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The activation strain model and molecular orbital theory.

Wolters LP, Bickelhaupt FM - Wiley Interdiscip Rev Comput Mol Sci (2015)

Bottom Line: Using these approaches, a causal relationship is revealed between the properties of the reactants and their reactivity, e.g., reaction barriers and plausible reaction mechanisms.These examples demonstrate how the methodology is applied to different research questions, how results are interpreted, and how insights into one chemical phenomenon can lead to an improved understanding of another, seemingly completely different chemical process.WIREs Comput Mol Sci 2015, 5:324-343. doi: 10.1002/wcms.1221.

View Article: PubMed Central - PubMed

Affiliation: Department of Theoretical Chemistry and Amsterdam Center for Multiscale Modeling (ACMM), VU University AmsterdamAmsterdam, The Netherlands; Dipartimento di Scienze Chimiche, Università degli Studi di PadovaPadova, Italy.

ABSTRACT

The activation strain model is a powerful tool for understanding reactivity, or inertness, of molecular species. This is done by relating the relative energy of a molecular complex along the reaction energy profile to the structural rigidity of the reactants and the strength of their mutual interactions: ΔE(ζ) = ΔE strain(ζ) + ΔE int(ζ). We provide a detailed discussion of the model, and elaborate on its strong connection with molecular orbital theory. Using these approaches, a causal relationship is revealed between the properties of the reactants and their reactivity, e.g., reaction barriers and plausible reaction mechanisms. This methodology may reveal intriguing parallels between completely different types of chemical transformations. Thus, the activation strain model constitutes a unifying framework that furthers the development of cross-disciplinary concepts throughout various fields of chemistry. We illustrate the activation strain model in action with selected examples from literature. These examples demonstrate how the methodology is applied to different research questions, how results are interpreted, and how insights into one chemical phenomenon can lead to an improved understanding of another, seemingly completely different chemical process. WIREs Comput Mol Sci 2015, 5:324-343. doi: 10.1002/wcms.1221.

No MeSH data available.


Related in: MedlinePlus

Different overlap situations for the metal d orbital with (a) a carbon–hydrogen bond and (b) a carbon–halogen bond.
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fig09: Different overlap situations for the metal d orbital with (a) a carbon–hydrogen bond and (b) a carbon–halogen bond.

Mentions: While the Pd–CH4 interaction gains in strength immediately at the beginning of the addition process, the interaction curves for the halomethanes ‘lag behind’, that is, are initially weaker but later on descend more steeply and arrive at equal, and even lower values. This lag is caused by the different orbital-electronic structure of the halomethane substrates as compared with methane. For each reaction, a significant part of the interaction results from donation from the Pd d orbitals to the antibonding σ* orbital of the bond being broken. Shown below in Figure 9 are the different overlap situations for a Pd d orbital with (a) a C–H σ* orbital and (b) a σ* orbital of C–X and a stretched C–X bond. For methane, the σ*C–H orbital consists of the out-of-phase combination of the CH3• sp3 lobe (again, ‘sp3’ serves merely as a general description of this orbital's shape and is not used in its strictly formal sense) with the H• 1s orbital, while for the halomethanes the σ*C–X orbital is the out-of-phase combination of the CH3• sp3 lobe with the halogen X•np orbital. The latter has an additional nodal plane centered at the halogen X, leading to partial cancelation of overlap with the Pd d orbitals in the beginning of the reaction, and a diminished (orbital) interaction energy. This delay in the build-up of stabilizing interactions causes the C–F activation barrier to be much higher than the C–H activation barrier, despite the similar bond strength. Later on, at larger C–X bond stretch (around 0.5 Å), the ΔEint curve for C–X activation catches up, as the Pd d orbitals then also favorably overlap with the σ*C–X orbital, leading to a strengthening of the interaction energy. This effect shows up not only when going from the C–H to the C–X bonds, but also along the C–X bonds from C–F to C–Cl to C–Br, which have increasingly stabilizing interaction energy curves. For the larger halogen atoms, less bond stretch is needed before good overlap with the Pd d orbital is achieved, and therefore the lag in the interaction energy term becomes less pronounced along this series. This is furthermore accompanied by the increased electron-accepting capability from the σ*C–F to the σ*C–Br orbitals, due to a lowering of the orbital energy along this series.


The activation strain model and molecular orbital theory.

