Charge-shift bond

Last updated April 10, 2024

In theoretical chemistry, the charge-shift bond is a proposed new class of chemical bonds that sits alongside the three familiar families of covalent, ionic, and metallic bonds where electrons are shared or transferred respectively.^[1]^[2] The charge shift bond derives its stability from the resonance of ionic forms rather than the covalent sharing of electrons which are often depicted as having electron density between the bonded atoms. A feature of the charge shift bond is that the predicted electron density between the bonded atoms is low. It has long been known from experiment that the accumulation of electric charge between the bonded atoms is not necessarily a feature of covalent bonds.^[3]

An example where charge shift bonding has been used to explain the low electron density found experimentally is in the central bond between the inverted tetrahedral carbon atoms in [1.1.1]propellanes. Theoretical calculations on a range of molecules have indicated that a charge shift bond is present, a striking example being fluorine, F₂, which is normally described as having a typical covalent bond.^[2] The charge shift bond (CSB) has also been shown to exist at the cation-anion interface of protic ionic liquids (PILs).^[4] The authors have also shown how CSB character in PILs correlates with their physicochemical properties.

Valence bond description

The valence bond view of chemical bonding that owes much to the work of Pauling is familiar to many, if not all, chemists. The basis of Pauling's description of the chemical bond is that an electron pair bond involves the mixing, resonance, of one covalent and two ionic structures. In bonds between two atoms of the same element, homonuclear bonds, Pauling assumed that the ionic structures make no appreciable contribution to the overall bonding. This assumption followed on from published calculations for the hydrogen molecule in 1933 by Weinbaum and by James and Coolidge^[5] that showed that the contribution of ionic forms amounted to only a small percentage of the H−H bond energy. For heteronuclear bonds, A−X, Pauling estimated the covalent contribution to the bond dissociation energy as being the mean of the bond dissociation energies of homonuclear A−A and X−X bonds. The difference between the mean and the observed bond energy was assumed to be due to the ionic contribution. The calculation for HCl is shown below.^[5]

	Actual H−H	Actual Cl−Cl	H−Cl_cov Covalent bond energy H−Cl, arithmetic mean (H−H) and (Cl−Cl)	H−Cl_act Actual H−Cl	"Ionic contribution" H−Cl_act – H−Cl_cov
Bond dissociation energy(kcal mol⁻¹)	103.5	57.8	80.6	102.7	22.1

The ionic contribution to the overall bond dissociation energy was attributed to the difference in electronegativity between the A and X, and these differences were the starting point for Pauling's calculation of the individual electronegativities of the elements. The proponents of charge shift bond bonding re−examined the validity of Pauling's assumption that ionic forms make no appreciable contribution to the overall bond dissociation energies of homonuclear bonds. What they found using modern valence bond methods was that in some cases the contribution of ionic forms was significant, the most striking example being F₂, fluorine, where their calculations indicate that the bond energy of the F−F bond is due wholly to the ionic contribution.^[2]

Calculated bond energies

The contribution of ionic resonance structures has been termed the charge−shift resonance energy, RE_cs, and values have been calculated for a number of single bonds, some of which are shown below:^[2]

	Covalent contribution kcal mol⁻¹	RE_cs kcal mol⁻¹	% RE_cs contribution
H−H	95.8	9.2	8.8
Li−Li	18.2	2.8	13.1
H₃C−CH₃	63.9	27.2	30.2
H₂N−NH₂	22.8	43.8	65.7
HO−OH	–7.1	56.9	114.3
F−F	–28.4	62.2	183.9
Cl−Cl	–9.4	48.7	124.1
H−F	33.2	90.8	73.2
H−Cl	57.1	34.9	37.9
H₃C−Cl	34.0	45.9	57.4
H₃Si−Cl	37.0	65.1	63.8

The results show that for homonuclear bonds the charge shift resonance energy can be significant, and for F₂ and Cl₂ show it is the attractive component whereas the covalent contribution is repulsive. The reduced density along the bond axis density is apparent using ELF, electron localization function, a tool for determining electron density.^[2]^[6]

The bridge bond in a propellane

The bridge bond (inverted bond between the bridgehead atoms which is common to the three cycles) in a substituted [1.1.1]propellane has been examined experimentally.^[7] A theoretical study on [1.1.1]propellane has shown that it has a significant RE_cs stabilisation energy.^[8]

