Seeing an
orbital
A paper published in the research journal
Nature in 1999 (Direct observation of
d-orbital holes and Cu-Cu bonding in
Cu2O, J. M. Zuo, M. Kim, M.
O'Keeffe, and J. C. H. Spence, Nature,
401, 49, 1999) caused great excitement in
the chemistry world. Through a series of very
careful experiments using modern X-ray and electron
diffraction methods, these authors measured the
electron distribution around Cu atoms in solid
Cu2O, copper(I) oxide. They found direct
evidence for a charge distribution around each Cu
atom that should look familiar to you:
This picture, adapted from their article, shows
electron density as both contours of constant
density (the dashed lines) and shades of color.
Look at the dark purple region: it's a
dz2 orbital!
Interestingly, in this compound, they found that
this orbital was empty. The way they
analyzed their data let them highlight regions in
space that had less electron density than a
spherical Cu+ ion would have. (Remember:
Cu has the electron configuration [Ar]
3d10 4s1 so that
bare Cu+ has the spherical
[Ar] 3d10 configuration.) They
were able to interpret this loss of electron
density in a way that showed Cu+
in the solid was non-spherical. Why? Because the
electrons that should have been in this orbital
were off dancing with another
Cu+, forming Cu–Cu
bonds throughout the solid!
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