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Cyclopropane[1]
Cyclopropane - displayed formula
Cyclopropane - displayed formula
Cyclopropane - skeletal formula
Cyclopropane - skeletal formula
Names
Preferred IUPAC name
Cyclopropane[2]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.000.771 Edit this at Wikidata
KEGG
UNII
UN number 1027
  • InChI=1S/C3H6/c1-2-3-1/h1-3H2 checkY
    Key: LVZWSLJZHVFIQJ-UHFFFAOYSA-N checkY
  • InChI=1/C3H6/c1-2-3-1/h1-3H2
    Key: LVZWSLJZHVFIQJ-UHFFFAOYAL
  • C1CC1
Properties
C3H6
Molar mass 42.08 g/mol
Appearance Colorless gas
Odor Sweet, ethereal
Density 1.879 g/L (1 atm, 0 °C)
680 g/L (liquid)
Melting point −128 °C (−198 °F; 145 K)
Boiling point −32.9 °C (−27.2 °F; 240.2 K)
502 mg/L
Vapor pressure 640 kPa (20 °C)
1350 kPa (50 °C)
Acidity (pKa) ~46
−39.9·10−6 cm3/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Highly flammable
Asphyxiant
GHS labelling:
GHS02: Flammable
Danger
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 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 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard SA: Simple asphyxiant gas. E.g. nitrogen, helium
1
4
0
SA
495 °C (923 °F; 768 K)
Explosive limits 2.4 % (lower)
10.4 % (upper)
Safety data sheet (SDS) Air Liquide
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 ?)

Cyclopropane is the cycloalkane with the molecular formula (CH2)3, consisting of three methylene groups (CH2) linked to each other to form a triangular ring. The small size of the ring creates substantial ring strain in the structure. Cyclopropane itself is mainly of theoretical interest, but many cyclopropane derivatives are of commercial or biological significance.[3]

Cyclopropane was used as a clinical inhalational anesthetic from the 1930s through the 1980s. The substance's high flammability poses a risk of fire and explosions in operating rooms due to its tendency to accumulate in confined spaces, as its density is higher than that of air.

History

[edit]

Cyclopropane was discovered in 1881 by August Freund, who also proposed the correct structure for the substance in his first paper.[4] Freund treated 1,3-dibromopropane with sodium, causing an intramolecular Wurtz reaction leading directly to cyclopropane.[5] The yield of the reaction was improved by Gustavson in 1887 with the use of zinc instead of sodium.[6] Cyclopropane had no commercial application until Henderson and Lucas discovered its anaesthetic properties in 1929;[7] industrial production had begun by 1936.[8] In modern anaesthetic practice, it has been superseded by other agents.

Anaesthesia

[edit]

Cyclopropane was introduced into clinical use by the American anaesthetist Ralph Waters who used a closed system with carbon dioxide absorption to conserve this then-costly agent. Cyclopropane is a relatively potent, non-irritating and sweet smelling agent with a minimum alveolar concentration of 17.5%[9] and a blood/gas partition coefficient of 0.55. This meant induction of anaesthesia by inhalation of cyclopropane and oxygen was rapid and not unpleasant. However at the conclusion of prolonged anaesthesia patients could suffer a sudden decrease in blood pressure, potentially leading to cardiac dysrhythmia: a reaction known as "cyclopropane shock".[10] For this reason, as well as its high cost and its explosive nature,[11] it was latterly used only for the induction of anaesthesia, and has not been available for clinical use since the mid-1980s. Cylinders and flow meters were colored orange (now orange is used for the anesthetic gas enflurane).

Pharmacology

[edit]

Cyclopropane is inactive at the GABAA and glycine receptors, and instead acts as an NMDA receptor antagonist.[12][13] It also inhibits the AMPA receptor and nicotinic acetylcholine receptors, and activates certain K2P channels.[12][13][14]

Structure and bonding

[edit]
Further information: Bent bond
Orbital overlap in the bent bonding model of cyclopropane

The triangular structure of cyclopropane requires the bond angles between carbon-carbon covalent bonds to be 60°. The molecule has D3h molecular symmetry. The C-C distances are 151 pm versus 153-155 pm.[15][16]

Despite their shortness, the C-C bonds in cyclopropane are weakened by 34 kcal/mol vs ordinary C-C bonds. In addition to ring strain, the molecule also has torsional strain due to the eclipsed conformation of its hydrogen atoms. The C-H bonds in cyclopropane are stronger than ordinary C-H bonds as reflected by NMR coupling constants.

Bonding between the carbon centres is generally described in terms of bent bonds.[17] In this model the carbon-carbon bonds are bent outwards so that the inter-orbital angle is 104°.

