D-orbital Splitting: Coordination Compounds vs. Transition Metal Compounds

[jd12345]

Does this splitting of d-orbitals happen only in case of coordination compounds or does this happen when transition metals form compounds too?

I have only studied about splitting of d-orbitals in case of ligands(i.e. when they form complexes). Does this happen when they form compounds too)
Like when in Fe2O3, MnO2 etc do the d-orbitals split?

 

[DrDu]

Does this splitting of d-orbitals happen only in case of coordination compounds or does this happen when transition metals form compounds too?

I have only studied about splitting of d-orbitals in case of ligands(i.e. when they form complexes). Does this happen when they form compounds too)
Like when in Fe2O3, MnO2 etc do the d-orbitals split?

Yes, it does in all kind of compounds. e.g. the oxygens arrange in these compound so as to form an octahedral or tetrahedral coordination. There is no fundamental difference between coordinative, covalent and ionic bonding which are limiting cases of the situation in a general bond.

 

[tex43]

There is a fumdemental difference between covalent and coordinate bonding. In the former and electron pair is shared between two atoms due to large overlap of adjacent atomic orbitals and the bond is very strong in the order of 300 to 450 kj per mole, whereas in the latter minimal overlap occurs with consequential less sharing if any of an electron pair. Coordinate bonds are much weaker with a range 0.1 to 20 kj per mole. Take as an example bonding between molecules of water and a copper 2+ ion. These coordinate bonds are broken easily by heat, whereas the genuine covalent bonds between the oxygen and hydrogen atoms within the water molecules are not.
Ionic bonds are completely different being the elctrostatic attraction between two oppositely charged ions formed by electron loss and gain.

 

[DrDu]

Now take e.g. MnO4-, can you tell apart the contribution of ionic, coordinative and covalent bonding?

 

[tex43]

Yes. Ionic bonding exists between this ion as a whole and a positively charged cation. Bonding between the oxygen atoms and manganese is best described as multiple via overlap between 3d orbitals on manganese and hybrid sp3 orbitals on oxygen and is thus covalent. The negative charge arises naturally from the number of electrons involved. An alternative explanation can carried out using resonance theory. As I understand it there is no coordinate bonding in the manganate (VII) ion.

 

[jd12345]

I think we deviated from the initial question -
coming back - so d orbitals do split in normal compounds too, not just coordination compounds right?

 

[DrDu]

I think we deviated from the initial question -
coming back - so d orbitals do split in normal compounds too, not just coordination compounds right?

Yes, they do.

 

[tex43]

Quite right, d orbitals are split by any crystal field surrounding the central ion. My only input was to correct the statement that there was no fundamental difference between covalent, coordinate and ionic bonding.

 

[Borek]

In HS chemistry bonds are idealized and can be classified, real bonds don't care about classification and they are never ideal - purely covalent, ionic or coordinate. From this point of view DrDu statement is perfectly correct.

 

[tex43]

Real bonds are classified in many ways, strength is one interaction is another etc. All bonds can be "thought of" as attraction between charges its how we describe that attraction that makes the difference. Clearly the attraction between two hydrogen atoms in a hydrogen molecule is different from the attraction between a molecule of water and a transition metal ion, or indeed even the attraction between the poles of a magnet.