Covalent and dative bonds

Most chemists would rationalize the existence of amine borane as due to the fact that the B atom in monomeric BH3 has only six electrons in the valence shell. The formation of the NB bond is described as due to donation of the electron lone pair on N to the electron poor boron atom. A species formed in this manner from two relatively stable chemical entities, is referred to as a complex or coordination compound. The chemical species providing the electron pair is referred to as the electron donor or the Lewis base. The bonding partner is referred to as the electron acceptor or the Lewis acid. The new bond that has been formed between the donor and acceptor atoms, has been referred to as an electron donor-acceptor bond or as a dative bond. 

Should dative bonds be counted when we define the valency of the boron and nitrogen atoms There does not seem to be general agreement between chemists at this point. Some (including the author of this book) do not include dative bonds in the valency, others count both normal and dative bonds and describe both the B and the N atom as tetravalent. Some define the valency of an atom in terms of the number of electrons used for bond formation. This means that the N atom in amine borane is pentavalent while the B atom remains trivalent [2]. 

If the two electrons between the N and B atoms in the Lewis structure of H3N BH3 were equally shared between the two atoms, the net electric charge on the acceptor (BH3) would be equal to one elementary negative charge and the net charge on the donor to one elementary positive charge. If these net charges were located on the B and N atoms, the electric dipole moment of the complex would be 

Note that the negative pole would be on the B atom which is less electronegative than the N atom. The observed dipole moment of 5.22 Debye suggest that the two electrons are not equally shared, they continue to be closer to the N than to the B atom. In the following we shall nevertheless indicate a dative bond by an arrow pointing from the donor to the acceptor atom N - B. 

The shorter bond distance in the solid phase may be interpreted as follows The crystal is stabilized by favorable Coulomb interactions, particularly dipole-dipole interactions between neighboring molecules. The greater the transfer of negative charge from the electron donor to the electron acceptor, the greater the electric dipole moment. The crystalline environment will therefore favor larger electron transfer between donor and acceptor, and hence a stronger and shorter bond. 

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Figure 13 AI-0 covalent and dative bonds in alkoxyaluminum dimers.

The normal covalent and dative bonds differ in several major aspects ... 

The distinction between normal covalent and dative bonds is underscored by interatomic distances when both types are present in the same compound, as shown in 12 and 13, below [45] ... 

The frequently observed equalization of the interatomic distances in the cages does not enable distinction between covalent and dative bonds, and this approach may look somewhat artificial to the synthetic chemist or crystallographer. But, from the view point of supramolecular chemistry we believe this formalism to be a... 

Similarly one might suspect that a singly, doubly, or triply charged metal cation is a much better electron acceptor than the same atom within a neutral molecule. We suspect, therefore, that the difference between covalent and dative bonds may be blurred in many ionic complexes. 

A secondary bond , as defined by Alcock [6-8], is an interaction between two atoms characterized by a distance longer than the sum of the covalent radii but shorter than the sum of the van der Waals radii of the corresponding atoms. Such secondary interactions are weaker than normal covalent or dative bonds, but strong enough to connect individual molecules and to modify the coordination geometry of the atoms involved. They are often present in a crystal, thus resulting in self-assembled supermolecules or supramolecular architectures. For gold complexes,... 

As is obvious from the drawings, the E atoms in these simple molecules retain a lone pair of electrons may be used for a dative bond to another ML fragment (10-13). Data on complexes of types 12 and 13 are summarized in Tables III and IV, respectively. The dilemma that arises is that from the electron counting procedures one would expect two types of bonds— covalent and dative (see Ref. 62 for a discussion of the differences between dative and covalent bonding)—for complexes 11-13, but the ML fragments... 

Chemical bonds in inorganic chemistry are not limited to covalency. Very important are the dative donor-acceptor) bonds. Unlike normal covalent bonds, which are formed by pairing of electrons with a one-electron contribution from each atom, dative bonds are two-electron bonds formed by donation of an electron pair from one atom to another. There is a general tendency to assume that the two-electron bonds between a certain pair of atoms e.g. boron and nitrogen in H2B-NH2 and HsB- NH3) are identical, regardless of the origin of electrons (i.e. covalent and dative), but it has been pointed out that a distinction between the two types should be made. ... 

Lithium iodide forms a solid complex with ammonia, Li(NH3)4l, but the related hydrate, alcoholate and amine complexes are less stable. These complexes presumably involve ion-dipole bonds (p. 115), the nitrogen lone pairs surrounding the Li+ some covalent character (dative bonding) is also permissible if s and p orbitals on the Li are invoked. The chloride, bromide and iodide of lithium are much more soluble in alcohol and ether than those of the other alkali metals, but this is not always a reliable indication of covalent character. The property is employed in separating lithium from sodium. 

From the valence bond point of view, formation of a complex involves reaction between Lewis bases (ligands) and a Lewis acid (metal or metal ion) with the formation of coordinate covalent (or dative) bonds between them. The model utilizes hybridization of metal s, p, and d valence orbitals to account for the observed structures and magnetic properties of complexes. For example, complexes of Pd(ll) and Pt(Il) are usually four-coordinate, square planar, and diamagnetic, and this arrangement is often found for Ni(II) complexes as well. Inasmuch as the free ion in the ground state in each case is paramagnetic (d, F), the bonding picture has to... 

While nonbonded electron pairs in molecules do not enter into covalent bonding in the usual sense, they may exhibit a secondary kind of valency by being transferred into vacant molecular orbitals in suitable acceptor molecules. This results in the transformation of a coordination complex in which the bond formed between the electron-pair donor and the acceptor is said to be a coordinate covalent or dative bond. Brpnsted basicity is the simplest example of coordinate covalent bond formation. A Brpnsted base donates a pair of nonbonded electrons to a vacant Is orbital of a hydrogen ion to form the conjugate acid. The o-bond formed between the base and the hydrogen ion results in the loss of identity of the nonbonded pair previously localized on the base. The formation of coordination complexes has significance in the interpretation of spectra of compounds having nonbonded electron pairs. 

Exports of structural descriptors, SMILES and InChl, provide chemical structure information in a simple tab-delimited text file containing CID or SID and either the isomeric SMILES or InChl strings. Given the very nature of the formats of SMILES and InChl, not all chemical structure information can be identically represented. For example, SMILES encodes only covalent bonds, while PubChem supports the additional concepts of ionic, complex, and dative bonds. Most small molecules in PubChem can be reproducibly interconverted between InChl, SMILES, and PubChem ASN.l formats without loss of chemical structure information. 

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