Factors causing charge shift bonding

Analysis of a number of compounds where charge shift resonance energy is significant shows that in many cases elements with high electronegativities are involved and these have smaller orbitals and are lone pair rich. Factors that reduce the covalent contribution to the bond energy include poor overlap of bonding orbitals, and the lone pair bond weakening effect where repulsion due to the Pauli exclusion principle is the main factor.^[2] There is no correlation between the charge−shift resonance energy RE_cs and the difference between the electronegativities of the bonded atoms as might be expected from the Pauling bonding model, however there is a global correlation between RE_cs and the sum of their electronegativities which can be accounted for in part by the lone pair bond weakening effect.^[2] The charge-shift nature of the inverted bond in [1.1.1]propellanes has been ascribed to the Pauli repulsion due to the adjacent "wing" bonds destabilising of the covalent contribution.

Experimental evidence for charge-shift bonds

The interpretation of experimentally determined electron density in molecules often uses AIM theory. In this the electron density between the atomic nuclei along the bond path are calculated, and the bond critical point where the density is at a minimum is determined. The factors that determine the type of chemical bond are the Laplacian and the electron density at the bond critical point. At the bond critical point a typical covalent bond has significant density and a large negative Laplacian. In contrast a "closed shell" interaction as in an ionic bond has a small electron density and a positive Laplacian.^[2] A charge shift bond is expected to have a positive or small Laplacian. Only a limited number of experimental determinations have been made, compounds with bonds with a positive Laplacian are the N–N bond in solid N₂O₄,^[9]^[10] and the (Mg−Mg)²⁺ diatomic structure.^[11]^{[ disputed – discuss ]}

Related Research Articles

<span class="mw-page-title-main">Chemical bond</span> Association of atoms to form chemical compounds

A chemical bond is the association of atoms or ions to form molecules, crystals, and other structures. The bond may result from the electrostatic force between oppositely charged ions as in ionic bonds or through the sharing of electrons as in covalent bonds, or some combination of these effects. Chemical bonds are described has having different strengths: there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipole–dipole interactions, the London dispersion force, and hydrogen bonding.

A covalent bond is a chemical bond that involves the sharing of electrons to form electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs. The stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding. For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full valence shell, corresponding to a stable electronic configuration. In organic chemistry, covalent bonding is much more common than ionic bonding.

Electronegativity, symbolized as χ, is the tendency for an atom of a given chemical element to attract shared electrons when forming a chemical bond. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity, the more an atom or a substituent group attracts electrons. Electronegativity serves as a simple way to quantitatively estimate the bond energy, and the sign and magnitude of a bond's chemical polarity, which characterizes a bond along the continuous scale from covalent to ionic bonding. The loosely defined term electropositivity is the opposite of electronegativity: it characterizes an element's tendency to donate valence electrons.

<span class="mw-page-title-main">Ionic bonding</span> Chemical bonding involving attraction between ions

Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding, along with covalent bonding and metallic bonding. Ions are atoms with an electrostatic charge. Atoms that gain electrons make negatively charged ions. Atoms that lose electrons make positively charged ions. This transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be more complex, e.g. molecular ions like NH⁺
₄ or SO²⁻
₄. In simpler words, an ionic bond results from the transfer of electrons from a metal to a non-metal to obtain a full valence shell for both atoms.

Quantum chemistry, also called molecular quantum mechanics, is a branch of physical chemistry focused on the application of quantum mechanics to chemical systems, particularly towards the quantum-mechanical calculation of electronic contributions to physical and chemical properties of molecules, materials, and solutions at the atomic level. These calculations include systematically applied approximations intended to make calculations computationally feasible while still capturing as much information about important contributions to the computed wave functions as well as to observable properties such as structures, spectra, and thermodynamic properties. Quantum chemistry is also concerned with the computation of quantum effects on molecular dynamics and chemical kinetics.

<span class="mw-page-title-main">Octet rule</span> Chemical rule of thumb

The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens; although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, and the 18-electron rule for transition metals.

In chemistry, resonance, also called mesomerism, is a way of describing bonding in certain molecules or polyatomic ions by the combination of several contributing structures into a resonance hybrid in valence bond theory. It has particular value for analyzing delocalized electrons where the bonding cannot be expressed by one single Lewis structure. The resonance hybrid is the accurate structure for a molecule or ion; it is an average of the theoretical contributing structures.