Wolters LP, Bickelhaupt FM - Wiley Interdiscip Rev Comput Mol Sci (2015)

Different overlap situations for the metal d orbital with (a) a carbon–hydrogen bond and (b) a carbon–halogen bond.
© Copyright Policy - open-access
Related In: Results  -  Collection

License
Show All Figures
getmorefigures.php?uid=PMC4696410&req=5

fig09: Different overlap situations for the metal d orbital with (a) a carbon–hydrogen bond and (b) a carbon–halogen bond.
Mentions: While the Pd–CH4 interaction gains in strength immediately at the beginning of the addition process, the interaction curves for the halomethanes ‘lag behind’, that is, are initially weaker but later on descend more steeply and arrive at equal, and even lower values. This lag is caused by the different orbital-electronic structure of the halomethane substrates as compared with methane. For each reaction, a significant part of the interaction results from donation from the Pd d orbitals to the antibonding σ* orbital of the bond being broken. Shown below in Figure 9 are the different overlap situations for a Pd d orbital with (a) a C–H σ* orbital and (b) a σ* orbital of C–X and a stretched C–X bond. For methane, the σ*C–H orbital consists of the out-of-phase combination of the CH3• sp3 lobe (again, ‘sp3’ serves merely as a general description of this orbital's shape and is not used in its strictly formal sense) with the H• 1s orbital, while for the halomethanes the σ*C–X orbital is the out-of-phase combination of the CH3• sp3 lobe with the halogen X•np orbital. The latter has an additional nodal plane centered at the halogen X, leading to partial cancelation of overlap with the Pd d orbitals in the beginning of the reaction, and a diminished (orbital) interaction energy. This delay in the build-up of stabilizing interactions causes the C–F activation barrier to be much higher than the C–H activation barrier, despite the similar bond strength. Later on, at larger C–X bond stretch (around 0.5 Å), the ΔEint curve for C–X activation catches up, as the Pd d orbitals then also favorably overlap with the σ*C–X orbital, leading to a strengthening of the interaction energy. This effect shows up not only when going from the C–H to the C–X bonds, but also along the C–X bonds from C–F to C–Cl to C–Br, which have increasingly stabilizing interaction energy curves. For the larger halogen atoms, less bond stretch is needed before good overlap with the Pd d orbital is achieved, and therefore the lag in the interaction energy term becomes less pronounced along this series. This is furthermore accompanied by the increased electron-accepting capability from the σ*C–F to the σ*C–Br orbitals, due to a lowering of the orbital energy along this series.

Bottom Line: Using these approaches, a causal relationship is revealed between the properties of the reactants and their reactivity, e.g., reaction barriers and plausible reaction mechanisms.These examples demonstrate how the methodology is applied to different research questions, how results are interpreted, and how insights into one chemical phenomenon can lead to an improved understanding of another, seemingly completely different chemical process.WIREs Comput Mol Sci 2015, 5:324-343. doi: 10.1002/wcms.1221.

View Article: PubMed Central - PubMed

Affiliation: Department of Theoretical Chemistry and Amsterdam Center for Multiscale Modeling (ACMM), VU University AmsterdamAmsterdam, The Netherlands; Dipartimento di Scienze Chimiche, Università degli Studi di PadovaPadova, Italy.

ABSTRACT

The activation strain model is a powerful tool for understanding reactivity, or inertness, of molecular species. This is done by relating the relative energy of a molecular complex along the reaction energy profile to the structural rigidity of the reactants and the strength of their mutual interactions: ΔE(ζ) = ΔE strain(ζ) + ΔE int(ζ). We provide a detailed discussion of the model, and elaborate on its strong connection with molecular orbital theory. Using these approaches, a causal relationship is revealed between the properties of the reactants and their reactivity, e.g., reaction barriers and plausible reaction mechanisms. This methodology may reveal intriguing parallels between completely different types of chemical transformations. Thus, the activation strain model constitutes a unifying framework that furthers the development of cross-disciplinary concepts throughout various fields of chemistry. We illustrate the activation strain model in action with selected examples from literature. These examples demonstrate how the methodology is applied to different research questions, how results are interpreted, and how insights into one chemical phenomenon can lead to an improved understanding of another, seemingly completely different chemical process. WIREs Comput Mol Sci 2015, 5:324-343. doi: 10.1002/wcms.1221.

No MeSH data available.


Related in: MedlinePlus