The unusual structural properties of cyclopropane have spawned many theoretical discussions. One theory invokes σ-aromaticity: the stabilization afforded by delocalization of the six electrons of cyclopropane's three C-C σ bonds to explain why the strain of cyclopropane is "only" 27.6 kcal/mol as compared to cyclobutane (26.2 kcal/mol) with cyclohexane as reference with Estr=0 kcal/mol,[18][19][20] in contrast to the usual π aromaticity, that, for example, has a highly stabilizing effect in benzene. Other studies do not support the role of σ-aromaticity in cyclopropane and the existence of an induced ring current; such studies provide an alternative explanation for the energetic stabilization and abnormal magnetic behaviour of cyclopropane.[21]

Synthesis

[edit]

Cyclopropane was first produced via a Wurtz coupling, in which 1,3-dibromopropane was cyclised using sodium.[4] The yield of this reaction can be improved by the use of zinc as the dehalogenating agent and sodium iodide as a catalyst.[22]

BrCH2CH2CH2Br + 2 Na → (CH2)3 + 2 NaBr

Reactions

[edit]

Owing to the increased π-character of its C-C bonds, cyclopropane is often assumed to add bromine to give 1,3-dibromopropane, but this reaction proceeds poorly.[23] Hydrohalogenation with hydrohalic acids gives linear 1-halopropanes. Substituted cyclopropanes also react, following Markovnikov's rule.[24]

Electrophilic addition of HBr to cyclopropane

Cyclopropane and its derivatives can oxidatively add to transition metals, in a process referred to as C–C activation.

Safety

[edit]

Cyclopropane is highly flammable. However, despite its strain energy it does not exhibit explosive behavior substantially different from other alkanes.

See also

[edit]

References

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

Cyclopropane

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Cyclopropane is a three-membered cycloalkane with the molecular formula C₃H₆, consisting of three carbon atoms connected in a planar ring structure that results in significant angle strain due to bond angles of approximately 60°, far from the ideal 109.5° tetrahedral geometry.[1] This strain imparts unique reactivity to the molecule, making it more susceptible to ring-opening reactions compared to larger cycloalkanes.[1] As a colorless, odorless-to-petroleum-like gas at room temperature, it has a boiling point of -33 °C and a melting point of -128 °C, with a density of about 0.69 g/cm³ in liquid form.[1] Highly flammable and explosive when mixed with air or oxygen, cyclopropane requires careful handling due to its low ignition energy and wide explosive limits (2.4–10.4% in air).[2] First synthesized in 1882 by Austrian chemist August Freund through the reaction of 1,3-dibromopropane with sodium, cyclopropane remained a laboratory curiosity for decades until its anesthetic properties were accidentally discovered in 1928 by Canadian researchers Velyien E. Henderson and George H. W. Lucas at the University of Toronto, who noted its potent effects while investigating contaminants in propylene samples.[3] The compound was introduced to clinical anesthesia in 1930 by Ralph M. Waters at the University of Wisconsin, marking the beginning of its widespread use as an inhalational agent valued for rapid induction, minimal physiological disturbance, and suitability for pediatric and obstetric cases.[3] By the 1930s and 1940s, it became one of the most popular general anesthetics, supplied in portable cylinders by companies like E. R. Squibb & Sons, and was administered in over 7,000 documented cases at institutions like Rochester General Hospital between 1935 and 1940 with low complication rates when properly managed.[3] However, its extreme flammability led to numerous operating room explosions, prompting its gradual replacement by non-flammable alternatives like halothane starting in the late 1950s, after which production ceased for medical purposes.[3] Beyond its historical role in anesthesia, cyclopropane serves as a key building block in organic synthesis, particularly for constructing strained ring systems in pharmaceuticals and agrochemicals, where the cyclopropane motif enhances potency, metabolic stability, and selectivity in drugs such as antibiotics and insecticides.[1] Its reactivity facilitates innovative synthetic methods, including recent palladium-catalyzed approaches to form cyclopropane rings in complex molecules, underscoring ongoing research into its applications despite limited commercial availability today. Safety protocols emphasize its hazards as a compressed, asphyxiant gas, incompatible with oxidizers, and it is now primarily handled in research settings under strict ventilation and ignition controls.[1]

