Lecture #33
  CURMUDGEON GENERAL'S WARNING. These "slides" represent highlights from lecture and are neither complete nor meant to replace lecture. It is advised not to use these as a reliable means to replace missed lecture material. Do so at risk to healthy academic performance in 09-105.
Lecture Outline Transition Metal Complexes

Crystal Field theory

Tetrahedral geometry

Hybrid orbitals and bonding

Leaving the discussion of coordination number = 6 (octahedral geometries) and proceeding to coordination number four.
Now it becomes clear that the square planar geometry correlates with the observation of extra stability of a transition metal comlex ion with a d8 configuration on the central atom
Returning to an octahedral geometry, we look at the effect of having a weak ligand plus five stronger ligands. The weak ligand is assumed to be on the z-axis which changes the stability of the transition metal ion valence orbitals directed that way
Co3+ transition metal complex ions with one of the ligands varying shifts the wavelength of light absorbed to longer wavelengths as that one ligand becomes "weaker" in the field it produces. Transition moves "easiest-to-move electron" up to next available empty orbital.
Measured optical properties (colors) of Co3+ transition metal complex ions. What do you estimate the color of the last complex to be? This last complex is meant to be just the trans geometric isomer.
The tetrahedral crystal field geometry (without demonstrating how to generate the result) turns out to be the exact inverse orbital order of that seen in the octahedral geometry.
Crystal field theory told us nothing about the actual bonds. To begin our discussion of bonding in transition metal complexes, we return to consideration of hybrid atomic orbitals. Here is a review of hybrid atomic orbitals for atoms which can have expanded octets as would be the case for our transition metal ions.
The highest occupied atomic orbital in Cr3+ and the lowest unoccupied orbitals are shown. The d2sp3 hybrids can be constructed from appropriate choices here.
A review of d2sp3 hybrid atomic orbitals
The ammonia ligands bind to the transition metal ion. In theory, the bond forms in the region of overlap between the transition metal ion's hybrid atomic d2sp3 orbital and the ligand's hybrid sp3 atomic orbital. Two electrons go here, both from the ligand, to form a sigma bond.
Looking in just the xy plane where the sigma bonds to four NH3's are located as is the dxy non-bonding orbital. The ligand atomic orbital which would overlap with the transition metal hybrid orbital (a non-bonding interaction) is shown on the x- and y-axes.  
Fe(H2O)62+ is our next example.
If we allow iron's six valence electrons to occupy the 3d orbital as indicated, each water will contribute a lone pair of electrons to a bond between its sp3 orbital and iron's d2sp3 orbital formed as indicated
Fe(CO)62+ is our next illustration. Since CO is a strong field ligand, we should not be surprised to see a low-spin complex ion electron configuration show up by necessity.
Bonding of the CO ligands to the d2sp3 hybrid orbitals chosen as illustrated is consistent with the complex ion being diamagnetic.
Looking in just the xy plane, we see that the electron in the dxy nonbonding orbital is repelled by chlorine's lone pair in its p-orbital, weaking Cl's bond to the metal.  
Can we now explain why CO (and CN-) are strong field ligands?  
Let's remind ourselves of the underlying valence electron structure associated with CO, including the empty anti-bonding molecular orbitals.  
With CO sigma bonded to the metal just as NH3 and Cl- were, the CO has its empty pi-anti-bonding orbital placed so that a delocalized molecular orbital in conjunction with the metal's dxy becomes possible. Note the constructive interference possibility.  
 Here, then is the delocalized orbital into which the previously though-to-be-non-bonding electron is placed, giving the CO bond to the metal some partial pi-bond character...making it a stronger bond!  
ZnCl42- is the next example.
Ten valence electrons fill the zinc 3d orbital. The four ligands, each contributing a lone pair to bond formation, utilize the hybrids formed from Zn's next four orbitals, the 4s and the three 4p's corresponding to sp3 hybrids and the observed tetrahedral geometry for complexes of this type.
(This was not included from an earlier lecture.) Here's another attempt to show how crystal field splitting works, but with coordination number 2, also known as linear geometry.
This orbital's electron, pointing directly at the ligands (with their electrons) is raised in energy.
On the other hand, this orbital does not point towards the ligands and is unaffected by their presence. Likewise for the dxy.