<span class="mw-page-title-main">Chemical polarity</span> Separation of electric charge in a molecule

In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end.

Lewis structures – also called Lewis dot formulas, Lewis dot structures, electron dot structures, or Lewis electron dot structures (LEDs) – are diagrams that show the bonding between atoms of a molecule, as well as the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently bonded molecule, as well as coordination compounds. The Lewis structure was named after Gilbert N. Lewis, who introduced it in his 1916 article The Atom and the Molecule. Lewis structures extend the concept of the electron dot diagram by adding lines between atoms to represent shared pairs in a chemical bond.

In chemistry, valence bond (VB) theory is one of the two basic theories, along with molecular orbital (MO) theory, that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the atomic orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule.

In chemistry, a hypervalent molecule is a molecule that contains one or more main group elements apparently bearing more than eight electrons in their valence shells. Phosphorus pentachloride, sulfur hexafluoride, chlorine trifluoride, the chlorite ion, and the triiodide ion are examples of hypervalent molecules.

In chemistry, orbital hybridisation is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. For example, in a carbon atom which forms four single bonds, the valence-shell s orbital combines with three valence-shell p orbitals to form four equivalent sp³ mixtures in a tetrahedral arrangement around the carbon to bond to four different atoms. Hybrid orbitals are useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Usually hybrid orbitals are formed by mixing atomic orbitals of comparable energies.

In chemistry, the valence or valency of an atom is a measure of its combining capacity with other atoms when it forms chemical compounds or molecules. Valence is generally understood to be the number of chemical bonds that each atom of a given chemical element typically forms. Double bonds are considered to be two bonds, triple bonds to be three, quadruple bonds to be four, quintuple bonds to be five and sextuple bonds to be six. In most compounds, the valence of hydrogen is 1, of oxygen is 2, of nitrogen is 3, and of carbon is 4. Valence is not to be confused with the related concepts of the coordination number, the oxidation state, or the number of valence electrons for a given atom.

In chemistry, heterolysis or heterolytic fission is the process of cleaving/breaking a covalent bond where one previously bonded species takes both original bonding electrons from the other species. During heterolytic bond cleavage of a neutral molecule, a cation and an anion will be generated. Most commonly the more electronegative atom keeps the pair of electrons becoming anionic while the more electropositive atom becomes cationic.

In chemistry, a formal charge, in the covalent view of chemical bonding, is the hypothetical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. In simple terms, formal charge is the difference between the number of valence electrons of an atom in a neutral free state and the number assigned to that atom in a Lewis structure. When determining the best Lewis structure for a molecule, the structure is chosen such that the formal charge on each of the atoms is as close to zero as possible.

Modern valence bond theory is the application of valence bond theory (VBT) with computer programs that are competitive in accuracy and economy with programs for the Hartree–Fock or post-Hartree-Fock methods. The latter methods dominated quantum chemistry from the advent of digital computers because they were easier to program. The early popularity of valence bond methods thus declined. It is only recently that the programming of valence bond methods has improved. These developments are due to and described by Gerratt, Cooper, Karadakov and Raimondi (1997); Li and McWeeny (2002); Joop H. van Lenthe and co-workers (2002); Song, Mo, Zhang and Wu (2005); and Shaik and Hiberty (2004)

The bond valencemethod or mean method is a popular method in coordination chemistry to estimate the oxidation states of atoms. It is derived from the bond valence model, which is a simple yet robust model for validating chemical structures with localized bonds or used to predict some of their properties. This model is a development of Pauling's rules.

An intramolecular force is any force that binds together the atoms making up a molecule or compound, not to be confused with intermolecular forces, which are the forces present between molecules. The subtle difference in the name comes from the Latin roots of English with inter meaning between or among and intra meaning inside. Chemical bonds are considered to be intramolecular forces which are often stronger than intermolecular forces present between non-bonding atoms or molecules.

The covalent radius of fluorine is a measure of the size of a fluorine atom; it is approximated at about 60 picometres.