Structure and properties

Bonding and ring strain

Cyclopropane possesses the molecular formula C3H6C_3H_6 and adopts a planar, triangular ring structure in which the three carbon atoms are connected by single bonds, forcing the C-C-C bond angles to 60°—a significant deviation from the ideal tetrahedral geometry of 109.5° expected for sp3sp^3-hybridized carbons.[4] This compressed geometry arises from the constraints of forming a stable three-membered ring, resulting in a highly symmetric molecule with D3hD_{3h} point group symmetry, where all carbon-hydrogen bonds are equivalent. The unusual bonding in cyclopropane is best explained by the bent bond theory, originally proposed by Coulson and Moffitt in 1949, which describes the C-C bonds as curved, "banana-shaped" orbitals rather than straight sigma bonds aligned along the internuclear axis. In this model, the hybrid orbitals on each carbon form an inter-orbital angle of approximately 104°, allowing better overlap despite the acute ring angle and leading to a C-C bond length of 1.51 Å—shorter than the typical unstrained sp3sp^3 C-C bond of 1.54 Å but indicative of the strain effects. This bending enhances orbital overlap in the plane perpendicular to the ring but weakens the bonds overall, contributing to the molecule's reactivity. The total ring strain energy in cyclopropane amounts to 27.6 kcal/mol, as determined from experimental heats of combustion and computational analyses, and is partitioned into angle strain from the distorted bond angles and torsional strain from the eclipsed conformation of the adjacent C-H bonds.[5] Angle strain alone accounts for much of this destabilization, as the 60° angles impose significant compression on the sp3sp^3 hybrids, while torsional strain further elevates the energy due to the inability to achieve staggered conformations. Compared to larger cycloalkanes like cyclohexane, which minimize strain through chair conformations, cyclopropane's rigidity exemplifies how small rings amplify these energetic penalties. Quantum mechanical descriptions reveal that the carbon hybridization in cyclopropane lies between sp3sp^3 and sp2sp^2, with greater p-character in the C-C bonding orbitals and increased s-character in the C-H bonds, resembling ethylene more closely than ethane. This partial sp2sp^2-like hybridization facilitates a degree of delocalization in the sigma framework, often termed σ-aromaticity, first proposed by Dewar in 1979 to account for the unexpectedly modest strain relative to predictions from pure angle distortion. The three parallel Walsh orbitals above and below the ring plane contribute to this stabilizing cyclic conjugation of the six sigma electrons, providing a conceptual bridge to π-aromatic systems like benzene. Spectroscopic techniques confirm these structural and bonding features. Infrared spectroscopy shows a characteristic C-C symmetric stretching mode at approximately 1020 cm⁻¹, lower than typical alkane C-C stretches due to the weakened, bent bonds. In ¹H NMR, the six hydrogens are magnetically equivalent by symmetry, appearing as a singlet at δ 0.22 ppm, shifted upfield from alkane protons owing to the ring current effects and high s-character in C-H bonds. These observations underscore the unique electronic environment imposed by the strained ring.

Physical characteristics

Cyclopropane is a colorless gas at standard conditions, characterized by a sweet, ethereal odor that is mildly pungent. This appearance and sensory property make it distinct among small hydrocarbons.[6] Key physical data for cyclopropane include a boiling point of -32.8 °C at 760 mmHg and a melting point of -127.4 °C. The density of the gas is 1.879 g/L at 0 °C and 1 atm, reflecting its lightweight nature compared to air. The critical temperature is 124.7 °C, above which it cannot be liquefied regardless of pressure. These values indicate a compound that remains gaseous at ambient temperatures, with the low boiling point partly attributable to ring strain effects.[1][1][1] Cyclopropane shows limited solubility in water, approximately 0.038 g/100 mL at 35 °C, consistent with its nonpolar character. In contrast, it is highly soluble in organic solvents, such as ethanol where solubility exceeds 1 g/100 mL, and ether. This solubility profile facilitates its handling in non-aqueous environments.[1][1] Thermodynamic properties include a standard enthalpy of formation of +53.3 kJ/mol for the gas phase, highlighting the energetic cost of its strained structure. The standard heat of combustion is -2091 kJ/mol, providing a measure of its energy content. Vapor pressure is notably high at 5410 mm Hg (722 kPa) at 25 °C, underscoring its volatility. The phase diagram features a triple point at approximately -127.6 °C and 0.00078 bar, with a critical pressure of 54.2 atm, delineating a compact liquid phase region under moderate compression and cooling.[7][7][1]

