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- Is there a crystal-field/ligand-field theory explanation to the spectrochemical series of transition metal ions within a period?
The spectrochemical series of metals, under the circumstances that same ligands are used and that it is in an octahedral coordination, is given by:
Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < V3+ < Co3+ < Mn4+ < Mo3+ < Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+
When I was skimming through a textbook to teach my lab members spectrochemical series, I was stuck about the ordering of the metals. Textbooks usually state that crystal-field increases in strength with increasing oxidation number, and that it also increases down a group in the periodic table. I can understand this since the former decreases the ionic radii, and the latter increases the distribution of the d-orbital, allowing ligands to interact stronger (=stronger crystal-field).
Now, I know that crystal-field theory or ligand-field theory doesn't fully explain coordination chemistry, but there are something I definitely have problem understanding.
For example, how would the metals within the same period with same oxidation be explained? I can understand that 4th period divalent metal ions will often have high-spin electron configuration. In that sense, Mn2+ and beyond should have electrons in the anti-bonding eg orbitals, which will decrease the interaction between the metal and the ligand, lowering the crystal-field. However, I can't explain Ni2+ < Co2+ < Fe2+. Unless the t2g orbitals are increasing in orbital energy with increasing atomic number, it should be the other way around because ionic radius should decrease with increasing atomic number.
Some of the elements also seem to not follow the trends of oxidation, such as Mn4+ and Pd4+. What would be the factor? Is it something that can be explained by crystal-field/ligand-field theory or not? If not, then that's fine. I'm just wondering for educational purposes.
Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < V3+ < Co3+ < Mn4+ < Mo3+ < Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+
When I was skimming through a textbook to teach my lab members spectrochemical series, I was stuck about the ordering of the metals. Textbooks usually state that crystal-field increases in strength with increasing oxidation number, and that it also increases down a group in the periodic table. I can understand this since the former decreases the ionic radii, and the latter increases the distribution of the d-orbital, allowing ligands to interact stronger (=stronger crystal-field).
Now, I know that crystal-field theory or ligand-field theory doesn't fully explain coordination chemistry, but there are something I definitely have problem understanding.
For example, how would the metals within the same period with same oxidation be explained? I can understand that 4th period divalent metal ions will often have high-spin electron configuration. In that sense, Mn2+ and beyond should have electrons in the anti-bonding eg orbitals, which will decrease the interaction between the metal and the ligand, lowering the crystal-field. However, I can't explain Ni2+ < Co2+ < Fe2+. Unless the t2g orbitals are increasing in orbital energy with increasing atomic number, it should be the other way around because ionic radius should decrease with increasing atomic number.
Some of the elements also seem to not follow the trends of oxidation, such as Mn4+ and Pd4+. What would be the factor? Is it something that can be explained by crystal-field/ligand-field theory or not? If not, then that's fine. I'm just wondering for educational purposes.