Bonding in coordination compounds

In this chapter, we discuss mostly the bonding in mononuclear homoleptic complexes ML using two simple models. The first, called crystal field theory (CFT), assumes that the bonding is ionic i.e., it treats the interaction between the metal ion (or atom) and ligands to be purely electrostatic. In contrast, the second model, namely the molecular orbital theory, assumes the bonding to be covalent. A comparison between these models will be made. 

With some ligands, such as Cl , Ni(ll) forms four-coordinate complexes that are paramagnetic and tetrahedral. For these cases, VB theory assumes the d orbital occupation of the complex to be the same as that of the free ion, which eliminates the possibility that valence-level d orbitals can accept electron pairs from the ligands. Hybrid orbitals of either the sp or sd type (the latter involving n-level d orbitals) or a combination of the two provide the proper symmetry for the a bonds as well as 

The valence bond picture for six-coordinate octahedral complexes involves d2spi hybridization of the metal (Fig. 1 l.lc. d). The specific d orbitals that meet the symmetry requirements for the metal-ligand cr bonds are the dj and (page 396). As with the four-coordinate ri8 complexes discussed above, the presence of unpaired electrons in some octahedral compounds renders the valence level (n - i)d orbitals unavailable for bonding. This is true, for instance, for paramagnetic [CoFfi]3 (Fig. I l.lc). In these cases, the VB model invokes participation of n-level d orbitals in the hybridization. 

One difficulty with the VB assumption of electron donation from ligands to metal ions is the buildup of formal negative charge on the metal. Since this is a problem that arises, in one form or another, in all complete treatments of coordination compounds, the following discussion is appropriate to all current bonding models. 

Although the above values involve very rough approximations, they do indicate qualitatively how buildup of excessive negative charge on a metal can destabilize a complex. Within the group of complexes shown. Be(H20)4)2+ and [AI(H20)J3+ are/ stable, but the other two are not. Four water molecules effectively neutralize the +2 ionic charge of beryllium, but six water molecules donate loo much electron density. 

9 An exception to this mlc is the paramagnetic complex. [R.PtOlNR1 l Ni R - [3if R - Pr1). in which the central Ni(ll) is bound to two oxygen and two nitrogen atoms in a planar arrangement. Sec Frdmmcl. T. Peters. W. Wunderlich, H. Kuchen, W. Anttew, Chert, int. Ed, Enfti. 1992, ii. 612-613. 

The bonding in coordination compounds is most easily described in the manner first proposed by the Oxford chemist N. V. Sidgwick in 1923. Sidgwick suggested that 

This is illustrated below, where I have taken the single black lines in Werner s formulae to be ionic bonds and used my s5mibol for a covalent bond (Section 8.3)  

This he called the effective atomic number rule . It is now called the inert/noble gas rule , or the 18-electron rule . The latter is based on the number of outer electrons (i.e. electrons outside an [Ar], [Kr], or [Xe4f ] core)  

Sidgwick recognized that, when the above formula for [PtClj is drawn as a Lewis structure with only dots, all the Pt-Cl bonds are the same (Pt Cl), as observed experimentally. He accordingly also wrote the 

This effectively places the negative charge on the platinum atom. 

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Comparative studies [1127] of the kinetics of decomposition of similar salts containing related pyridine ligands have been used to investigate the strength of M—N bonds in coordination compounds. Non-isothermal DSC measurements were used to determine values of E for the reactions... 

Oxygen bonding in coordination compounds. B. Jezowska-Trzebiatowska, Coord. Chem. Rev.,... 

Although the simple valence-bond approach to the bonding in coordination compounds has many deficiencies, it is still useful as a first attempt to explain the structure of many complexes. The reasons why certain ligands force electron pairing will be explored in Chapter 17, but it is clear that high- and low-spin complexes have different magnetic character, and the interpretation of the results of this technique will now be explored. 

DeKock, R. L., and Gray, H. B. (1980). Chemical Structure and Bonding. Benjamin/Cummings, Menlo Park, CA. Chapter 6 presents a good introduction to bonding in coordination compounds. 

In previous chapters, we have presented a great deal of information about structure and bonding in coordination compounds. This chapter will be devoted to describing some of the important chemistry in the broad areas of organometallic complexes and those in which there are metal-metal bonds. The body of literature on each of these topics is enormous, so the coverage here will include basic concepts and a general survey. 

Chapter 11 Coordination Chemistry Bonding, Spectra, and Magnetism 387 Bonding in Coordination Compounds 391 Valence Bond Theory 391 Crystal Field Theory 394 Molecular Orbital Theory 413 Electronic Spectra of Complexes 433 Magnetic Properties of Complexes 459... 

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