Pretty in Red? - Colour and the Spectrochemical Series

Many but by no means all coordination complexes absorb light in the visible region of the spectrum, leading to their distinctive colour. The crystal field and ligand field theory that we have developed to some extent earlier in this textbook provided a reasonable interpretation of colour. The ligands lead to, for octahedral geometry, a stabilization of diagonal (t2g) orbitals by -4D (-0.4A) and destabilization of axial (eg) orbitals by +6D, (+0.64) and a separation of Ao; for the vast majority of complexes A lies in the range from ~7000 to ~40 000 cm which places it in the near-infrared-visible-near-ultraviolet regions. Energy is required to promote an electron from a lower to a higher level, and where the energy gap between levels is equivalent to that of a region of the visible light spectrum, in achieving an excited state a select part of the coloured light spectrum is absorbed; we see the residue as colour in the complex. If we examine the octahedral splitting diagram for all of the first row transition metal ions in an octahedral field (Figure 7.4), we can appreciate the concept and even understand why some compounds are colourless.

The simplest case is d', as found for Ti3³⁺ aq. Light of the appropriate energy is absorbed, and the electron promoted from the t2g to the eg level, to produce an excited state. If the excited state decays by transferring energy to its immediate environment in various ways (such as through molecular collisions) rather than re-emitting the light energy, then we have achieved removal of part of the coloured light spectrum, and the complex appears coloured. The Zn²⁺ aq has a full complement of electrons so that electron promotion is forbidden and the compound is predicted and observed to be colourless. Further, Mn²⁺, is extremely pale in colour which we can interpret if we assign a special stability to the half-filled set of orbitals so that electron promotion is more difficult (or less 'allowed') because it would require a theoretically forbidden change in spin of the promoted electron

Figure 7.4

The high spin d-electron configurations for d1-d10 systems, with examples of aqua metal ions of these configurations, and their colours. (Although not shown in the figure, A, varies across the group of ions selected as examples.)

for it to occupy an orbital already occupied by one electron of the same spin, and hence the colour is very weak. This is a modest and not entirely satisfactory interpretation, but at least a beginning. In effect, as soon as more than one electron is present in the d orbitals, we must consider the full set of electrons rather than an individual electron. By doing this a much more satisfactory and successful interpretation of electronic spectroscopy results - but this is a task for an advanced text.

The capacity of the crystal and ligand field models to respond to and accommodate ligand-directed influences is notable. This is best exemplified by the spectrochemical series for ligands. We have discussed this earlier in Chapter 3.3. Put simply, the capacity of ligands to split the d subshell of transition metal ions is a variable, and thus produces different A for different ligands. The capacity of a particular ligand relative to others is not affected much by the particular metal ion involved, so the capacity of ligands to split can be ordered fairly consistently, as given below for an abbreviated list of ligands:

I-<Br-<CI-<F- <OH2 <NH3 <NO2- < CN- <CO.

In this spectrochemical series for ligands, those at the left split the d-orbital set least, thus favouring high-spin complexes, whereas those at the right split the d-orbital set most, thus favouring low-spin complexes. This behaviour is observable experimentally. The colour of complexes depends on the set of ligands bound, and particularly the type of donor group. Subtle variations can be detected; for example, low-spin do ColL6 ions vary from yellow (with L = CN) to orange (with L = NH₃) to blue (with L = OH₂) even though these ligands lie at one end of the series.

We can see the influence of ligands on the splitting by examining a straight-forward example the d3 system. For Cr(III) the variation in colour with ligand type is clearly identified, as complexes absorb in the visible region of the electromagnetic spectrum. The size of Δ_O for d3 is easily determined from the position of the lowest-energy absorption maxima and follows the trend shown in Table 7.6. It is clear that there is a variation with ligand that is rather substantial; further as the charge on the metal ion increases the size of the splitting increases. In effect, with complex geometry and ligand donor set held constant the variation with metal ion for a range of ligands is relatively constant and of the order for selected metal ions

Pt4+ Ir3+ Rh3+ Co3+ Cr3+ Fe3+ Fe2+ Co2+ Ni2+ Mn2+.

