14. Silicon

Silicon has the same structure as carbon, a hybridized, tetrahedral sp3 configuration. Since it is a larger atom, the four degenerate 3rd shell electrons form a less coherent quantum system than they do in carbon. The larger size gives silicon different (more semi-metallic) properties than carbon. (The wireframe indicates the boundary of the n=3 shell.)

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The outer spheres above simply indicate the directions of maximum electron density. The orbitals themselves will be more like spherical tetrahedra that can only occupy volume within their shell. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital and will decrease toward the nodal regions that divide the orbitals, where electron density will be lowest. Like argon, silicon features two nested spherical tetrahedra — an inner 2nd shell comprising 4 di-electrons, and an outer 3rd shell comprising 4 single electrons.

Silicon’s 3sp3 single-electron orbitals surrounding the 2sp3 shell.

Silicon can make 4 bonds, like carbon, and its natural crystal form has the same network covalent structure as diamond. This gives it a high melting point, but not the same strength or hardness as diamond because silicon does not bond with itself as strongly as carbon does. We also therefore do not see many organic molecule analogues involving silicon in place of carbon.

With a lower electronegativity than carbon, oxygen bonds more strongly to silicon than to carbon, making silicates (sand) one of the primary components of earth’s crust. Due to its larger size and less coherent hybridization, along with the presence of a d-orbital in the 3rd shell into which to expand the electron octet, silicon is also capable of octahedral bonding. Silicon-based materials (like silicates) can therefore have more complex crystalline structures than carbon-based oxides.

Since silicon becomes less of an electrical resistor as temperature increases, it is a semiconductor, with important uses in electronics and photovoltaics.


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