonitrile to cyanide (generally about 2–12 hours).[11]

Cases of acetonitrile poisoning in humans (or, to be more specific, of cyanide poisoning after exposure to acetonitrile) are rare but not unknown, by inhalation, ingestion and (possibly) by skin absorption.[20] The symptoms, which do not usually appear for several hours after the exposure, include breathing difficulties, slow pulse rate, nausea, and vomiting. Convulsions and coma can occur in serious cases, followed by death from respiratory failure. The treatment is as for cyanide poisoning, with oxygen, sodium nitrite, and sodium thiosulfate among the most commonly used emergency treatments.[20]

It has been used in formulations for nail polish remover, despite its toxicity. At least two cases have been reported of accidental poisoning of young children by acetonitrile-based nail polish remover, one of which was fatal.[22] Acetone and ethyl acetate are often preferred as safer for domestic use, and acetonitrile has been banned in cosmetic products in the European Economic Area since March 2000.[23]

Metabolism and excretion

[edit]
Compound Cyanide, concentration in brain (μg/kg) Oral LD50 (mg/kg)
Potassium cyanide 700 ± 200 10
Propionitrile 510 ± 80 40
Butyronitrile 400 ± 100 50
Malononitrile 600 ± 200 60
Acrylonitrile 400 ± 100 90
Acetonitrile 28 ± 5 2460
Table salt (NaCl) — 3000
Ionic cyanide concentrations measured in the brains of Sprague-Dawley rats one hour after oral administration of an LD50 of various nitriles.[24]

In common with other nitriles, acetonitrile can be metabolised in microsomes, especially in the liver, to produce hydrogen cyanide, as was first shown by Pozzani et al. in 1959.[25] The first step in this pathway is the oxidation of acetonitrile to glycolonitrile by an NADPH-dependent cytochrome P450 monooxygenase. The glycolonitrile then undergoes a spontaneous decomposition to give hydrogen cyanide and formaldehyde.[19][20] Formaldehyde, a toxin and a carcinogen on its own, is further oxidized to formic acid, which is another source of toxicity.

The metabolism of acetonitrile is much slower than that of other nitriles, which accounts for its relatively low toxicity. Hence, one hour after administration of a potentially lethal dose, the concentration of cyanide in the rat brain was 1⁄20 that for a propionitrile dose 60 times lower (see table).[24]

The relatively slow metabolism of acetonitrile to hydrogen cyanide allows more of the cyanide produced to be detoxified within the body to thiocyanate (the rhodanese pathway). It also allows more of the acetonitrile to be excreted unchanged before it is metabolised. The main pathways of excretion are by exhalation and in the urine.[19][20][21]

See also

[edit]

References

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

Acetonitrile

View on Grokipedia
from Grokipedia
Acetonitrile is a colorless, volatile, and with the CH₃CN (or C₂H₃N) and a molecular weight of 41.05 g/mol, known for its ether-like or aromatic and high in and most organic solvents. It has a of 81.6°C (179°F), a of -45.4°C (-49°F), a of 0.786 g/cm³ at 20°C, and a of 5.5°C (42°F), making it less dense than but capable of forming vapors when heated. Primarily produced as a byproduct of manufacturing through processes like the ammoxidation of , acetonitrile serves as a key industrial chemical with an annual global production exceeding 200,000 metric tons. As a , acetonitrile is extensively used in , , and (HPLC) due to its ability to dissolve a wide range of compounds without donating protons. It plays a critical role in the production of pharmaceuticals, agrochemicals, dyes, plastics, and batteries, including lithium-ion batteries, as well as in the extraction of hydrocarbons and the separation of fatty acids from oils. Additionally, it is employed in the manufacture of photographic films, perfumes, rubber products, and pesticides, and historically as a and alcohol denaturant. Despite its utility, acetonitrile is toxic and poses significant health and safety risks; it is harmful if swallowed, inhaled, or absorbed through the skin, with an odor threshold of 170 ppm, and metabolizes in the body to , potentially causing , , and . It reacts violently with strong oxidizing agents and requires careful handling in well-ventilated areas to mitigate and hazards from its low and vapor density greater than air. Environmentally, acetonitrile enters the atmosphere primarily from automobile exhaust and industrial emissions, where it can persist but is biodegradable under aerobic conditions.

