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From: email@example.com (Bill Sloman)
Subject: Re: Thermal resistance--does this sound right?
Date: 10 Sep 2002 04:56:49 -0700
References: <3D781177.firstname.lastname@example.org> <3D78271F.email@example.com> <firstname.lastname@example.org> <3D7A7F2F.F04BCD59@ieee.org>
NNTP-Posting-Date: 10 Sep 2002 11:56:49 GMT
"Sir Charles W. Shults III" wrote in message news:...
> Diamond is more strongly cross-linked (more bonds) than graphite. It
> therefore takes more energy to oxidize it because you have more bonds to
Diamond has a tetrahydral structure, which maximally separates the
four bonds in the steric space available. All of the bonds are the
same and all are symmetrical about the straight line joining the
relevant atoms (so no "pi" character).
Graphite has a layered structure, with the three in-plane bonds at 120
degrees to one another and the fourth - out-of-plane bond - at 90
degrees to all three.
The in-plane bonds are quite a lot shorter and stronger than the
out-of-plane bond. The assymmetry means that these bonds have an
appreciable "pi" character.
Cabon's four valence electrons aren't equally distributed between the
bonds - the out-of-plane bond gets less than its fair share, and the
in-plane bonds get the excess as a delocalised "pi" cloud, which is
the source of graphite's relatively high electrical conductivity.
The overall bond strength in graphite is lower, which is why hyrogen
atoms preferrentially reduce any deposited graphite back to methane,
but in so far as any of the bonds involved have a multiple character,
it is the graphite in-plane bonds.
Good practical point, wrong theoretical explanation (if they haven't
made any drastic changes to the theory since I was taught it back in
Bill Sloman, Nijmegen (ex-chemist and the son of two chemists)
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