Whereas this latter effect can be understood in part on the basic of electrostatic arguments (ion charge and size), the variation with ligand type poses more of a problem. That there is an apparent correlation between the donor atom position in the Periodic Table and splitting energy (C> N > O> F and also F > Cl> Br) is very difficult to reconcile using a simple electrostatic view, such as met in the crystal field theory. Moreover, there is a clear trend with position in the Periodic Table for metals, with the splitting for 4d and 5d metal ions significantly larger than those for 3d. This is exemplified for the cobalt triad, where for [M(NH3)6]+ complexes, A. varies from Co (22 900 cm-1) to Rh (34 000 cm-1) to Ir (41, 200 cm¹). Again, the crystal field model struggles to interpret this observation, and we are drawn into the different ligand field model to provide the better interpretation.

Apart from the position of absorption bands, it is notable that the intensity of bands varies with coordination geometry. For perhaps the two most common shapes, octahedral and tetrahedral, it is noted that the intensity of absorbance bands for tetrahedral complexes are invariable greater (up to 50-fold) than those for octahedral complexes (Figure 7.5), although both are significantly smaller than absorbances of organic chromophores. This is the result of the transitions involving d electrons moving between d orbitals (called d-d transitions) being partly forbidden under the theory that governs our understanding of their behaviour, with the 'forbiddenness' relaxed by different effects. For tetrahedral shape, there is greater relaxation of the rules. As a general guide, for a particular coordination number the lower the symmetry the more relaxation of the rules applies and hence the larger the absorbance bands.

Figure7.5

The electronic spectra of (at left) a tetrahedral versus an octahedral complex and (at right) the shift in the spectrum of an octahedral complex on replacement of just one donor group by another with a different position in the spectrochemical series.

As a guide to how absorbance band intensity varies with selection rules that allow or forbid the transition, experimental results for some simple compounds can be compared. The high-spin octahedral d' [Mn(OH₂)₆]²⁺ ion, which is spin forbidden and Laporte for- bidden, has a very low intensity band (Emax 0.1 M1 cm^-1); the octahedral d' [Ti(OH₂)₆]³⁺ ion, which is spin allowed and Laporte forbidden, has a moderate-sized band (Emax 10 M^-1 cm^-1); the d❜ tetrahedral [CoCl4]2ion, which is spin allowed and partially Laporte al- lowed has a markedly greater band intensity (Emax 500 M1 cm^-1) whereas the octahedral do [TiCl2ion, which is spin allowed and Laporte allowed (i.e. a charge-transfer spec- trum) has a very large band intensity (Emax 10000 M^-1 cm^-1). Above, a Laporte allowed transition is one that occurs between different orbital types, such as s→p p→d or d→f; as a consequence a d→d transition is Laporte forbidden although some symmetry-based relaxation rules may operate.

It is also noted that the absorbance bands we see in the electronic spectra of d-block complexes are broad. This arises because complexes are constantly undergoing an array of molecular vibrations and rotations that, for example, are changing bond lengths slightly and thus influencing the size of Δ_O in the process. Because absorption of a photon of light is an extremely fast process compared with these minor internal structural changes in the complex, the form of the complex at the particular instant of photon capture is itself 'captured' leading to a range of energies associated with different vibrational and rotational states, so that we see an averaged outcome and a broad peak. It is notable that f-block elements display sharp absorbance bands, as the f orbitals involved are more 'buried' and overall there is little influence of rotational and vibrational motion in that block of the Periodic Table.

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مواضيع ذات صلة

Peak Performance– Illustrating Selected Physical Methods

Probing the Life of Complexes– Using Physical Methods

Organometallic Synthesis

Reactions of Coordinated Ligands

Metal-directed Reactions

Reactions Involving the Metal Oxidation State

Catalysed Reactions

Chiral Complexes

Substitution Reactions without using a Solvent

Substitution Reactions in Nonaqueous Solvents

Reactions Involving the Coordination Shell

Block f : Inner Transition Metals (Lanthanoids and Actinoids)