Properties

Molecular structure

Acetonitrile has the CH₃CN and a molecular weight of 41.05 g/mol. The adopts a linear structure along its C-C≡N backbone, with the methyl carbon bonded to three atoms in a tetrahedral arrangement. The C-C bond length is approximately 1.46 , the C≡N length is approximately 1.16 , and the bond angles around the cyano group are approximately 180°, consistent with the sp hybridization of the cyano carbon. The carbon atom in the (CH₃-) is sp³ hybridized, forming four bonds: three C-H bonds and one C-C bond. In contrast, the cyano carbon is sp hybridized, utilizing two sp hybrid orbitals to form bonds with the methyl carbon and , while the remaining p orbitals form the pi bonds of the with . The atom is also sp hybridized, with its occupying an sp orbital. This electronic arrangement results in a significant dipole moment of 3.92 D, primarily arising from the electronegative atom pulling toward itself in the polar C≡N bond. The of acetonitrile is represented as H₃C–C≡N, where the cyano carbon shares one with the and a with ; the bears a of electrons. The group exhibits , with two major contributing structures that delocalize the : \chemfig{H_3C-C#N: <-> H_3C-C^{+}=[::60]N^{-}} This resonance stabilizes the molecule and enhances the polarity of the C≡N bond, with the cyano carbon bearing a partial positive charge and the nitrogen a partial negative charge.

Physical properties

Acetonitrile is a colorless liquid at room temperature, exhibiting a faint ether-like odor due to trace impurities or its inherent chemical nature. Its is 81.6 °C at standard , while the is -45.4 °C, indicating a relatively wide liquid range suitable for various and industrial applications. The of acetonitrile is 0.786 g/cm³ at 20 °C, which is lower than that of , contributing to its utility as a less dense . The of acetonitrile is 1.344 at 20 °C, a value that reflects its in spectroscopic analyses. Its is 0.34 cP at 25 °C, making it a low-viscosity fluid that flows easily. Additionally, the is 73 mmHg at 20 °C, and the heat of vaporization is 29.75 kJ/mol, parameters that influence its volatility and evaporation behavior in open systems. Thermodynamic properties include a of 91.7 J/mol·K for the liquid phase at 25 °C, which describes its capacity to store .
PropertyValueConditionsSource
AppearanceColorless liquid, ether-like odorRoom temperaturePubChem
81.6 °C1 atmNIST
-45.4 °C-NIST
0.786 g/cm³20 °CNIST
1.34420 °CNIST
0.34 cP25 °CNIST
73 mmHg20 °CNIST
Heat of vaporization29.75 kJ/molNIST
(liquid)91.7 J/mol·K25 °CNIST

Chemical properties

Acetonitrile features a (-C≡N), which is characterized by a between carbon and , enabling it to undergo reactions due to the electrophilic nature of the carbon atom in the cyano group. A prominent example is the of acetonitrile under acidic conditions, where it reacts with in the presence of to yield acetic acid and : CH3CN+2H2O+H2SO4→CH3COOH+NH4HSO4\text{CH}_3\text{CN} + 2\text{H}_2\text{O} + \text{H}_2\text{SO}_4 \rightarrow \text{CH}_3\text{COOH} + \text{NH}_4\text{HSO}_4 This reaction proceeds via initial addition of water to form an intermediate, followed by further , highlighting the reactivity of the group toward nucleophiles. Acetonitrile exhibits considerable , particularly resistance to oxidation under standard conditions, which contributes to its utility in various chemical environments. However, it is susceptible to in the presence of strong acids or bases, converting to or acetic acid depending on the conditions, as the is cleaved by nucleophilic attack. The pKa of its conjugate acid (CH₃CNH⁺) is approximately -10, indicating that acetonitrile is a very and occurs only under highly acidic conditions. In coordination chemistry, acetonitrile serves as a in complexes, coordinating through the on the atom of the cyano group to form stable bonds with metal centers. This behavior is observed in various complexes, such as those with or , where the ligand can be substituted or reduced while bound to the metal. Acetonitrile lacks geometric isomers due to its linear structure and simple molecular framework, with no possibility for cis-trans configurations. Tautomeric forms, such as potential imine-like structures, are negligible under normal conditions, as the energy barrier for such interconversions is prohibitively high in this aliphatic . Spectroscopically, acetonitrile is readily identified by characteristic signals from the group. In (IR) spectroscopy, the C≡N stretching vibration appears as a sharp, intense peak at approximately 2250 cm⁻¹, a hallmark of the . In (NMR) spectroscopy, the ¹H NMR spectrum shows a singlet at 2.0 ppm for the methyl protons, reflecting their equivalent environment. The ¹³C NMR spectrum displays signals at 1.3 ppm for the CH₃ carbon and 118.3 ppm for the CN carbon, with the latter shifted downfield due to the of the .

