Inorganic chemistry,ring strain is a type of instability that exists whenbonds in amolecule formangles that deviate from their normal values as a result of being part of aring. This type ofstrain is most commonly discussed for small rings such ascyclopropanes andcyclobutanes, whose internal angles are substantially smaller than the idealized value of approximately 109°. Because of their high strain, theheat of combustion for these small rings is elevated.^[1][2]

Ring strain results from a combination ofangle strain,conformational strain orPitzer strain (torsional eclipsing interactions), andtransannular strain, also known asvan der Waals strain orPrelog strain. The simplest examples of angle strain are small cycloalkanes such as cyclopropane and cyclobutane.

Ring strain energy can be attributed to the energy required for the distortion of bond and bond angles in order to close a ring.^[3]

The avoidance or reduction of ring strain can help direct or accelerate chemical reactions.

Ring strain theory was first developed by German chemist Adolf von Baeyer in 1890. Previously, the only types of strain believed to exist were torsional and steric; however, Baeyer's theory became based on the interactions between the two strains. Baeyer's theory was based on the assumption that the rings in cyclic compounds were flat, and the simple geometry required for atoms in a planar structure.

Around the same time,Hermann Sachse [de] postulated that rings were not flat, and potentially existed in other folded or twistedconformations. Ernst Mohr later combined the two theories to explain the stability of six-membered rings and their frequency in nature, as well as the energy levels of other ring structures.^[4]

In alkanes, optimum overlap ofatomic orbitals is achieved at 109.5°, the mathematically ideal angle fortetrahedral geometry. The most common cyclic compounds have five or six carbons in their ring.^[5]Adolf von Baeyer received aNobel Prize in 1905 for the discovery of the Baeyer strain theory, which was an explanation of the relative stabilities of cyclic molecules in 1885.^[5]

Angle strain occurs whenbond angles deviate from the ideal bond angles to achieve maximum bond strength in a specificchemical conformation. Angle strain typically affects cyclic molecules, which lack the flexibility of acyclic molecules.

Angle strain destabilizes a molecule, as manifested in higher reactivity and elevatedheat of combustion. Maximum bond strength results from effective overlap of atomic orbitals in achemical bond. A quantitative measure for angle strain isstrain energy. Angle strain andtorsional strain combine to create ring strain that affects cyclic molecules.^[5]

${\displaystyle {\text{C}}_n{\text{H}}_{2n}+\frac{3n}{2}{\text{O}}_2^{}⟶n{\text{CO}}_2^{}+n{\text{H}}_2^{}\text{O}-\mathrm{\Delta }{H}_{\text{combustion}}}$

Normalized energies that allow comparison of ring strains are obtained by measuring permethylene group (CH₂) of the molar heat of combustion in the cycloalkanes.^[5]

ΔH_combustion per CH₂ − 658.6 kJ = strain per CH₂

The value 658.6 kJ per mole is obtained from an unstrained long-chain alkane.^[5]

Cycloalkanes generally have less ring strain than cycloalkenes, which is seen when comparing cyclopropane and cyclopropene.^[7]

This sectionmay beconfusing or unclear to readers. In particular, This section conflates angle deviation (now 120° vs previously discussed 109.5°) with forced non-planarity or twisted alkenes (conformational and/or improper-dihedral issues).. Please helpclarify the section. There might be a discussion about this onthe talk page.(September 2025) (Learn how and when to remove this message)

Cyclic alkenes are subject to strain resulting from distortion of the sp²-hybridized carbon centers. Illustrative isC₆₀ where the carbon centres are pyramidalized. This distortion enhances the reactivity of this molecule. Angle strain also is the basis ofBredt's rule which dictates that bridgehead carbon centers are not incorporated in alkenes because the resulting alkene would be subject to extreme angle strain.

Small trans-cycloalkenes have so much ring strain they cannot exist for extended periods of time.^[8] For instance, the smallest trans-cycloalkane that has been isolated istrans-cyclooctene. Trans-cycloheptene has been detected viaspectrophotometry for minute time periods, and trans-cyclohexene is thought to be an intermediate in some reactions. No smaller trans-cycloalkenes are known. On the contrary, while small cis-cycloalkenes do have ring strain, they have much less ring strain than small trans-cycloalkenes.^[8]

In general, the increased levels of unsaturation in alkenes leads to higher ring strain. Increasing unsaturation leads to greater ring strain in cyclopropene.^[7] Therefore, cyclopropene is an alkene that has the most ring strain between the two mentioned. The differing hybridizations and geometries between cyclopropene and cyclopropane contribute to the increased ring strain. Cyclopropene also has an increased angle strain, which also contributes to the greater ring strain. However, this trend does not always work for every alkane and alkene.^[7]

