# Stability of Sulfides - backbonding?

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 possess 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 ...

1) Is the above excerpt describing pi-backbonding?

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?

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.

• Silver oxide and silver sulfide have roughly similar energy of formation. Given higher energy of atomization per atom for $\ce{O2}$ molecule in comparison to $\ce{S8}$, this suggests stronger bonding in $\ce{Ag2O}$ rather than in $\ce{Ag2S}$. – permeakra Jan 20 '15 at 0:30
• @permeakra my professor makes the argument that Ag2S is MUCH less soluble than Ag2O in water (Ksp of Ag2S is ~10^-51 while Ksp of Ag2O does not even approach that by a few trillion trillion). Does his argument hold water? I was just thinking that this only suggests that one is less polar than the other (like dissolves like). – Dissenter Jan 20 '15 at 0:33
• How Ksp of Ag2O can be measured, if $\ce{O^{2-}}$ does not exist in water solutions? – permeakra Jan 20 '15 at 7:52
• @permeakra Ksp of Ag2O is not necessarily defined as [Ag+]^2[O^2-] ... goldbook.iupac.org/S05742.html – Dissenter Jan 20 '15 at 15:29

• Actually, yes .. it's part of an old convention for grouping of spectral lines .. "sharp", "principal", "diffuse" and "fundamental" .. as it became known that these grouping arose from atomic levels with different angular momentum quantum numbers, the numbers themselves took on those designations: $s$ for $l=0$, $p$ for $l=1$, $d$ for $l=2$ and so on ... but it's just coincidental that "diffuse" is also qualitatively descriptive for the nature of the d-orbitals in my answer to your question. – dtmoore1971 Jan 12 '15 at 0:47