Production

Industrial synthesis

Acetonitrile is predominantly produced on an industrial scale as a of manufacturing through the ammoxidation of , a process developed by the of (Sohio) in the 1950s and now widely adopted globally. In this method, reacts with and oxygen over a metal oxide , typically phosphomolybdate, at temperatures of 400–500°C to yield as the main product, alongside and byproducts including and . The overall reaction for formation is CH₃CH=CH₂ + NH₃ + ³/₂ O₂ → CH₂=CHCN + 3 H₂O, with formed via parallel pathways involving and cyanation, resulting in co-production yields of 40–60 kg of per metric ton of , equivalent to approximately 4–6% by weight. This integrated process accounts for over 90% of global supply, tying its availability closely to the ~7 million metric tons annual production of for polymers and fibers. Historically, prior to the dominance of the Sohio process, acetonitrile was synthesized through the of , an derived from acetic acid and . The reaction, CH₃CONH₂ → CH₃CN + H₂O, was facilitated industrially using dehydrating agents such as or catalytic supports like alumina at elevated temperatures around 300–400°C. This method, employed in the mid-20th century, offered a dedicated route independent of production but was less economical due to higher raw material costs and energy demands, leading to its decline as the ammoxidation byproduct became the preferred source. Following extraction from the reaction mixture via absorption and , crude acetonitrile undergoes multi-stage purification primarily through to remove , , and other impurities. Azeotropic or pressure-swing is commonly used to break the acetonitrile- azeotrope, achieving final purities exceeding 99.9% suitable for applications. Global annual production of acetonitrile stands at approximately 200,000 metric tons as of the 2020s, with major producers in and scaling output in line with demand. A significant supply disruption occurred during 2008–2009, when reduced production capacity—driven by the global economic crisis lowering demand for downstream products like synthetic fibers—combined with plant maintenance issues and temporary shutdowns in key facilities, particularly in and the , led to a sharp acetonitrile shortage. This event caused prices to escalate dramatically from about $1.50 per kg to peaks of $25 per kg, impacting pharmaceutical and analytical sectors reliant on the solvent and prompting temporary shifts to alternatives or recycling methods.

Laboratory preparation

In laboratory settings, acetonitrile is commonly prepared by the dehydration of using strong dehydrating agents such as (P₂O₅) or (SOCl₂). The reaction involves heating with P₂O₅, which removes the elements of water from the amide group to form the : \ceCH3CONH2+P2O5−>CH3CN+HPO3+...\ce{CH3CONH2 + P2O5 -> CH3CN + HPO3 + ...} This method is straightforward and suitable for small-scale synthesis, producing acetonitrile in good yields after isolation. Another approach starts with the reaction of with to form as an intermediate, followed by dehydration to acetonitrile. The initial step is the nucleophilic acyl substitution where attacks the carbonyl carbon of , yielding and : \ceCH3COCl+NH3−>CH3CONH2+NH4Cl\ce{CH3COCl + NH3 -> CH3CONH2 + NH4Cl} The is then dehydrated using P₂O₅ or SOCl₂ as described above, ultimately affording acetonitrile via the loss of water. This two-step process allows for controlled preparation in educational or research laboratories. A less common laboratory method involves the catalytic reaction of with , typically promoted by metal catalysts such as cobalt-nickel, to produce acetonitrile, , and : \ceCH3CH2OH+NH3−>CH3CN+H2O+2H2\ce{CH3CH2OH + NH3 -> CH3CN + H2O + 2 H2} This dehydrogenative amination is more specialized for applications. Across these methods, yields typically range from 70% to 90%, depending on reaction conditions and scale. The crude product is purified by under reduced pressure ( approximately 82°C at , lower under ) to minimize and remove impurities like or unreacted starting materials.