In some molecules, torsional strain can contribute to ring strain in addition to angle strain. One example of such a molecule iscyclopropane. Cyclopropane's carbon-carbon bonds form angles of 60°, far from the preferred angle of 109.5° angle in alkanes, so angle strain contributes most to cyclopropane's ring strain.^[9] However, as shown in theNewman projection of the molecule, the hydrogen atoms are eclipsed, causing some torsional strain as well.^[9]

In cycloalkanes, each carbon is bondednonpolarcovalently to two carbons and two hydrogen. The carbons havesp³ hybridization and should have ideal bond angles of 109.5°. Due to the limitations of cyclic structure, however, the ideal angle is only achieved in a six carbon ring —cyclohexane inchair conformation. For other cycloalkanes, the bond angles deviate from ideal.

Molecules with a high amount of ring strain consist of three, four, and some five-membered rings, including:cyclopropanes,cyclopropenes,cyclobutanes,cyclobutenes, [1,1,1]propellanes, [2,2,2]propellanes,epoxides,aziridines,cyclopentenes, andnorbornenes. These molecules havebond angles between ring atoms which are more acute than the optimal tetrahedral (109.5°) and trigonal planar (120°)bond angles required by their respective sp³ and sp² bonds. Because of the smallerbond angles, the bonds have higher energy and adopt more p-character to reduce the energy of the bonds. In addition, the ring structures of cyclopropanes/enes and cyclclobutanes/enes offer very little conformational flexibility. Thus, the substituents of ring atoms exist in aneclipsed conformation in cyclopropanes and between gauche and eclipsed in cyclobutanes, contributing to higher ring strain energy in the form of van der Waals repulsion.

monocycles

• cyclopropane (29 kcal/mol), C₃H₆ — the C-C-C bond angles are 60° whereas tetrahedral 109.5° bond angles are expected.^[6] The intense angle strain leads to nonlinear orbital overlap of its sp³ orbitals.^[5] Because of the bond's instability, cyclopropane is more reactive than other alkanes.^[5] Since any three points make a plane and cyclopropane has only three carbons, cyclopropane is planar.^[6] The H-C-H bond angle is 115° whereas 106° is expected as in the CH₂ groups of propane.^[6]
• cyclobutane (26.3 kcal/mol), C₄H₈ — if cyclobutane were completely square planar, its bond angles would be 90° whereas tetrahedral 109.5° bond angles are expected.^[5] However, the actual C-C-C bond angle is 88° because it has a slightly folded form to relieve some torsional strain at the expense of slightly more angle strain.^[5] The high strain energy of cyclobutane is primarily from angle strain.^[6]
• cyclopentane (7.4 kcal/mol), C₅H₁₀ — if it was a completely regular planar pentagon its bond angles would be 108°, but tetrahedral 109.5° bond angles are expected.^[5] However, it has an unfixed puckered shape that undulates up and down.^[5]
• cyclohexane (1.3 kcal/mol), C₆H₁₂ — Although the chair conformation is able to achieve ideal angles, the unstablehalf-chair conformation has angle strain in the C-C-C angles which range from 109.86° to 119.07°.^[10]

Bicyclics^[11]

• bicyclo[1.1.0]butane (66.3 kcal/mol),C₄H₆
• bicyclo[1.2.0]pentane (54.7 kcal/mol),C₅H₈
• [[bicyclo[1.3.0]hexane]] (26 kcal/mol),C₆H₁₀
• norbornane (16.6 kcal/mol),C₇H₁₂

Ring strain can be considerably higher inbicyclic systems. For example,bicyclobutane, C₄H₆, is noted for being one of the most strained compounds that is isolatable on a large scale; its strain energy is estimated at 63.9 kcal mol⁻¹ (267 kJ mol⁻¹).^[12][13]

Cyclopropane has a lesser amount of ring strain since it has the least amount of unsaturation; as a result, increasing the amount of unsaturation leads to greater ring strain.^[7] For example, cyclopropene has a greater amount of ring strain than cyclopropane because it has more unsaturation.

The potential energy and unique bonding structure contained in the bonds of molecules with ring strain can be used to drive reactions inorganic synthesis. Examples of such reactions arering opening metathesis polymerisation, photo-induced ring opening ofcyclobutenes,nucleophilic ring-opening ofepoxides andaziridines, andcycloalkyne cycloaddition reactions seen in several types ofclick chemistry.

Increased potential energy from ring strain also can be used to increase the energy released by explosives or increase their shock sensitivity.^[14] For example, the shock sensitivity of the explosive1,3,3-Trinitroazetidine could partially or primarily explained by its ring strain.^[14]

• Alkane stereochemistry

• Chemical bonding
• Physical organic chemistry

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