Pauling's principle of electroneutrality states that each atom in a stable substance has a charge close to zero. It was formulated by Linus Pauling in 1948 and later revised. The principle has been used to predict which of a set of molecular resonance structures would be the most significant, to explain the stability of inorganic complexes and to explain the existence of π-bonding in compounds and polyatomic anions containing silicon, phosphorus or sulfur bonded to oxygen; it is still invoked in the context of coordination complexes. However, modern computational techniques indicate many stable compounds have a greater charge distribution than the principle predicts.

References

↑ Sini, Gjergji; Maitre, Philippe; Hiberty, Philippe C.; Shaik, Sason S. (1991). "Covalent, ionic and resonating single bonds". Journal of Molecular Structure: THEOCHEM. 229: 163–188. doi:10.1016/0166-1280(91)90144-9. ISSN 0166-1280.
1 2 3 4 5 6 7 8 Shaik, Sason; Danovitch, David; Wei, Wu & Hiberty, Phillippe.C. (2014) [1st. Pub. 2014]. "Chapter 5: The Valence Bond Perspective of the Chemical Bond" . In Frenking, Gernod & Shaik, Sason (eds.). The Chemical Bond . Wiley-VCH.^{[ failed verification ]}
↑ Dunitz, Jack D.; Seiler, Paul (1983). "The absence of bonding electron density in certain covalent bonds as revealed by x-ray analysis". Journal of the American Chemical Society. 105 (24): 7056–7058. doi:10.1021/ja00362a007. ISSN 0002-7863.
↑ Patil, Amol Baliram; Bhanage, Bhalchandra Mahadeo (2016). "Modern ab initio valence bond theory calculations reveal charge shift bonding in protic ionic liquids". Physical Chemistry Chemical Physics. 18: 15783–15790. doi:10.1039/C6CP02819E.
1 2 The Nature of the Chemical bond, L. Pauling, 1940, 2d edition, pp. 49−59, Oxford University Press
↑ Shaik, Sason; Danovich, David; Silvi, Bernard; Lauvergnat, David L.; Hiberty, Philippe C. (2005). "Charge−Shift Bonding—A Class of Electron-Pair Bonds That Emerges from Valence Bond Theory and Is Supported by the Electron Localization Function Approach". Chemistry: A European Journal. 11 (21): 6358–6371. doi:10.1002/chem.200500265. ISSN 0947-6539. PMID 16086335.
↑ Messerschmidt, Marc; Scheins, Stephan; Grubert, Lutz; Pätzel, Michael; Szeimies, Günter; Paulmann, Carsten; Luger, Peter (2005). "Electron Density and Bonding at Inverted Carbon Atoms: An Experimental Study of a [1.1.1]Propellane Derivative". Angewandte Chemie International Edition. 44 (25): 3925–3928. doi:10.1002/anie.200500169. ISSN 1433-7851.
↑ Shaik, Sason; Danovich, David; Wu, Wei; Hiberty, Philippe C. (2009). "Charge-shift bonding and its manifestations in chemistry". Nature Chemistry. 1 (6): 443–449. Bibcode:2009NatCh...1..443S. doi:10.1038/nchem.327. ISSN 1755-4330.
↑ Messerschmidt, Marc; Wagner, Armin; Wong, Ming Wah; Luger, Peter (2002). "Atomic Properties of N₂O₄ Based on Its Experimental Charge Density". Journal of the American Chemical Society. 124 (5): 732–733. doi:10.1021/ja011802c. ISSN 0002-7863.
↑ Tsirelson, Vladimir G.; Shishkina, Anastasia V.; Stash, Adam I.; Parsons, Simon (2009). "The experimental and theoretical QTAIMC study of the atomic and molecular interactions in dinitrogen tetroxide" (PDF). Acta Crystallographica Section B. 65 (5): 647–658. doi:10.1107/S0108768109028821. hdl: 20.500.11820/5fa0a31e-7a10-466e-a0f3-239f685217e6 . ISSN 0108-7681.
↑ Platts, James A.; Overgaard, Jacob; Jones, Cameron; Iversen, Bo B.; Stasch, Andreas (2011). "First Experimental Characterization of a Non-nuclear Attractor in a Dimeric Magnesium(I) Compound". The Journal of Physical Chemistry A. 115 (2): 194–200. Bibcode:2011JPCA..115..194P. doi:10.1021/jp109547w. ISSN 1089-5639. PMID 21158464.

This page is based on this Wikipedia article
Text is available under the CC BY-SA 4.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.