Synthesis

Classical methods

The initial laboratory synthesis of cyclopropane was achieved by August Freund in 1881 through an intramolecular variant of the Wurtz coupling reaction. Freund treated 1,3-dibromopropane with sodium metal in dry ether, generating the three-membered ring via double dehalogenation and carbon-carbon bond formation, with reported yields of approximately 20-30%. This method marked the first preparation of the compound, confirming its cyclic trimethylene structure, though it suffered from modest efficiency due to competing elimination reactions. An improvement on Freund's approach was introduced by Gustav Gustavson in 1887, replacing sodium with zinc dust to enhance selectivity and yield. The reaction involves 1,3-dibromopropane (or analogous dihalides) with activated zinc in a solvent like ethanol or ether, typically at elevated temperatures, producing cyclopropane in yields of 50-60%. Further refinement came in the 1930s with the addition of sodium iodide as a catalyst, which facilitates halogen exchange to form more reactive iodides in situ, boosting yields to over 90% in optimized conditions. This zinc-mediated reduction became the standard classical route, emphasizing the need for anhydrous conditions to minimize side reactions. By the 1930s, the zinc-based method was scaled for industrial production to meet demand for cyclopropane as an anesthetic gas. However, these classical methods were hampered by inherently low to moderate yields in unoptimized runs, formation of side products such as propene via elimination, and the requirement for high-purity 1,3-dihalopropane precursors, which were costly to prepare from allyl halides or glycerol derivatives. These approaches remain relevant for the synthesis of unsubstituted cyclopropane.

Modern advancements

Since the mid-20th century, the Simmons–Smith reaction has emerged as a cornerstone for stereospecific cyclopropanation, involving the treatment of alkenes with diiodomethane and a zinc–copper couple to effect methylene transfer, yielding cyclopropanes with high fidelity in stereochemistry preservation.[8] This method, adaptable for synthesizing cyclopropane and its simple derivatives through intermediates like propene-derived precursors, routinely achieves yields exceeding 80% under optimized conditions, making it suitable for scalable production of unsubstituted or monosubstituted cyclopropanes.[9] Its chemoselectivity toward alkenes minimizes side reactions, even with functionalized substrates, positioning it as a reliable post-1950 advancement over earlier less efficient routes.[10] Complementing this, carbene insertion methods utilizing diazomethane decomposition have gained prominence for [2+1] cycloadditions to alkenes, particularly when catalyzed by palladium or copper complexes, enabling efficient construction of cyclopropane rings with control over regioselectivity.[11] These catalytic systems, often employing zerovalent Pd or Cu species, facilitate the reaction under mild conditions, with palladium catalysts demonstrating superior efficiency for a broad range of unsaturated substrates, including those leading to simple cyclopropane derivatives.[12] Yields in these processes typically range from moderate to high, depending on the alkene substitution, and the approach has been refined for stereocontrol in derivative synthesis since the 1960s.[13] Advancements from 2020 to 2025 have further enhanced efficiency and sustainability, exemplified by organoelectrocatalytic cyclopropanation of alkenyl trifluoroborates with methylene compounds, which proceeds intermolecularly under metal-free electrochemical conditions to deliver functionalized cyclopropanes with broad substrate tolerance.[14] This method, reported in late 2024, achieves high yields and enantioselectivities for simple derivatives, addressing limitations in traditional carbenoid transfers by leveraging anodic oxidation for carbene generation.[15] Similarly, Oxone®/KI-promoted protocols have enabled Michael-initiated ring closure (MIRC) for spirocyclopropane synthesis from α,β-unsaturated carbonyls and sulfur ylides, offering a mild, oxidant-driven route with yields up to 90% for strained spiro systems relevant to cyclopropane motifs.[16] In pharmaceutical contexts, modular routes employing pinacol boronate intermediates have streamlined access to cyclopropyl amino acids, as demonstrated by ruthenium-catalyzed enantioselective cyclopropanation followed by borylation, providing scalable building blocks with >95% ee and gram-scale throughput.[17] Industrial scalability has benefited from continuous flow processes, which mitigate safety risks associated with batch cyclopropanations in reactions like nickel-catalyzed electroreductive variants using gem-dichloroalkanes.[18] These flow systems, optimized as of 2025, support larger-scale production of cyclopropane intermediates for agrochemical and pharmaceutical applications, with electrochemical adaptations enabling sustainable operation under ambient conditions.[19] Analyses of 2025 manufacturing highlight these innovations as contributing to economic viability in the sector.[20]

