MadSci Network: Chemistry
Query:

Re: Why are diamonds a very poor conductor of electricity, where as Graphite,

Date: Thu Feb 26 11:41:57 1998
Posted By: Maria Gelabert, postdoctoral associate, Rutgers University
Area of science: Chemistry
ID: 887794550.Ch
Message:

I personally find this to be a very interesting question, since the differences 
between graphite and diamond are so fascinating.  At first it is difficult to 
believe that these two substances are even the same element.  

Diamond is a semiconductor.  In order for electrons to conduct electricity, they 
must overcome an energy barrier known as the band gap.  This gap is defined as 
the energy difference between the valence band (lower energy) and the conduction 
band (higher energy).  The valence band is filled with electrons, whereas the 
conduction band is empty.  When the electrons obtain energy from an applied 
voltage, they can jump to the conduction band.  This is how semiconductors and 
insulators conduct electricity; their electrical conductivity is low (< 100 mho/
cm).  The value of the band gap determines whether a substance is a 
semiconductor or an insulator.  Scientists argue about where the breakoff point 
is, but generally semiconductors have band gaps around 5 eV (1 eV = 1.6x10e-19 
J) or less.  Hence, since diamond has a band gap of 5 eV, it is considered a 
wide-band-gap semiconductor.

Graphite, on the other hand, is metallic.  There is no band gap (you can think 
of it as zero band gap if you wish), so in this case the electrons don't need 
much energy at all to move through the solid.  Electrical conductivities of 
metals are usually > 100 mho/cm.  Graphite is sometimes called a "bad metal", a 
term that means its conductivity is at least 1000 times lower than copper 
(conductivity 10e6 mho/cm).  It's good enough, though, to be used as electrode 
material.

So far, I've only told you what the properties of the bulk materials diamond and 
graphite are.  I haven't really told you why, on the atomic scale, would one be 
a semiconductor and the other a metal.  Well, in both polymorphs, carbon-carbon 
bonds are fundamentally what enable electrons to pass through the solid.  The 
distances between carbon atoms play a role in that conductivity.  In diamond, 
the C-C distance is 1.546 angstroms, whereas in graphite, the C-C distance (in 
the hexagonal sheets) is 1.421 angstroms, a difference of roughly 10%.  You 
could imagine, then, that the closer the carbon atoms are to each other, the 
more electronic overlap between those atoms, and therefore the easier electrons 
will travel "through" the bonds.

A note about their structures.  Diamond has a cubic structure where all the 
carbon atoms are tetrahedrally coordinated by 3 other carbon atoms.  The carbon 
atoms in graphite are trigonally coordinated to make hexagonal sheets.  The C-C 
distance within the sheets is quite short (as mentioned, 1.421 A), but the 
distances between the sheets is considerably longer (> 3 A).  Graphite is 
therefore termed an anisotropic conductor because good conductivity takes place 
in two out of three directions. The conductivity measured on a bulk graphite 
sample is an average of the conductivity in three directions.  These sheets, 
incidentally, are responsible for graphite being used in pencils.  



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