Bonding in coordination complexes

Driving off the water in this way also destroys the color, turning it from a beautiful deep blue to a nondescript pale yellow. If the anhydrous salt is now dissolved in water, the blue color now pervades the entire solution. It is apparent that the presence of water is somehow necessary for the copper(II) ion to take on a blue color, but why should this be  

A very common lab experiment that most students carry out is to add some dilute ammonia to a copper sulfate solution. At first, the solutions turns milky as the alkaline ammonia causes the precipitation of copper hydroxide Cu2+ + 2 OH Cu(OH)2(s). But if more ammonia is added, the cloudiness disappears and the solution assumes an intense deep blue color that makes the original solution seem pale by comparison. The equation for this reaction is usually given as 

The new product is commonly known as the copper-ammonia complex ion, or more officially, hexamminecopper(II). This equation is somewhat misleading, however, in that it implies the formation of a new complex where none existed before. In fact, since about 1895 it has been known that the ions of most transition metals dissolve in water to form complexes with water itself, so a better representation of the reaction of dissolved copper with ammonia would be 

In effect, the ammonia binds more tightly to the copper ion than does water, and it thus displaces the latter when it comes into contact with the hexaaquocopper(II) ion, as the dissolved form of Cu2+ is properly known. 

FIGURE 8.30 (a) Hybrid orbitals formed from linear combinations of Pm P r and orbitals. The pair of orbitals that point along the positive and negative z-axes are the same except for their orientation in space they are called axial orbitals. The set of three orbitals in the x-y plane are equivalent to one another and are called equatorial orbitals, (b) This set of hybrid orbitals can be used to describe the bonding in PF5, for example. 

We construct the first set of hybrid orbitals from one s atomic orbital, the three p atomic orbitals and the d atomic orbital they are called dsp hybrid orbitals. The principal quantum numbers of the participating atomic orbitals depend on the particular metal atom under consideration for Co, they would be the 3d, 4s, and 4p atomic orbitals. The dsp hybrid orbitals in the most general case are written out as 

The second set of hybrid orbitals we constrnct are the d sp hybrids these are six eqnivalent orbitals directed toward the vertices of an octahedron (Fig. 8.31a). They describe the structures and bonding in all of the octahedral coordination complexes discussed in Section 8.6, as well as that in SFg, which we show in Fig-nre 8.31b. 

FIGURE 8.31 (a) Hybrid orbitals formed from linear combinations of d, Px, Pj and orbitals. All 

The failure of crystal field theory and VB theory to explain the spectrochemical series stimulated the development of ligand field theory, which applies qualitative methods of molecular orbital theory to describe the bonding and structure of coordination complexes. The terms ligand field theory and molecular orbital theory are often used interchangeably in inorganic chemistry today. 

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Spectroscopic studies of metal-phosphorus bonding in coordination complexes. J. G. Verkade, Coord. Chem. Rev., 1972, 9, 1-106 (308). 

The current model of bonding in coordination complexes developed gradually between 1930-1950, and has largely superseded the hybridization model discussed previously. It is essentially a simplified adaptation of molecular orbital theory which focuses on the manner in which the electric field due to the unpaired electrons on the ligands interact with the five different d orbitals of the central ion. 

Any theory of bonding in coordination complexes must explain the experimental be-... 

The chief drawbacks to the crystal field approach are in its concept of the repulsion of orbitals by the ligands and its lack of any explanation for bonding in coordination complexes. As we have seen in all our discussions of molecular orbitals, any interaction between orbitals leads to both higher and lower energy molecular orbitals. The purely electrostatic approach does not allow for the lower (bonding) molecular orbitals, and thus fails to provide a complete picture of the electronic structure. 

Pi bonding in coordination complexes is possible when the ligand has p or tt molecular orbitals available. Because the effects are smaller for occupied orbitals, we will first treat the more important case of ligands with empty tt orbitals, or TT-acceptor ligands. 

Including tt bonding in molecular orbital theory is essential to understand the nature of bonding in coordination complexes and to account for the trends observed in the spectrochemical series. 

Bryan, R.F., Greene, P.T., Newlands, M.J. and Field, D.S. (1970) Metal-metal bonding in coordination complexes. Part X. Preparation, spectroscopic properties,... 

A simple covalent cr-bonding model employing five d orbitals and the next available s and three p orbitals, with appropriate hybridization to match a particular shape, allows a limited interpretation of bonding in coordination complexes. 

As with a polyatomic ligand, when a free ion becomes covalently bound, there is a decrease in its structural symmetry. This results in the removal of the degeneracy of some vibrations, causing new bands to be observed in the infrared (and Raman) region. Hence, infrared spectroscopy may be u.sed to distinguish between ionic and covalent bonds in coordination complexes. 

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