The 3-center 4-electron (3c–4e) bond is a model used to explain bonding in certain hypervalent molecules such as tetratomic and hexatomic interhalogen compounds, sulfur tetrafluoride, the xenon fluorides, and the bifluoride ion. [1] [2] It is also known as the Pimentel–Rundle three-center model after the work published by George C. Pimentel in 1951, [3] which built on concepts developed earlier by Robert E. Rundle for electron-deficient bonding. [4] [5] An extended version of this model is used to describe the whole class of hypervalent molecules such as phosphorus pentafluoride and sulfur hexafluoride as well as multi-center π-bonding such as ozone and sulfur trioxide.
There are also molecules such as diborane (B2H6) and dialane (Al2H6) which have three-center two-electron bond (3c-2e) bonds.
While the term "hypervalent" was not introduced in the chemical literature until 1969, [6] Irving Langmuir and G. N. Lewis debated the nature of bonding in hypervalent molecules as early as 1921. [7] [8] While Lewis supported the viewpoint of expanded octet, invoking s-p-d hybridized orbitals and maintaining 2c–2e bonds between neighboring atoms, Langmuir instead opted for maintaining the octet rule, invoking an ionic basis for bonding in hypervalent compounds (see Hypervalent molecule, valence bond theory diagrams for PF5 and SF6). [9]
In a 1951 seminal paper, [3] Pimentel rationalized the bonding in hypervalent trihalide ions (X−
3, X = F, Br, Cl, I) via a molecular orbital (MO) description, building on the concept of the "half-bond" introduced by Rundle in 1947. [4] [5] In this model, two of the four electrons occupy an all in-phase bonding MO, while the other two occupy a non-bonding MO, leading to an overall bond order of 0.5 between adjacent atoms (see Molecular orbital description).
More recent theoretical studies on hypervalent molecules support the Langmuir view, confirming that the octet rule serves as a good first approximation to describing bonding in the s- and p-block elements. [10] [11]
The σ molecular orbitals (MOs) of triiodide can be constructed by considering the in-phase and out-of-phase combinations of the central atom's p orbital (collinear with the bond axis) with the p orbitals of the peripheral atoms. [12] This exercise generates the diagram at right (Figure 1). Three molecular orbitals result from the combination of the three relevant atomic orbitals, with the four electrons occupying the two MOs lowest in energy – a bonding MO delocalized across all three centers, and a non-bonding MO localized on the peripheral centers. Using this model, one sidesteps the need to invoke hypervalent bonding considerations at the central atom, since the bonding orbital effectively consists of two 2-center-1-electron bonds (which together do not violate the octet rule), and the other two electrons occupy the non-bonding orbital.
In the natural bond orbital viewpoint of 3c–4e bonding, the triiodide anion is constructed from the combination of the diiodine (I2) σ molecular orbitals and an iodide (I−) lone pair. The I− lone pair acts as a 2-electron donor, while the I2 σ* antibonding orbital acts as a 2-electron acceptor. [12] Combining the donor and acceptor in in-phase and out-of-phase combinations results in the diagram depicted at right (Figure 2). Combining the donor lone pair with the acceptor σ* antibonding orbital results in an overall lowering in energy of the highest-occupied orbital (ψ2). While the diagram depicted in Figure 2 shows the right-hand atom as the donor, an equivalent diagram can be constructed using the left-hand atom as the donor. This bonding scheme is succinctly summarized by the following two resonance structures: I—I···I− ↔ I−···I—I (where "—" represents a single bond and "···" represents a "dummy bond" with formal bond order 0 whose purpose is only to indicate connectivity), which when averaged reproduces the I—I bond order of 0.5 obtained both from natural bond orbital analysis and from molecular orbital theory.
More recent theoretical investigations suggest the existence of a novel type of donor-acceptor interaction that may dominate in triatomic species with so-called "inverted electronegativity"; [13] that is, a situation in which the central atom is more electronegative than the peripheral atoms. Molecules of theoretical curiosity such as neon difluoride (NeF2) and beryllium dilithide (BeLi2) represent examples of inverted electronegativity. [13] As a result of unusual bonding situation, the donor lone pair ends up with significant electron density on the central atom, while the acceptor is the "out-of-phase" combination of the p orbitals on the peripheral atoms. This bonding scheme is depicted in Figure 3 for the theoretical noble gas dihalide NeF2.
The valence bond description and accompanying resonance structures A—B···C− ↔ A−···B—C suggest that molecules exhibiting 3c–4e bonding can serve as models for studying the transition states of bimolecular nucleophilic substitution reactions. [12]
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.
The noble gases are the naturally occurring members of group 18 of the periodic table: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). Under standard conditions, these elements are odorless, colorless, monatomic gases with very low chemical reactivity and cryogenic boiling points.
In theoretical chemistry, a conjugated system is a system of connected p-orbitals with delocalized electrons in a molecule, which in general lowers the overall energy of the molecule and increases stability. It is conventionally represented as having alternating single and multiple bonds. Lone pairs, radicals or carbenium ions may be part of the system, which may be cyclic, acyclic, linear or mixed. The term "conjugated" was coined in 1899 by the German chemist Johannes Thiele.
