Re: Stable 5-coordinate carbon

From: JohnWW (JohnWW_at_x.co.nz)
Date: 07/31/04 

Date: Sat, 31 Jul 2004 23:36:55 +1200

I stand by my observation as to the implications of the Pauli Exclusion
Principle. It rules out the possibility of pentavalent carbon, and
indeed of pentacoordinated carbon unless there is delocalization of
electrons over 3 or more atoms with the average order of the covalent
bonds formed by the carbon atom being not more than 4/5.

As for "CH5+", formed as an adduct of H2 to the carbonium ion CH3+, this
would not contain pentacoordinated carbon. If it is isolatable, the H2
would donate, via ONE only of the H atoms, its electron pair to the
empty orbital on the C, over which and the two H atoms from the H2 this
electron pair would be delocalized in a 3-center bond.

BTW Regarding the possibility of NF5, not with 5 covalent bonds for the
above reason (and for steric reasons also), but possibly as NF4+F-,
electron affinity is far from being everything when it comes to salts of
resonance-stabilized oxy-anions of strong acids like NO3-, ClO4-, MnO4-,
SO4--, as well as fluoro-acids like BF4-, PF6-, etc., in which the
negative charges are delocalized over more than one O or F.

John W.

Angelo wrote:
>
> JohnWW wrote:
>
> > Because the formation of 5 covalent bonds to carbon, using 2s and 2p
> > orbitals to form a trigonal bipyramidal structure, would violate the
> > Pauli exclusion principle, noting also that the energy levels of the
>
> Please excuse me, but I can't see such a violation. The Pauli
> exclusion principle states that no more than 2 electrons could
> be in every given orbital, provided their spins are opposite.
> In a 5-coordinate C atom we have the same (at max) 8 valence
> shell electrons in no less than 4 orbitals. Let's look at simple
> example, i.e. (CH5)(+). This may be seen as an adduct of (CH3)(+)
> with H2. So, we have 3 (quasi) normal C-H bonds and the
> initially empty orbital on C that impinges on the electron density of
> the H2 filled, bonding molecular orbital. How many electrons in
> how many orbitals?
>
> > next orbitals, 3s and 3p, are much too high to be usable, there is no
> > way that the bonds formed by pentacoordinated carbon could be ordinary
> > bonds of unit order. They would have to be of order 4/5 or less, in
> > which the electrons are delocalized over more than 2 atoms.
> >
> > For the same reasons, attempts in recent years to synthesize NF5 by
> > direct combination have not succeeded, in spite of the existence of
> > pentavalent N oxo-compounds in which pentavalence is achieved with one
>
> IMO it's not for the same reason, but raw stuff of ionization
> energy and electron affinity mismatching, i.e. they are too
> much different. F(-) (E.A. ca. 3.5 eV) yields typical ionic
> compounds with alkali metals (I.E. 5 eV or less) thanks to
> to coulombic energy gain at equilibrium distance of the ion
> pair and the additional Madelung lattice energy.
>
> > ionic bond entailing a positive charge on the N. It could conceivably
> > exist as a partly ionic compound, [NF4]+F-, comparable to NO2F, but the
> > NF4+ cation is even more difficult to obtain than NO2+ (which itself is
> > much more difficult to isolate than NH4+ and organic-substituted
> > ammonium salts, existing only in compounds with anions of the strongest
> > oxy-acids as in solid N2O5 (nitronium nitrate) and nitronium
>
> N2O5 is ionic in the solid state, but it's a marginal stability.
> Its vapor is discrete molecular. Also, it exists (although metastable)
> as a molecular solid. To achieve decent stability with (NO2)(+)
> (the I.E. of NO2 is 9.6 eV)
> one has to pair it with an anion from a far stronger acid than HNO3,
> e.g. (BF4)(-), (AsF6)(-), and so on.
>
> > perchlorate), and it would have to be somehow isolated with F- as the
> > balancing anion.
>
> Sorry, perhaps I don't understand what you mean,
> but if you say that (NO2)(+), or even (NF4)(+) could be
> isolated with F(-), well that's simply impossible.
> You'd need (SbF6)(-) for reasonable stability.
> BTW, the E.A. of SbF6 is 6 eV.
>
> > John W.
> >
> > P.S. My real email: JohnWW(at)xtra.co.nz (replace the (at) with @)
> >
> > Uncle Al wrote:

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