Applications

Solvent uses

Acetonitrile is widely used as a in and , characterized by its high dielectric constant of 37.5 at 20°C, which facilitates the dissolution of many ionic salts while avoiding strong of anions through hydrogen bonding. This property enhances the nucleophilicity of anionic species, making it ideal for SN2 reactions, such as the displacement of alkyl halides by nucleophiles, where polar aprotic solvents like acetonitrile accelerate the reaction rate compared to protic alternatives. In (HPLC), acetonitrile serves as a key mobile phase component for reverse-phase separations, valued for its transparency in the UV region down to 190 nm, low that supports high flow rates, and compatibility with aqueous buffers. Its eluotropic strength and with water enable efficient for analyzing pharmaceuticals, biomolecules, and environmental samples. Acetonitrile functions as an entrainer in processes to separate azeotropic mixtures, such as benzene-cyclohexane, by selectively altering the relative volatilities of components through liquid-liquid equilibria interactions. This application leverages its polarity to preferentially solvate aromatic compounds, facilitating industrial-scale purification of hydrocarbon streams. In , acetonitrile is incorporated into non-aqueous electrolytes for -ion batteries, typically with lithium salts like LiPF₆, due to its low (0.34 cP), high ionic conductivity (up to 35 mS/cm in formulations), and wide electrochemical stability window exceeding 4 V. These attributes support fast-charging capabilities and operation across a broad range, as demonstrated in ternary solvent systems including carbonates. Acetonitrile also plays a role in , particularly in solid-phase methods, where its full miscibility with water aids in swelling and washing steps, while its aprotic nature promotes efficient coupling of hindered without side reactions. It serves as a greener alternative to DMF in Fmoc-based protocols, yielding high-purity peptides comparable to or exceeding traditional solvents.

Other industrial applications

Acetonitrile serves as a key extractive in the refining industry for the purification of 1,3- from crude C4 hydrocarbon streams obtained via . In this process, known as extractive , acetonitrile's high selectivity and polarity allow it to preferentially dissolve 1,3- while separating it from less polar components such as n-butane, , and butenes, enabling the production of high-purity butadiene for and other applications. This method is widely adopted in industrial plants, including those in , due to its efficiency and compatibility with existing infrastructure. Beyond extraction, acetonitrile functions as a chemical intermediate and in the synthesis of various pharmaceuticals and agrochemicals. It participates in reactions to form nitrogen-containing compounds, such as amides and heterocycles, which are building blocks for active pharmaceutical ingredients, including certain antidiabetic and antiviral drugs. For instance, acetonitrile can be incorporated into molecular frameworks through Ritter-type reactions or as a source in multi-component syntheses, contributing to the development of complex therapeutic agents. Its role as a reactant enhances reaction efficiency in electrochemical and catalytic processes tailored for pharmaceutical production. In , acetonitrile is employed as a in workflows, particularly for the desalting and enrichment of biomolecules. A common application involves acetonitrile precipitation to selectively remove high-molecular-weight proteins and salts from complex biological samples like serum, allowing for the isolation and of low-molecular-weight peptides and proteins with minimal interference. This technique improves efficiency and sensitivity in , facilitating proteomic studies and discovery. Acetonitrile is also utilized in the production of dyes and plastics, serving as a solvent and intermediate in the synthesis of various colorants and polymers. Additionally, it finds application in the manufacture of photographic films, where it aids in coating and steps; in perfumery as a solvent for extracting and formulating fragrances; in rubber product for and ; and in production as a reaction medium or intermediate. Furthermore, acetonitrile is employed in the separation of fatty acids from oils through processes like or extraction, enhancing the purification of oils and oleochemicals. In battery applications, acetonitrile acts as a co-solvent in electrolytes for lithium-ion batteries, enhancing ionic conductivity and enabling operation across a wide range from -40°C to 60°C. By blending with solvents and additives like fluoroethylene carbonate, these formulations promote faster lithium-ion transport and improved cycle life for electric vehicles and portable electronics.

