Mostly playing around in molcalc these days, not playing in the lab, but I think I found something interesting enough to share anyway.
Look at this molecule (1). Why the hell does it have such a large dipole of 25 Debye?! Surely something is wrong here. Furthermore I really don't get the
bonding situation here even when looking at the calculated orbitals. The related molecule (2) seems normal with a dipole around 10 Debye so why does the calculations seem to fail for (1)? Could this be reflected in
reality in any way? Both of these molecules have been synthesized according to the litterature, by the way, through reacting alkali perchloramides
with the respective halogens.
Here are some more subtle peculiarities and differences I noticed.
The nitrogen atom on (1) is pyramidialized but seems to be missing its negative charge present in (2) even though it is supposed to be less than
-1.
The pendant halogen atoms in (1) have equal bond lengths yet one has twice the charge unlike in (2).
The perchloryl group in (1) seems to have negative charge between the three Oxygen atoms unlike in (2).
Obviously there is a Halogen-Halogen bond in (2) but there does not seem to be one in (1).
I just don't get where this huge dipole in (1) comes from. Any ideas besides wonky model?Texium - 28-2-2024 at 11:14
Perhaps the models fail for these molecules that would very much not want to exist. Can you link the literature reference for their synthesis?Σldritch - 28-2-2024 at 11:46
Primary N-perchlorylamines formed alkali metal salts which reactedI with chlorine and bromine to give the halo derivatives. The pKa Of N-perchloryl-t-
but lamine was 6.82.
[Edited on 28-2-2024 by Σldritch]bnull - 28-2-2024 at 12:05
Are the Cl-O bonds single bonds in the first molecule?DraconicAcid - 28-2-2024 at 13:55
That's an amino version of perchloric acid. I expect that yes, it's going to be very polar and not very stable.
clearly_not_atara - 28-2-2024 at 14:53
There's a negative charge on the oxygens, a positive charge on the hydrogens, and the separation is large. At least that's my reading. The number is
still pretty unusual. The perchloryl group might not be modeled well by this software. bnull - 28-2-2024 at 16:28
Would you care to try the same calculations again, this time starting with perchloric acid and making the appropriate substitutions? I suspect that
the issue is (related to) the valence of the chlorine.
When I searched for perchloric acid (HClO4), it gave me H4ClO4 (chlorine (iv) hydroxide, orthochloric acid, pick your
choice). I had to manually insert the double bonds; otherwise, a pop up would appear when I tried to calculate properties:
Quote:
There are 35 electrons, with charge 0. Only multiplicity 1 allowed.
When I changed two Cl-O bonds from single to double to form H2ClO4 and tried to calculate properties, another pop up
came:
Quote:
Explicit valence for atom Cl, 6, is greater than permitted
The only compounds that worked were perchloric acid (here) and H3ClO4 (here and here), the latter of which I hope is inexistent.
I tested with chloric acid to see if the valence is the problem. NH2ClO2, chlorylamine (I suppose that's the name): here. NHCl(ClO2), N-chloryl-monochloramine: here. NCl2(ClO2), N-chloryl-dichloramine: here. Not a single issue.
Now try tungstic acid (H2WO4). Here comes a pop up:
Quote:
Error. Unable to optimize molecule
The program is buggy.
[Edited on 29-2-2024 by bnull]
Edit: Fixed typos.
[Edited on 29-2-2024 by bnull]DraconicAcid - 28-2-2024 at 16:37
H2ClO4 is going to be a radical.
I'm not surprised it couldn't do tungstic acid- transition metals tend to be a challenge for such programs.bnull - 28-2-2024 at 17:41
Indeed. Radicals are not accepted. H3CO gives electrons with charge 0, for example. There must be bonds or charges, even if the resulting
compound is non-existent or hypothetical.
I had the same issue a decade ago when simulating organometallics.
[Edited on 29-2-2024 by bnull]j_sum1 - 28-2-2024 at 20:16
I threw this molecule into webmo (webmo.net -- go for the demo). It did not seem to render well -- or more to the point, I did not trust it because
the bond lengths seemed to be too uniform.
I then used the webmo app on my phone. It did look a bit different. Unfortunately this app does not calculate dipole moment (unless I missed
something.) Interestingly, the molecule crashed the app when I set it doing some calculations. This has never happened before. I am going to
conclude that the molecule is really unstable.bnull - 29-2-2024 at 07:19
If someone else wants to try other perchloryl compounds, I have attached a section from volume 19 of Advances in Inorganic Chemistry and
Radiochemistry.
so why does the calculations seem to fail for (1)?
Hmm..
You're comparing a molecule with bromine, which is less electronegative than a molecule with chlorine.
Debye units require that you know an 'amount' of charge and it's physical location and separation.
It's predicated, mathematically on a fiction, where the 'center' of the positive and negative charges are located at two exact locations.
I don't have your level of intuition for how these bonds ought to behave. But, it's obvious that the program has to make an arbitrary/vague decision
about how to convert a 3D spatial charge into a 2D debye model of the 'center of *two* hypothetical charges' to replace a distribution of charges
where there are much more than two centers.
I also know that the energy of the neighboring atoms (bond energy) determines a frequency of vibration.
FTIR spectra kind of makes that obvious, and Linus Pauling comments on it in his book "The nature of the chemical Bond."
If I were to consider the two molecules you've linked (as a crude approx) to have the same order of bond lengths, then I would expect the chlorine
atom to have a larger debye moment than a bromine based on chloring displacing *more* charge over approximately the same order of distance.
But, since the "Bromine" molecule looks normal to you; then it should be obvious that it's already at the 'extreme' of normal at 10 Debye. But,
Chlorine (qualitatively), would move more charge.
I think Bromine:Chlorine has a dipole moment of 0.518 Dalton. Using that as a blind guess at an order of magnitude; I would expect the dipole moment
for the two molecules to differ by more than 2 atoms x .518Dalton, 1 Dalton.
I would have expected something like 10Daltons for the Bromine version, and 11 to 12 Daltons for the Chlorine.
What amount of dipole moment would you (without Molcalc), have expected the Chlorinated molecule to formed given that the Bromine molecule really does
have a 10 Debye moment?