Reactions and reactivity

Ring-opening reactions

Cyclopropane's ring strain facilitates electrophilic ring-opening reactions, where the C-C bonds break to afford linear products, often analogous to alkene additions but driven by the relief of approximately 28 kcal/mol of strain energy. Hydrohalogenation of cyclopropane with HX (where X = Cl or Br) proceeds to give 1-halopropanes, such as n-propyl chloride or n-propyl bromide, in a regioselective manner that favors the primary halide due to the stability of the resulting linear structure and strain relief. For instance, the reaction with HBr yields 1-bromopropane as the major product under acidic conditions. This addition follows Markovnikov orientation, with the proton adding to one carbon and the halide to the adjacent position.[21] The mechanism involves initial electrophilic protonation of the cyclopropane ring by H⁺ from HX, forming a corner-protonated cyclopropane species as an intermediate or transition state. This unsymmetrical structure features partial positive charge on the adjacent carbons due to the bent bonds, leading to rapid C-C bond cleavage and generation of a primary carbocation (CH₃CH₂CH₂⁺). The halide ion (X⁻) then attacks the carbocation at the terminal carbon in an Sₙ2-like fashion, completing the ring opening. The transition state for protonation exhibits asynchronous bond breaking, with the C-C bond elongation and partial carbocation development, as supported by stereochemical studies showing inversion at the nucleophilic attack site.[22] The overall reaction is exemplified by:
CX3HX6+HBr→CHX3CHX2CHX2Br \ce{C3H6 + HBr -> CH3CH2CH2Br}
Hydrogenolysis represents another key ring-opening pathway, where cyclopropane is catalytically reduced to propane using palladium on carbon (Pd/C) under mild conditions, typically at approximately 175 °C and atmospheric pressure, achieving quantitative conversion. The mechanism entails adsorption of the cyclopropane and H₂ on the Pd surface, followed by strain-assisted C-C bond cleavage and stepwise hydrogenation, often involving surface-bound intermediates that allow for hydrogen-deuterium exchange when D₂ is employed. This process highlights the role of metal catalysts in promoting selective bond scission without skeletal rearrangement.[23] Electrophilic additions to unsubstituted cyclopropane are generally inefficient compared to alkenes. For example, Br₂ exhibits poor reactivity under standard conditions (room temperature, dark), with no significant ring-opening addition to form 1,3-dibromopropane observed, due to the insufficient π-character of the bent C-C bonds for effective bromonium ion formation. Instead, free-radical pathways dominate under UV irradiation, leading to substitution or partial addition products. Thermal pyrolysis of cyclopropane, however, induces clean ring opening to propene via a biradical mechanism at temperatures above 400 °C, involving initial C-C bond homolysis to a 1,3-diradical intermediate, followed by 1,2-hydrogen migration; this unimolecular process has an activation energy of about 65 kcal/mol and follows first-order kinetics.[24]

Catalytic transformations

Cyclopropane serves as a valuable synthon in organic synthesis due to its ability to undergo catalytic transformations that functionalize the ring while preserving its strained structure, allowing for the construction of complex molecules with precise stereocontrol. Transition metal catalysts, particularly those involving iridium and rhodium, enable selective C-H activation at the methylene positions of cyclopropane, facilitating deuteration or arylation without ring disruption. For instance, iridium complexes with chiral bidentate boryl ligands promote enantioselective C(sp³)-H borylation, targeting the less substituted methylene carbons to yield borylated cyclopropanes in up to 98% enantiomeric excess.[25] Similarly, rhodium catalysts achieve enantioselective silylation of cyclopropyl C-H bonds, providing hydrido(silyl) ethers as versatile intermediates for further derivatization. These activations leverage the ring strain to lower the energy barrier for C-H cleavage, enabling site-selective modifications that are challenging in unstrained alkanes. Cross-coupling reactions further highlight cyclopropane's utility, with nickel catalysts enabling the formation of cyclopropyl arenes from aryl halides while maintaining ring integrity. A notable example is the reductive cross-coupling of cyclopropylamine NHP esters with (hetero)aryl halides, catalyzed by nickel under mild conditions, which proceeds rapidly in less than 2 hours and tolerates a broad range of functional groups to afford 1-arylcyclopropylamines in good yields.[26] This method avoids air- or heat-sensitive reagents and directly accesses motifs prevalent in pharmaceuticals, demonstrating cyclopropane's role in building diversity through ring-intact C-C bond formation. Recent advancements in asymmetric catalysis from 2020 to 2025 have expanded access to stereodefined cyclopropane derivatives. Enantioselective [2+1] cycloadditions, such as the sulfoximine-mediated Johnson-Corey-Chaykovsky reaction applied to menthyl acrylates, provide trans-cyclopropyl esters with high diastereoselectivity (up to 9:1 dr) and scalability for lead optimization in drug discovery.[27] Additionally, catalytic methods for alkylidenecyclopropane synthesis via strain-relieving prototropic shifts using bifunctional iminophosphorane catalysts achieve up to 99% ee, enabling the preparation of enantioenriched precursors for bioactive compounds like pyrethroid insecticides.[28] These developments underscore the growing synthetic versatility of cyclopropane in enantioselective transformations. A representative cross-coupling example is depicted below, where a cyclopropylamine NHP ester reacts with an aryl halide (Ar-X) under nickel catalysis to form 1-arylcyclopropylamine:
(c-C3H4NH2-NHP ester)+Ar-X→[Ni],reductantAr-c-C3H4NH2 \text{(c-C$_3$H$_4$NH$_2$-NHP ester)} + \text{Ar-X} \xrightarrow{[\text{Ni}], \text{reductant}} \text{Ar-c-C$_3$H$_4$NH$_2$}
This reaction exemplifies the precision of modern catalysis in appending aryl groups to the cyclopropane ring.[26]