Safety and toxicology

Health hazards

Acetonitrile exhibits moderate through oral, dermal, and inhalation routes, primarily due to its into (HCN) in the body. The oral LD50 in rats is 2,460 mg/kg, while the inhalation LC50 in rats is 17,100 ppm over 4 hours. These values indicate that significant exposure can lead to severe systemic effects, including , which arises from the metabolic release of rather than direct action of the parent compound. Recent incidents include a 2025 outbreak in from acetonitrile-adulterated alcoholic drinks, resulting in 20 hospitalizations and 2 fatalities. Symptoms of acute acetonitrile poisoning typically manifest with a delayed onset of 2 to 12 hours after exposure, attributed to the time required for metabolic conversion to . Initial signs include headache, nausea, dizziness, and vomiting, progressing to more severe manifestations such as (bluish skin discoloration due to oxygen deprivation), weakness, , , and . In severe cases, respiratory depression, convulsions, , and death may occur if cyanide levels become critically elevated. The toxicity mechanism involves hepatic metabolism of acetonitrile via enzymes, which oxidize it to cyanomethanol (HO-CH₂-CN), followed by further breakdown to ion (CN⁻) and (HCHO). This process, represented as CH₃CN → HO-CH₂-CN → CN⁻ + HCHO, occurs primarily in the liver and is oxygen- and NADPH-dependent. The released inhibits in the mitochondrial , disrupting and leading to and tissue hypoxia. Once formed, cyanide is detoxified by the enzyme rhodanese, which converts it to using as a sulfur donor; is then excreted primarily in the . The biological half-life of unmetabolized acetonitrile in humans is approximately 32 hours, while that of is about 15 hours, allowing for prolonged exposure to toxic metabolites in cases of significant or . Chronic exposure to acetonitrile may result in potential damage to the liver and kidneys, as well as effects such as numbness and tremors, based on animal studies and limited human data indicating accumulation in these organs. The International Agency for Research on Cancer (IARC) has not classified acetonitrile as to its carcinogenicity to humans due to insufficient data.

Environmental and handling considerations

Acetonitrile is biodegradable under aerobic conditions in and , serving as the primary degradation pathway in these media. In , the estimated aerobic is 2-8 weeks. It demonstrates ready biodegradability, achieving over 60% degradation within 10 days and 70% after 21 days in standard tests using non-adapted . Due to its log Kow value of -0.34, acetonitrile exhibits low potential and is not expected to concentrate in organisms. Acetonitrile is regulated as a hazardous air pollutant under the U.S. Clean Air Act by the Environmental Protection Agency. In the , it is registered under regulation, with detailed dossiers on its environmental and safety profiles. The sets a of 40 ppm as an 8-hour time-weighted average for workplace air. As a , acetonitrile has a of 2 °C (closed cup) and an of 524 °C, necessitating careful handling to avoid ignition sources. It should be used exclusively in well-ventilated areas such as chemical fume hoods to minimize exposure. Storage requires tightly sealed containers in a cool, dry place away from light, heat, and incompatible materials. For spill response, absorb the liquid with inert materials like vermiculite or dry sand, ensure adequate ventilation, and avoid confined spaces. Acetonitrile is incompatible with strong oxidizers and acids, which can lead to hazardous reactions. Disposal involves incineration at approved facilities or treatment as hazardous waste in accordance with local environmental regulations.

References

  1. https://pubchem.ncbi.nlm.nih.gov/compound/Acetonitrile
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  10. https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Organic_Chemistry_(Morsch_et_al.)/20%3A_Carboxylic_Acids_and_Nitriles/20.07%3A_Chemistry_of_Nitriles
  11. https://www.chemistrysteps.com/the-mechanism-of-nitrile-hydrolysis-to-carboxylic-acid/
  12. https://pure.mpg.de/rest/items/item_3009418/component/file_3010027/content
  13. https://organicchemistrydata.org/hansreich/resources/pka/
  14. https://pmc.ncbi.nlm.nih.gov/articles/PMC8659010/
  15. https://www.spectroscopyonline.com/view/organic-nitrogen-compounds-iv-nitriles
  16. https://www.chemicalbook.com/SpectrumEN_75-05-8_1hnmr.htm
  17. https://www.chemicalbook.com/SpectrumEN_75-05-8_13CNMR.htm
  18. https://www.sciencedirect.com/science/article/abs/pii/S0926860X15301630
  19. https://patents.google.com/patent/US4179462A/en
  20. https://www.sciencedirect.com/science/article/abs/pii/B978044453711950167X
  21. https://www.chemanalyst.com/industry-report/acetonitrile-market-721
  22. https://pure.qub.ac.uk/files/918400/PN%25202.pdf
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  24. https://pubchem.ncbi.nlm.nih.gov/compound/Acetamide
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