Stability of Sulfides: Lewis Acid-Base Covalent Bonding

This is therefore a "sneaky way" to get people to get an education degree, or two.In summary, the passage discusses the stability of sulfides in terms of Lewis acid-base covalent bonding and pi-backbonding. It suggests that sulfides can be stronger Lewis bases than oxides due to their lower electronegativity and that this type of bonding is more favorable for softer metals in lower oxidation states. However, the depiction of pi-bonding in the diagrams may not accurately reflect the actual bonding in these compounds. The excerpt could be improved by clarifying the role of electronegativity in the "drifting" of electron density and avoiding incorrect implications about Bronsted acidity and basicity.
  • #1
Qube
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Homework Statement



Textbook Excerpt said:
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:

dj038.png


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.
 
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Qube said:
1) Is the above excerpt describing pi-backbonding?
It's a little hard to say what it's trying to describe.
Qube said:
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.
Qube said:
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.
Qube said:
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.
Qube said:
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.
 

What are sulfides and how are they formed?

Sulfides are compounds that contain a sulfur atom bonded to a metal or another nonmetal. They are formed through the reaction of a sulfide ion with a metal ion or through the combination of a sulfur atom with another nonmetal element.

What is the significance of stability in sulfides?

Stability in sulfides refers to the ability of the compound to resist decomposition or chemical reactions. This is important for the practical applications of sulfides, such as in the production of batteries or as catalysts in chemical reactions.

How does Lewis acid-base covalent bonding contribute to the stability of sulfides?

Lewis acid-base covalent bonding is a type of chemical bonding in which a Lewis acid (electron acceptor) and a Lewis base (electron donor) share a pair of electrons. In sulfides, this type of bonding contributes to the stability of the compound by creating strong covalent bonds between the sulfur and the other elements, making it more difficult for the compound to break apart.

What factors affect the stability of sulfides?

The stability of sulfides can be affected by several factors, including the strength of the Lewis acid-base covalent bonds, the size and charge of the metal ion or nonmetal element, and the surrounding environmental conditions such as temperature and pressure.

How do researchers study the stability of sulfides?

Researchers study the stability of sulfides through various methods, including theoretical calculations, laboratory experiments, and analysis of the physical and chemical properties of the compound. These studies help to understand the factors that influence stability and inform the design of more stable sulfide compounds for specific applications.

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