Applications

Anesthetic applications

Cyclopropane was first recognized for its anesthetic potential in 1929 by pharmacologists Velyien E. Henderson and George H. W. Lucas at the University of Toronto, who demonstrated its ability to produce surgical anesthesia in experimental animals without significant toxicity.[29] This discovery stemmed from observations during studies on propylene, where cyclopropane emerged as the active component responsible for the observed effects. Clinical adoption followed in the early 1930s under the leadership of anesthetist Ralph M. Waters at the University of Wisconsin, who pioneered its use in human patients through a closed-circuit delivery system that minimized waste and enhanced safety.[3] Waters' work established cyclopropane as a viable alternative to ether and chloroform, leading to widespread use in surgical settings until the mid-20th century. The pharmacological profile of cyclopropane as an inhalational anesthetic involves modulation of neuronal ion channels and receptors, contributing to its rapid onset of unconsciousness and immobility. It functions as an antagonist at NMDA receptors, inhibiting excitatory glutamatergic transmission, while also suppressing AMPA and nicotinic acetylcholine receptors; concurrently, it activates two-pore domain potassium (K2P) channels, promoting hyperpolarization and neuronal inhibition.[30] The minimum alveolar concentration (MAC) required to prevent movement in 50% of patients is 9.2 vol%, reflecting its moderate potency compared to other agents.[31] Its low blood/gas partition coefficient of 0.46 facilitates swift induction and emergence by allowing quick alveolar uptake and minimal accumulation in blood.[31] Cyclopropane undergoes negligible hepatic metabolism, with over 99% eliminated unchanged via the lungs, reducing risks of toxic byproducts. Key advantages of cyclopropane included its high potency, non-irritating nature to airways, and pleasant sweet odor, which improved patient tolerance during mask induction.[32] It provided stable cardiovascular and respiratory function at clinical doses, with minimal postoperative nausea or cognitive impairment, making it suitable for a range of procedures including obstetrics and pediatrics. However, drawbacks limited its long-term viability; notable among these was "cyclopropane shock," a sudden postoperative hypotension attributed to sympathetic blockade and vasodilation upon abrupt discontinuation.[33] By the late 1950s, cyclopropane was largely supplanted by safer, non-flammable alternatives such as halothane, though its historical role influenced modern anesthetic practices emphasizing rapid recovery. Administration typically involved mixtures of 20-50% cyclopropane in oxygen or oxygen-nitrous oxide carriers, delivered via semi-closed or closed-circuit systems to achieve 1.0-1.5 MAC for maintenance after initial induction at higher concentrations.[34] This dosing regimen capitalized on its volatility and low solubility for controlled depth of anesthesia, often without supplemental agents for short surgeries.

Synthetic and industrial uses

Cyclopropane serves as a key building block for incorporating cyclopropyl motifs into pharmaceutical compounds, enhancing metabolic stability, potency, and selectivity due to the ring's unique three-dimensional structure. For instance, Nirmatrelvir, a component of the COVID-19 treatment Paxlovid, features a bicyclic gem-dimethyl cyclopropyl proline moiety that contributes to its efficacy as a SARS-CoV-2 main protease inhibitor. Recent research from 2020 to 2025 has highlighted the growing prevalence of fused-cyclopropane rings in drug discovery, where they provide structural novelty and improved pharmacokinetic properties, such as in inhibitors for cancer (e.g., Akt), Alzheimer's (BACE-1), and viral diseases. These motifs appear in clinical candidates like NTQ1062 and PF-00835231, demonstrating enhanced target binding and permeability compared to non-fused analogs.[35][36][37] In natural product synthesis, cyclopropane-containing compounds, particularly those derived from microbial sources like lipids, have been targeted through advanced total synthesis strategies to elucidate their biological roles. Recent efforts include the synthesis of microbe-derived cyclopropane lipids that activate host nuclear receptors, aiding in understanding host-pathogen interactions and immune modulation. These syntheses leverage novel methodologies, such as stereoselective cyclopropanation, to construct complex structures found in bacterial lipids, as reviewed in progress from 2016 to 2024 with applications extending into 2025.[38][39] Industrial production of cyclopropane has seen advancements in scalable manufacturing processes to meet demand in specialty chemicals. Market reports indicate its use as an intermediate in agrochemicals, where cyclopropane derivatives like cyclopropanecarboxylic acid enhance the efficacy of pesticides and herbicides through improved bioactivity and stability. Additionally, it functions as a precursor for polymer materials, contributing to the development of advanced resins and coatings via ring-opening polymerization pathways. Recent plant setups emphasize sustainable production methods to support these applications.[20][40][41] Emerging research explores cyclopropane derivatives, such as fatty acids, for applications in food authentication, particularly in dairy products where they serve as biomarkers for production practices like silage feeding. Studies in 2025 have differentiated specific cyclopropane fatty acids (e.g., dihydrosterculic and lactobacillic acids) in milk using GC-MS and NMR, enabling traceability for premium cheeses like Parmigiano Reggiano. While synthetic methods, including potential electrocatalytic approaches, are under investigation to produce these markers for analytical standards, their microbial origins underscore ongoing synthetic challenges.[42][43]

