Coordinate covalent bonds orbital

A coordinate covalent bond, represented by an arrow, is one in which both electrons come from the same atom that is, the bond can be regarded as being formed by the overlap of an orbital containing two electrons with an empty one. Thus trimethylamine oxide would be represented... 

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). 

Dative covalent bonds, or coordinate covalent bonds, are those in which electrons are shared (as in all covalent bonds), but in which both electrons involved in each bond are contributed from the same atom. Such bonds occur in organometallic compounds of transition metals having vacant d orbitals. It is beyond the scope of this book to discuss such bonding in detail the reader needing additional information should refer to works on organometallic compounds.12 The most common organometallic compounds that have dative covalent bonds are carbonyl compounds, which are formed from a transition metal and carbon monoxide, where the metal is usually in the -1, 0, or +1 oxidation state. In these compounds the carbon atom on the carbon monoxide acts as an electron-pair donor ... 

When atoms possess an incomplete outer shell (e.g., nonpaired electrons), yet their net charge is zero, attraction between such atoms takes place because of their strong tendency to complete their outer electron orbital shell by sharing their unpaired electrons. This gives rise to a covalent bond. One example of a covalent bond is the bimolecular chlorine gas (Cl2) (Fig. 1.1). Covalent bonding is a characteristic of some nonmetals or metalloids (bimolecular molecules), but may also arise between any two atoms when one of the atoms shares its outer-shell electron pair (Lewis base) with a second atom that has an empty outer shell (Lewis acid). Such bonds are known as coordinated covalent bonds or polar covalent bonds. They are commonly weaker than the covalent bond of two atoms which share each other s unpaired outer-shell electrons (e.g., F2 and 02). Coordinated covalent bonds often involve organometallic complexes. 

Hybridization can also help explain the existence and structure of many inorganic molecular ions. Consider, for example, the zinc compounds shown here. At the top is shown the electron configuration of atomic zinc, and just below it, of the divalent zinc ion. Notice that this ion has no electrons at all in its 4-shell. In zinc chloride, shown in the third row, there are two equivalent chlorine atoms bonded to the zinc. The bonding orbitals are of sp character that is, they are hybrids of the 4s and one 4p orbital of the zinc atom. Since these orbitals are empty in the isolated zinc ion, the bonding electrons themselves are all contributed by the chlorine atoms, or rather, the chlor ide ions, for it is these that are the bonded species here. Each chloride ion possesses a complete octet of electrons, and two of these electrons occupy each sp bond orbital in the zinc chloride complex ion. This is an example of a coordinate covalent bond, in which the bonded atom contributes both of the electrons that make up the shared pair. 

Nitrogen has three half-occupied p orbitals available for bonding, all perpendicular to one another. Since the nitrate ion is known to be planar, we are forced to assume that the nitrogen outer electrons are sp2 hybridized. The addition of an extra electron fills all three hybrid orbitals completely. Each of these filled sp2 orbitals forms a a bond by overlap with an empty oxygen 2pz orbital this, you will recall, is an example of coordinate covalent bonding, in which one of the atoms contributes both of the bonding electrons. 

The empty sp3 hybrid orbital overlaps with a filled orbital of F which holds two electrons, F + BF3 — BF4 (coordinate covalent bonding)... 

The interaction between a metal ion and a ligand can be viewed as a Lewis acid-base reaction, with the ligand donating a lone pair of electrons to an empty orbital on the metal ion to form a coordinate covalent bond ... 

The ion has a vacant orbital, which accepts a share in the lone pair on nitrogen. The formation of a covalent bond by the sharing of an electron pair that is provided by one atom is called coordinate covalent bond formation. This type of bond formation is discussed again in Chapters 10 and 25. 

Electron configurations of the elements of the three li-transition series are given in Table 25-1 and in Appendix B. Most li-transition metal ions have vacant d orbitals that can accept shares in electron pairs. Many act as Lewis acids by forming coordinate covalent bonds in coordination compounds (coordination complexes, or complex ions). Complexes of transition metal ions or molecules include cationic species (e.g., [Cr(OH2)( ]5+, [Co(NH3)g]3 +, [Ag(NH3)2]+), anionic species (e.g., [Ni(CN4)]2-, [MnCl ] ), and neutral species (e.g., [Fe(CO)5], [Pt(NH3)2Cl2]). Many complexes are very stable, as indicated by their low dissociation constants, (Section 20-6 and Appendix 1). 

These successful applications of Equation (3.1) to estimate bond energies are exceptions, rather than the rule. If AA is too large, so that ionic bonds are formed, then size factors will dominate the bonding. We want the bonding to be mainly due to electron transfer in one direction, but limited in extent. The best examples will be those where a coordinate covalent bond is formed. Charge transfer complexes should usually qualify, but only if similar molecules (and orbitals) are compared. 

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. 

In Section 10.4 we saw that the B atom in BF3 is xp -hybridized. The vacant, unhybridized 2p orbital accepts the pair of electrons from NH3. So BF3 functions as an acid according to the Lewis definition, even though it does not contain an ionizable proton. Note that a coordinate covalent bond is formed between the B and N atoms, as is the case in all Lewis acid-base reactions. 

A coordinate covalent bond is distinguished by the ligand donor atom donating both electrons (of a lone pair) to an empty orbital on the central atom to form the bond. 

Consider cobalt as an example, forming the cobalt(III) cation and subsequently an octahedral [Co(NH3)6]3+ complex ion where all Co—N bonds are identical. We commence with our basic model of the coordinate covalent bond, which requires the ligand donor group to supply a lone pair of electrons to an empty orbital on the metal. We can just about deal with this for the cobalt(III) cation, through a somewhat complicated process of ion formation, electron rearrangement, orbital hybridization and filling of empty hybrid orbitals, as depicted in Figure 3.5. 

Electron configurations of the elements of the three /-transition series are given in Table 25-1 and in Appendix B. Most i-transition metal ions have vacant d orbitals that can accept shares in electron pairs. Many act as Lewis acids by forming coordinate covalent bonds in coordination compounds (coordination complexes, or complexions). 

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... 

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