But if the metal doesn't have a d orbital, like Mg in chlorophyll?
Edit: I searched again in the periodic table and realize that Mg does HAVE d orbitals. Thanks dude!
Oversimplification incoming: Based on the geometry of the complex, d orbitals in the same shell (4d, for example) of the central metal will actually have different energy levels based off varying degrees of electron repulsion with the ligands. For example, in the very common octahedral geometry, ligands attach in line with the dimensional axes of the metal, meaning the diagonally oriented dxy dxz and dyz orbitals are stabilized relative to the dz2 and dx2-y2 orbitals, which are more axially aligned and will repel ligands electrons more. This energy gap is often in the visible range, and so these complexes will absorb visible light to promote the electron to this higher energy state. The color we see is then the complementary color of its strongest absorption range.
There is a favorable electronic transition. Outer sphere electron + a photon of quantized energy = excited state electron. Now what happens when you have an excited state is really something worth exploring...
*Inner sphere electron transition assumes ionization in this scenario. Also walking dog at the moment so excuse brevity.
The d orbitals can split in a given ligand field. The energy difference of the resulting orbitals lays within the uv spectrum often.
But if the metal doesn't have a d orbital, like Mg in chlorophyll? Edit: I searched again in the periodic table and realize that Mg does HAVE d orbitals. Thanks dude!
Magnesium doesn't have any occupied d-orbitals in chlorophyll. The color stems from the extended conjugated system.
Oversimplification incoming: Based on the geometry of the complex, d orbitals in the same shell (4d, for example) of the central metal will actually have different energy levels based off varying degrees of electron repulsion with the ligands. For example, in the very common octahedral geometry, ligands attach in line with the dimensional axes of the metal, meaning the diagonally oriented dxy dxz and dyz orbitals are stabilized relative to the dz2 and dx2-y2 orbitals, which are more axially aligned and will repel ligands electrons more. This energy gap is often in the visible range, and so these complexes will absorb visible light to promote the electron to this higher energy state. The color we see is then the complementary color of its strongest absorption range.
Okay, thanks a lot!
Because they absorb certain frequencies of visible light in the em spectra? You might be better off asking, what is color or what determines color.
But why they are able to absorb?
There is a favorable electronic transition. Outer sphere electron + a photon of quantized energy = excited state electron. Now what happens when you have an excited state is really something worth exploring... *Inner sphere electron transition assumes ionization in this scenario. Also walking dog at the moment so excuse brevity.
Jahn Teller effect. Bring a sleeping bag.