History

Discovery and early research

Cyclopropane was first synthesized in 1881 by the Austrian chemist August Freund while working at the University of Lemberg, through the reaction of 1,3-dibromopropane with sodium metal. This classical method involved treating the dihalide with sodium to form the three-membered ring, marking the initial preparation of the smallest stable cycloalkane. Freund reported the discovery and proposed the correct cyclic structure in his key publication the following year.[44] The structure of cyclopropane as trimethylene (the original name) was confirmed in 1885 by Adolf von Baeyer, who integrated it into his strain theory to explain the compound's reactivity and physical properties arising from bond angle deviation in small rings. Baeyer's theoretical framework highlighted the angular strain in cyclopropane's 60° bond angles, distinguishing it from larger cycloalkanes and open-chain hydrocarbons like propylene. This confirmation solidified the compound's identity as C₃H₆ with a triangular carbon skeleton. Early 20th-century characterization efforts further validated the molecular formula C₃H₆ through combustion and vapor density studies. Researchers measured the gas's density relative to air and analyzed combustion products—yielding carbon dioxide and water in ratios consistent with three carbon and six hydrogen atoms—providing empirical confirmation of the composition. These investigations, building on Freund's work, established cyclopropane's basic thermodynamic properties and purity. In the 1910s, pre-anesthetic research examined cyclopropane's flammability and structural integrity, with studies revealing its wide explosive limits in air and sensitivity to ignition. Chemists such as W. A. Bone and R. V. Wheeler investigated slow combustion processes, identifying intermediate products like formaldehyde and highlighting the ring's tendency to open under oxidative conditions, which informed early safety considerations for handling the gas.

Medical development and obsolescence

The anesthetic properties of cyclopropane were first identified in 1929 through experiments conducted by pharmacologists Velyien E. Henderson and George H. W. Lucas at the University of Toronto, who tested the gas on laboratory animals and reported its rapid onset of anesthesia with minimal physiological disruption compared to existing agents like ether.[45][46] These findings marked a breakthrough in inhalational anesthesia, as cyclopropane required lower concentrations for effective narcosis and allowed quicker recovery, prompting further clinical evaluation.[3] In 1930, Ralph M. Waters at the University of Wisconsin administered cyclopropane to humans for the first time during surgical procedures, confirming its efficacy and safety in controlled settings, which accelerated its adoption in medical practice.[3] By 1936, industrial-scale production had commenced, enabling widespread availability and integration into hospital anesthesia protocols, with manufacturers like E.R. Squibb scaling up synthesis from 1,3-dichloropropane via zinc reduction to meet demand.[47] This commercialization transformed cyclopropane into the dominant inhalational anesthetic of the era, supplanting ether in many operations due to its potency and reduced side effects.[3] During World War II, cyclopropane proved essential for battlefield surgery, as its gaseous form permitted easy transport in lightweight cylinders and portable apparatus, facilitating rapid anesthesia administration in forward medical units under austere conditions.[48] U.S. Army consultants emphasized its role in enabling efficient thoracic and abdominal procedures amid high casualty volumes, though concerns over explosion risks in oxygen-rich environments led to partial restrictions later in the conflict.[49] Its portability and hemodynamic stability made it indispensable for mobile surgical teams, contributing to lower perioperative mortality rates in combat zones.[50] The decline of cyclopropane began in the 1950s with the introduction of non-flammable alternatives like halothane in 1956, which offered similar potency without the explosion hazards that had caused fatal incidents in operating rooms.[51] By the 1960s and 1970s, safer halogenated agents such as enflurane and isoflurane further eroded its use, as hospitals prioritized risk mitigation and regulatory pressures mounted against flammable gases.[3] By the late 1970s, its clinical use had largely ceased in the U.S. due to these safety concerns and the superiority of modern anesthetics.[34] Despite its obsolescence, cyclopropane's legacy endures in the evolution of inhalational anesthesia techniques, having pioneered closed-circuit delivery systems and low-flow administration methods that remain standard today.[3] Its unique pharmacological profile—minimal metabolism and rapid equilibration—continues to inform studies on anesthetic mechanisms, with research into its effects on ion channels and neuroprotection extending into the 2000s. As of 2025, cyclopropane is no longer used clinically but remains of interest in research for its unique properties.[32]

