Stability of Sulfides: Lewis Acid-Base Covalent Bonding

[Qube]

Homework Statement

A stability model premised on Lewis acid-base covalent bonding can be proposed to rationalize the stability shown by sulfides such as CuS and Ag2S. For instance, although the oxide anion is a stronger Bronsted base than sulfide anion, sulfide anion can be (and usually is) a stronger Lewis base than oxide anion since S has a lower EN than O. Furthermore, recall that Lewis base strength depends on the Lewis acid with which the Lewis base interacts. For the sulfide anion, since sulfide anion possesses empty valence 3d orbitals, Lewis acid-base interactions are often very strong if the Lewis acid is a low oxidation state metal ion with a fairly large number of valence d electrons. for such cases, sulfide-to-cation sigma coordinate covalent bond formation couples with pi-type coordinate covalent bond formation because metal ion d-electron density "drifts" back to empty valence 3d orbitals on sulfide anion. This results in covalent bond interactions that have substantial, strong multiple ond character and formation of a metal-sulfide lattice which is rather macromolecular in character. In sum, such as lattice must be very stable ...

The attempt at a solution

1) Is the above excerpt describing pi-backbonding? It seems to be describing some form of backbonding because the electron density is moving away from the positively charged metal cation (rather unexpected based on superficial Columbic analysis).

2) Is pi-backbonding more favorable for metals in a low rather than a high oxidation state because such metals still have substantial electron density that needs to be stabilized - and preferably stabilized by something more electronegative than a metal?

3) Metal d-orbitals overlap with p or d-orbitals of the non-metal. How does this work? I always see diagrams such as these:

Is the overlap between the metal d-orbital and non-metal p-orbital as poor as depicted? Are the d-orbitals really at a 45 degree angle relative to the p-orbital?

3b) Why does there seem to be 4 bonds in the carbon monoxide ligand in the above diagram? Shouldn't there be instead 2 lines between C and O not 3 lines and two aligned p-orbitals? (Probably being a bit nit-picky here).

4) How well is the above excerpt written? I feel that two improvements could be made:

A) The passage seems to imply that Bronsted acidity and basicity do not depend on the corresponding base or acid. Wrong implication. HCl in HBr solvent won't be a strong acid.

B) Why electron density might just "drift" to the non-metal could be explained explicitly. I.e. electronegativity.

 

[Bystander]

1) Is the above excerpt describing pi-backbonding?

It's a little hard to say what it's trying to describe.

2) Is pi-backbonding more favorable for metals in a low

If one really has to think in terms of "backbonding," it's more favorable for a "softer" ion.

3) Metal d-orbitals overlap with p or d-orbitals of the non-metal

This is analogous to p-pi -- p-pi bonding at a higher level of occupation; the d orbitals include higher probabilities of finding electrons close to an atomic nucleus than do the s and p orbitals, making "delocalization" between a pair of nuclei with accessible and/or occupied d orbitals advantageous.

3b) Why does there seem to be 4 bonds in the carbon monoxide ligand

Someone is mixing depictions of pi-bonding with conservation of line drawings of tetravalent carbon.

4) How well is the above excerpt written?

You seem unimpressed. Check the front end of the book for the author's/authors' credentials; there's a very good possibility you will find "M. Ed." and/or "Ed. D." listed.