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.
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, 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 in chlorous acid 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.
A three-center two-electron (3c–2e) bond is an electron-deficient chemical bond where three atoms share two electrons. The combination of three atomic orbitals form three molecular orbitals: one bonding, one non-bonding, and one anti-bonding. The two electrons go into the bonding orbital, resulting in a net bonding effect and constituting a chemical bond among all three atoms. In many common bonds of this type, the bonding orbital is shifted towards two of the three atoms instead of being spread equally among all three. Example molecules with 3c–2e bonds are the trihydrogen cation and diborane. In these two structures, the three atoms in each 3c-2e bond form an angular geometry, leading to a bent bond.
In chemistry, triiodide usually refers to the triiodide ion, I−
3. This anion, one of the polyhalogen ions, is composed of three iodine atoms. It is formed by combining aqueous solutions of iodide salts and iodine. Some salts of the anion have been isolated, including thallium(I) triiodide (Tl+[I3]−) and ammonium triiodide ([NH4]+[I3]−). Triiodide is observed to be a red colour in solution.
In chemistry, Bent's rule describes and explains the relationship between the orbital hybridization and the electronegativities of substituents. The rule was stated by Henry A. Bent as follows:
Atomic s character concentrates in orbitals directed toward electropositive substituents.
The bifluoride ion is an inorganic anion with the chemical formula [HF2]−. The anion is colorless. Salts of bifluoride are commonly encountered in the reactions of fluoride salts with hydrofluoric acid. The commercial production of fluorine involves electrolysis of bifluoride salts.
The covalent radius of fluorine is a measure of the size of a fluorine atom; it is approximated at about 60 picometres.
A molecular orbital diagram, or MO diagram, is a qualitative descriptive tool explaining chemical bonding in molecules in terms of molecular orbital theory in general and the linear combination of atomic orbitals (LCAO) method in particular. A fundamental principle of these theories is that as atoms bond to form molecules, a certain number of atomic orbitals combine to form the same number of molecular orbitals, although the electrons involved may be redistributed among the orbitals. This tool is very well suited for simple diatomic molecules such as dihydrogen, dioxygen, and carbon monoxide but becomes more complex when discussing even comparatively simple polyatomic molecules, such as methane. MO diagrams can explain why some molecules exist and others do not. They can also predict bond strength, as well as the electronic transitions that can take place.
In molecular mechanics, VALBOND is a method for computing the angle bending energy that is based on valence bond theory. It is based on orbital strength functions, which are maximized when the hybrid orbitals on the atom are orthogonal. The hybridization of the bonding orbitals are obtained from empirical formulas based on Bent's rule, which relates the preference towards p character with electronegativity.
In chemistry, a halogen bond occurs when there is evidence of a net attractive interaction between an electrophilic region associated with a halogen atom in a molecular entity and a nucleophilic region in another, or the same, molecular entity. Like a hydrogen bond, the result is not a formal chemical bond, but rather a strong electrostatic attraction. Mathematically, the interaction can be decomposed in two terms: one describing an electrostatic, orbital-mixing charge-transfer and another describing electron-cloud dispersion. Halogen bonds find application in supramolecular chemistry; drug design and biochemistry; crystal engineering and liquid crystals; and organic catalysis.
In quantum chemistry, a natural bond orbital or NBO is a calculated bonding orbital with maximum electron density. The NBOs are one of a sequence of natural localized orbital sets that include "natural atomic orbitals" (NAO), "natural hybrid orbitals" (NHO), "natural bonding orbitals" (NBO) and "natural (semi-)localized molecular orbitals" (NLMO). These natural localized sets are intermediate between basis atomic orbitals (AO) and molecular orbitals (MO):
In chemistry, a chalcogen bond (ChB) is an attractive interaction in the family of σ-hole interactions, along with halogen bonds. Electrostatic, charge-transfer (CT) and dispersion terms have been identified as contributing to this type of interaction. In terms of CT contribution, this family of attractive interactions has been modeled as an electron donor interacting with the σ* orbital of a C-X bond of the bond donor. In terms of electrostatic interactions, the molecular electrostatic potential (MEP) maps is often invoked to visualize the electron density of the donor and an electrophilic region on the acceptor, where the potential is depleted, referred to as a σ-hole. ChBs, much like hydrogen and halogen bonds, have been invoked in various non-covalent interactions, such as protein folding, crystal engineering, self-assembly, catalysis, transport, sensing, templation, and drug design.
Linnett double-quartet theory (LDQ) is a method of describing the bonding in molecules which involves separating the electrons depending on their spin, placing them into separate 'spin tetrahedra' to minimise the Pauli repulsions between electrons of the same spin. Introduced by J. W. Linnett in his 1961 monograph and 1964 book, this method expands on the electron dot structures pioneered by G. N. Lewis. While the theory retains the requirement for fulfilling the octet rule, it dispenses with the need to force electrons into coincident pairs. Instead, the theory stipulates that the four electrons of a given spin should maximise the distances between each other, resulting in a net tetrahedral electronic arrangement that is the fundamental molecular building block of the theory.