Safety and handling

Health and toxicity risks

Cyclopropane acts as a simple asphyxiant, displacing oxygen in enclosed spaces and potentially leading to hypoxia even at concentrations below those causing direct toxic effects.[52] Acute exposure to elevated concentrations produces central nervous system (CNS) depression, manifesting as headache, dizziness, nausea, lightheadedness, and loss of coordination. Narcosis and anesthetic effects occur at concentrations typically above 5% based on historical use, potentially progressing to unconsciousness. Higher levels may induce convulsions, cardiac arrhythmias, respiratory distress, coma, and death, primarily through profound CNS and cardiovascular suppression. Unlike many volatile agents, cyclopropane exhibits minimal hepatotoxicity or nephrotoxicity, with adverse effects largely confined to the nervous and respiratory systems. No evidence of carcinogenicity has been established, and chronic toxicity data are limited, with no identified long-term organ damage from repeated low-level exposure.[53][1] These toxic effects stem from mechanisms akin to its historical anesthetic role, involving general CNS depression without specific antagonism at NMDA receptors, though it shares broad inhibitory actions on neuronal excitability common to inhalational agents. Occupational exposure limits have not been formally established by OSHA (no PEL) or ACGIH (no TLV); however, Protective Action Criteria (PAC) include PAC-1 at 600 ppm, PAC-2 at 4,000 ppm, and PAC-3 at 6,000 ppm for emergency response planning.[52][54] Industrial incidents involving cyclopropane leaks have reported dizziness and unconsciousness among workers, as seen in a fatal case of a 22-year-old male found deceased in a storage room due to intentional high-concentration inhalation, highlighting risks of rapid narcosis and asphyxiation in poorly ventilated areas. Such cases underscore the need for oxygen monitoring and ventilation in handling environments to prevent non-therapeutic exposures.[55]

Flammability hazards

Cyclopropane is a highly flammable gas that poses significant fire and explosion risks due to its low ignition energy and wide flammable range in air.[1] Its autoignition temperature is 500°C, and the flash point for the liquefied form is approximately -95°C, making it susceptible to ignition from common heat sources.[56] The lower explosive limit (LEL) is 2.4 vol% and the upper explosive limit (UEL) is 10.4 vol% in air, with a minimum ignition energy of 0.17 mJ at 6.3% concentration, which is low enough for static sparks or electrical arcs to initiate combustion.[54][57] Upon ignition, cyclopropane undergoes an exothermic combustion reaction, primarily producing carbon dioxide (CO₂) and water (H₂O), with a heat of combustion of 49.7 MJ/kg that releases substantial energy and can intensify fires rapidly.[58] The National Fire Protection Association (NFPA) rates cyclopropane with a flammability hazard of 4 (severe), health hazard of 1 (slight), and instability of 0 (minimal), emphasizing its extreme fire danger while noting relative chemical stability under normal conditions.[1] Safe handling requires storage in high-pressure cylinders to maintain it as a liquid, away from ignition sources, and the use of explosion-proof electrical equipment to prevent sparks in classified hazardous areas (Class I, Group C).[59] Modern safety data sheets recommend grounding and bonding containers during transfer, using non-sparking tools, and ensuring adequate ventilation to avoid accumulation of flammable vapors heavier than air.[60] For firefighting, shutting off the gas supply is critical, supplemented by water spray or foam to cool surrounding areas and prevent explosion propagation. Historical incidents include laboratory explosions attributed to static sparks igniting cyclopropane vapors, highlighting the need for rigorous grounding protocols; similar risks were evident in early medical applications where uncontrolled leaks led to over 300 operating room fires and explosions between 1930 and 1960.[34] Current guidelines from safety data sheets, such as Air Liquide's 2020 edition, stress prohibiting smoking, open flames, and hot surfaces near storage or use areas to mitigate these hazards.[60]

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