Spectroscopic study of Cu(II) Complexes: Crystal Field Theory Objective:
To study the effect of ligands on crystal field splitting energy ( E)
The color of coordination compounds of the transition elements is one of their characteristic properties. These colors are due to the absorption and subsequent emission of light in the visible part of the spectrum. Light in this region of the spectrum caused promotion of d–electrons from a lower to a higher energy level. The spectra which result are gen erally referred to as Electronic Spectra. Note that electronic transitions may also be effected by ultraviolet light. Because of the size of the quanta involved, electronic transitions in molecules are always accompanied by vibrational and rotational chang es, and hence a band spectrum is observed. In general, the bands which arise are much broader than bands in an infrared spectrum and are little used for identification purposes.
The crystal field theory of bonding in transition metal complexes has help appreciably to rationalize many of the physical properties of such complexes. Much of the data required for crystal field theory calculation is obtained from a study of the absorption spectra of transition metal complexes.
Regularly six–coordination is most readily pictured by placing the ligands at the plus and minus ends of the three coordinate axes. An electron in the d X2 - y2 and dz2 orbitals is therefore most effected by the field of the ligands and is raised in energy relative to an electron in the dxy, dyz and dzx orbitals. The combined energy level diagram is therefore composed of two upper orbitals, of equal energy, and three lower orbitals, which are also degenerate (Figure 4.1). The energy zero is conveniently taken as the weighted mean of the en ergies of these two sets of orbitals; the lower trio are thus stabilized by -2/5 E while the upper pair are destabilized by 3/5E, where E is the total energy separation. It is termed the crystal field splitting energy. The eg and t2g symbols are symmetry labels arising from Group theory and are now the most commonly used symbols.
Figure 4.1 Energy level diagram for the d orbital in an octahedral field. The energy gap E is often labeled 10 Dq or oct
Figure 4.2 Representation of the electronic spectra of Ti(III) complexes. The full line ( ) is a
representation of the spectrum of Ti(H2O)6 with a maximum at 20,000 cm . The Broken line (----) represents the spectrum of TiCl63-, with a maximum at 13,000 cm-1 The size of E is most readily measured spectrochemically by observing the energy of the electronic transitions between the t2g and eg orbitals. The energy usually lies in the visible or near ultra-violet region of the spectrum and it is such d-d transitions, which are responsible for
the colors of most transition metal compounds. The magnitude of E depends on the ligand and on the nature of the transition metal ion.
Consider the simplest possible case, namely a complex in which there is one electron in a d level, as in the ion [Ti(H2O)6]3+. The transition from the t2g to eg level of the single electron, give rise to single absorption band in the visible region (Figure 4.2). The spectrum appears as a single absorption band 20,000 cm-1. The water ligands have split the degeneracy of the free gaseous ion in two: the t2g and eg level, as decribed above, and in the ground state the electron occupies the t2g level. Irradiation of the complex with light of an appropriate frequency results in excitation of the electron in the lower t2g level to the higher eg level. Knowing that light of frequency 20,000 cm-1 is required the crystal field splitting energy for a d1 system in an octahedral field of water ligand can be calculated. However, this experiment, we will observe this phenomenon by using copper(II) complexes, due to their vivid colors.
4.1 Preparation of Cis–diaquabis(glycinato)copper(II)monohydrate Dissolve about...