[SiniMaitre1991-1] Sini, Gjergji; Maitre, Philippe; Hiberty, Philippe C.; Shaik, Sason S. (1991). "Covalent, ionic and resonating single bonds". Journal of Molecular Structure: THEOCHEM. 229: 163–188. doi:10.1016/0166-1280(91)90144-9. ISSN 0166-1280.

[The_Chemical_bond-2] 1 2 3 4 5 6 7 8 Shaik, Sason; Danovitch, David; Wei, Wu & Hiberty, Phillippe.C. (2014) [1st. Pub. 2014]. "Chapter 5: The Valence Bond Perspective of the Chemical Bond" . In Frenking, Gernod & Shaik, Sason (eds.). The Chemical Bond . Wiley-VCH.^{[ failed verification ]}

[DunitzSeiler1983-3] Dunitz, Jack D.; Seiler, Paul (1983). "The absence of bonding electron density in certain covalent bonds as revealed by x-ray analysis". Journal of the American Chemical Society. 105 (24): 7056–7058. doi:10.1021/ja00362a007. ISSN 0002-7863.

[PatilBhanage2016-4] Patil, Amol Baliram; Bhanage, Bhalchandra Mahadeo (2016). "Modern ab initio valence bond theory calculations reveal charge shift bonding in protic ionic liquids". Physical Chemistry Chemical Physics. 18: 15783–15790. doi:10.1039/C6CP02819E.

[Pauling_Nature_Chemical_bond_2d_edition-5] 1 2 The Nature of the Chemical bond, L. Pauling, 1940, 2d edition, pp. 49−59, Oxford University Press

[ShaikDanovich2005-6] Shaik, Sason; Danovich, David; Silvi, Bernard; Lauvergnat, David L.; Hiberty, Philippe C. (2005). "Charge−Shift Bonding—A Class of Electron-Pair Bonds That Emerges from Valence Bond Theory and Is Supported by the Electron Localization Function Approach". Chemistry: A European Journal. 11 (21): 6358–6371. doi:10.1002/chem.200500265. ISSN 0947-6539. PMID 16086335.

[MesserschmidtScheins2005-7] Messerschmidt, Marc; Scheins, Stephan; Grubert, Lutz; Pätzel, Michael; Szeimies, Günter; Paulmann, Carsten; Luger, Peter (2005). "Electron Density and Bonding at Inverted Carbon Atoms: An Experimental Study of a [1.1.1]Propellane Derivative". Angewandte Chemie International Edition. 44 (25): 3925–3928. doi:10.1002/anie.200500169. ISSN 1433-7851.

[ShaikDanovich2009-8] Shaik, Sason; Danovich, David; Wu, Wei; Hiberty, Philippe C. (2009). "Charge-shift bonding and its manifestations in chemistry". Nature Chemistry. 1 (6): 443–449. Bibcode:2009NatCh...1..443S. doi:10.1038/nchem.327. ISSN 1755-4330.

[MesserschmidtWagner2002-9] Messerschmidt, Marc; Wagner, Armin; Wong, Ming Wah; Luger, Peter (2002). "Atomic Properties of N₂O₄ Based on Its Experimental Charge Density". Journal of the American Chemical Society. 124 (5): 732–733. doi:10.1021/ja011802c. ISSN 0002-7863.

[TsirelsonShishkina2009-10] Tsirelson, Vladimir G.; Shishkina, Anastasia V.; Stash, Adam I.; Parsons, Simon (2009). "The experimental and theoretical QTAIMC study of the atomic and molecular interactions in dinitrogen tetroxide" (PDF). Acta Crystallographica Section B. 65 (5): 647–658. doi:10.1107/S0108768109028821. hdl: 20.500.11820/5fa0a31e-7a10-466e-a0f3-239f685217e6 . ISSN 0108-7681.

[PlattsOvergaard2011-11] Platts, James A.; Overgaard, Jacob; Jones, Cameron; Iversen, Bo B.; Stasch, Andreas (2011). "First Experimental Characterization of a Non-nuclear Attractor in a Dimeric Magnesium(I) Compound". The Journal of Physical Chemistry A. 115 (2): 194–200. Bibcode:2011JPCA..115..194P. doi:10.1021/jp109547w. ISSN 1089-5639. PMID 21158464.

[1]

[2]

[3]

[4]

[5]

[6]

[7]

[8]

[9]

[10]

[11]