Strain (chemistry)

In the field of chemistry, molecules can experience a specific type of internal stress known as strain. A molecule is considered strained when its chemical structure undergoes stress that raises its internal energy. This internal energy is the total energy stored within the molecular structure. A strained molecule possesses an additional amount of energy compared to a strain-free reference compound. You might imagine this extra energy as a compressed spring. Just as a compressed spring must be held in place to prevent its energy from releasing, a molecule is held in an energetically unfavorable conformation by its own chemical bonds.

Thermodynamics helps us understand how these molecules behave. The equilibrium between two different molecular shapes, or conformations, is determined by the difference in their Gibbs free energy. If a transformation results in a decrease in Gibbs free energy, the process is spontaneous. This means the molecule will naturally move toward the lower energy state, which is the more stable one. While enthalpy is usually the most important factor for determining stability, entropy can also play a role. For example, in n-butane, the anti conformation is more stable by 0.9 kcal mol⁻¹. However, because there are two possible gauche conformations and only one anti conformation, entropy contributes 0.4 kcal in favor of the gauche shape.

Scientists use several methods to measure this molecular strain. One common way is to look at the standard heat of formation, denoted as ΔfH°. This is the enthalpy change that occurs when a compound is formed from its separate elements. If the measured heat of formation differs from a prediction, that difference is often attributed to strain. Researchers also use experimental heats of combustion to determine strain energy. Another method is Benson group increment theory. This theory allows scientists to predict the expected internal energy of a molecule if the necessary group increments are available. If the experimental value differs from the prediction, the gap represents the strain energy.

One major type of strain is Van der Waals strain, which is also called steric strain. This occurs when atoms are forced to get closer than their Van der Waals radii allow. This specific form of strain happens when the interacting atoms are at least four bonds away from each other. The amount of steric strain depends heavily on the size of the groups involved. Bulky groups, such as tert-butyl groups, occupy much more space than smaller methyl groups and cause greater interactions.

Nobel laureate Herbert C. Brown studied this in the reaction of trialkylamines and trimethylboron. He found that increasing the size of the alkyl groups on the amine decreased the equilibrium constant due to steric strain.

Other specific forms of steric strain include syn-pentane strain and allylic strain. Syn-pentane strain occurs when two methyl-substituted bonds are rotated in opposite directions. This forces the two terminal methyl groups into close proximity, raising the energy by 3 kcal mol⁻¹. Allylic strain, or A1,3 strain, is closely related to this effect. In a compound like 2-pentene, an ethyl substituent can rotate so that its methyl group sits near the methyl group of the olefin. To avoid this strain, these compounds often adopt a more linear shape.

Torsional strain is the resistance a molecule feels when twisting its bonds. In cyclic molecules, this is also known as Pitzer strain. It occurs when atoms separated by three bonds are placed in an eclipsed conformation rather than a more stable staggered conformation. In ethane, the energy barrier to rotate between staggered conformations is approximately 2.9 kcal mol⁻¹. While it was once thought this was due to steric interactions, recent research suggests the staggered shape is stabilized by a hyperconjugative effect. In more complex molecules like butane, the strain energy comes from both steric interactions between methyl groups and angle strain.

Ring strain, or Baeyer strain, involves the geometry of the bonds. According to VSEPR theory, electrons prefer to be as far apart as possible. In many rings, the bond angles are compressed or expanded away from this optimal value. This is very common in small cycloalkanes. For example, cyclohexane is a benchmark because it has almost no strain energy. In contrast, methylcyclopentane has a higher energy state because of the ring strain in its five-membered ring.

Small rings experience the most intense angle strain. Cyclopropane is shaped like a triangle with bond angles of only 60°. This is much lower than the preferred 109.5° angle for an sp3 hybridized carbon. Cyclobutane also experiences significant strain with angles of approximately 88°. As ring size increases, the strain energy generally decreases.

Understanding these energy differences allows scientists to use molecular mechanics to calculate how molecules will react or change in different chemical systems.