Viewing a single comment thread. View all comments

FemChemist t1_iu4cs9r wrote

So, the example of diamond is a bit of an outlier for "covalent" bonds. More specifically, diamond is made up of covalent network bonds. These networks buck the *very general* trend of ionic bonds are stronger than covalent bonds.

But as another commenter said, one usually makes more direct comparisons, where you can actually compare energies. It's just easier to introduce trends before the more nuanced takes.

4

Live-Goose7887 t1_iu4i4j1 wrote

I'm confused by how diamond is an outlier. The bond dissociation energies in diamond are basically the same as any other C-C single bond.

2

FemChemist t1_iu4ougu wrote

Each individual bond, yes. But the vast network of these bonds adds up in ways that other covalent compounds with C-C bonds simply can't (not apples to apples, but ethane, for example).

This is just one way that talking about trends generally doesn't really tell the true story for a specific compound. Because of the sheer number of bonds, covalent network bonding is considered the strongest type of bonding (even though I'm not sure this comparison should really be made this way). And the trend is that ionic bonds are stronger than covalent bonds are stronger than intermolecular interactions, but when you start comparing actual compounds' values for bond strength, it is really easy to find exceptions.

3

Live-Goose7887 t1_iu5ryl4 wrote

Comparing network solids to small molecule compounds isn't what's confusing. Calling diamond an outlier is confusing, it's not even the most thermodynamically stable allotrope of carbon under ambient conditions. There are also countless other comparable network solids that behave like diamond (like boron nitride). My gripe is that there's absolutely nothing special about diamond from the standpoint of "bond strength," so calling it an outlier is weird.

3