Valence, coordinate covalence

A coordinate covalent bond forms between the aluminum and the hydrating water molecules. Aluminum is a Group lllA element, so the aluminum (111) cation (with a chcirge of -1-3) formally has no valence electrons. The oxygen of water has lone pairs. Therefore, water molecules most likely hydrate the cation by donating lone pairs to form coordinate covalent bonds. In this respect, you can call the water molecules ligands of the metal ion. 

These examples reveal that formal charges appear on an atom that does not have its usual covalence and does not have more than an octet of valence electrons. Formal charges always occur in a molecule or ion that can conceived to be formed as a result of coordinate covalent bonding. 

One final point about covalent bonds involves the origin of the bonding electron pair. Although most covalent bonds form when two atoms each contribute one electron, bonds can also form when one atom donates both electrons (a lone pair) to another atom that has a vacant valence orbital. The ammonium ion (NH4+), for example, forms when the two lone-pair electrons from the nitrogen atom of ammonia, NH3, bond to H +. Such bonds are called coordinate covalent bonds. 

According to the valence bond theory (Section 7.10), the bonding in metal complexes arises when a filled ligand orbital containing a pair of electrons overlaps a vacant hybrid orbital on the metal ion to give a coordinate covalent bond ... 

Valence bond theory describes the bonding in complexes in terms of two-electron, coordinate covalent bonds resulting from the overlap of filled ligand orbitals with vacant metal hybrid orbitals that point in the direction of the ligands sp (linear), sp3 (tetrahedral), dsp2 (square planar), and d2sp3 or sp3d2 (octahedral). 

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... 

Electron Concept. With the rise of the electron concept of valence, it was recognized that a covalent bond might arise in such a way as to neutralize the charges of two oppositely charged ions (normal covalence) or to create a partial charge on a hitherto uncharged atom (coordinate covalence). The former was represented by a solid line, the latter by an arrow with the head indicating the direction of the electron transfer (B- A = B- A) ... 

The shape of cation (CH3)4N" is tetrahedral. The nitrogen atom is at the centre of the cation. Nitrogen has five valence electrons. Three electrons form bonds to CH3, plus one lone pair forms coordinate covalent bond with... 

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. 

The B—N bond in the above compound is different from the covalent bonds discussed so far in the sense that both electrons are contribnted by the N atom. This type of bond is called a coordinate covalent bond (also referred to as a dative bond), defined as a covalent bond in which one of the atoms donates both electrons. Although the properties of a coordinate covalent bond do not differ from those of a normal covalent bond (because all electrons are alike no matter what their sonrce), the distinction is useful for keeping track of valence electrons and assigning formal charges. 

Valence bond (VB) theory, which helps explain bonding and structure in main-group compounds (Section 11.1), is also used to describe bonding in complex ions. In the formation of a complex ion, the filled ligand orbital overlaps the empty metal-ion orbital. The ligand (Lewis base) donates the electron pair, and the metal ion (Lewis acid) accepts it to form one of the covalent bonds of the complex ion (Lewis adduct) (Section 18.8). Such a bond, in which one atom in the bond contributes both electrons, is called a coordinate covalent bond, although, once formed, it is identical to any covalent single bond. Recall that the... 

Valence bond theory pictures bonding in complex ions as arising from coordinate covalent bonding between Lewis bases (ligands) and Lewis acids (metal ions). Ligand lone pairs occupy hybridized metal-ion orbitals to form complex ions with characteristic shapes. 

The introductory chapters of this book present some of the general principles and theories which help to interpret the experiments, for example valence theory, covalence and the periodic table, coordination, oxidation and reduction, and phase equilibria. The treatment given here is, of course, brief and should be supplemented by lectures and by further reading. It is assumed that the student is already familiar with the fundamentals of atomic structure, the periodic system, and the principles of chemical equilibrium that is, that he has had a year course in general chemistry and qualitative analysis and an... 

Thinking it Through Elements with low ionization energy readily lose electrons to form positive ions. Elements with high electron affinity readily accept electrons to form negative ions. Ionic bonds, correct choice (D), result when atoms exchange electrons. Covalent bonds, choice (A), result when atoms each contribute an electron to a shared pair. Coordinate covalent bonds, choice (C), are similar except one of the two atoms furnishes both electrons. Metallic bonds, choice (B), result when atoms free one or more valence electrons to the metal lattice. 

There are two distinct types of coordination compounds separated from one another by their reactivity, which is due to the nature of the bonding from the metal to the ligand. Complex compounds are bound by coordinate covalent bonds described by valence bond theory (see Chapter 6). Addition salts or double salts, however, are bound according to electrostatic interactions, or ionic bonds. (Turn to Chapter 8 for details on ionic bonds and salts.)... 

It is useflil to show the valence bond representations of the complexes [CoFe] and [Co(NH3)6], which can then be compared with representations from the crystal field and molecular orbital theories to be discussed later. First, we must know from experiment that [CoF ] contains four unpaired electrons, whereas [Co(NH3)g] has all of its electrons paired. Each of the ligands, as Lewis bases, contributes a pair of electrons to form a coordinate covalent bond. The valence bond theory designations of the electronic structures are shown in Figure 2.7. The bonding is described as being covalent. Appropriate combinations of metal atomic orbitals are blended together to give a new set of orbitals, called hybrid orbitals. 

Example 10-13. Sketch a qualitative MO diagram for the octahedral coordination compound, [Fe(H20) ]. For each coordinate covalent bond, the HjO ligands act as Lewis bases (cr-donors) and each donate a pair of electrons to the metal-ligand bond. Using only the [Pg.322]

As its name implies, a coordinate covalent bond (also known as a dative bond in the older literature) is a covalent bond where the valence electrons are shared between the metal and the ligand, but where the ligands donate both of the shared electrons to the bond. Thus, the ligand (L) coordinates (or adds on to) the metal (M), as shown in Figure 14.5. Coordination compounds are discussed in more detail in Chapter 15. 

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