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From: email@example.com (Bob Wilson)
Subject: Re: Thermal resistance--does this sound right?
Date: Wed, 11 Sep 2002 02:48:35 -0000
Organization: Your Organization
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References: <3D781177.firstname.lastname@example.org> <3D78271F.email@example.com> <firstname.lastname@example.org> <3D7A7F2F.F04BCD59@ieee.org> <email@example.com>
In article <firstname.lastname@example.org>,
>"Sir Charles W. Shults III" wrote in message
>> 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
Thanks Bill. THAT was the sort of answer I was looking for.
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