Copper covalent radius

Like copper, silver and gold have a single s electron outside the completed d shell, but in spite of the similarity in electronic structures and ionization potential, the chemistries of Ag, Au, and Cu differ more than might be expected. There are no simple explanations for many of the differences although some of the differences between Ag and Au may be traced to relativistic effects on the 6s electrons of the latter. The covalent radii of the triad follow the trend Cu < Ag Au, i.e., gold has about the same or a slightly smaller covalent radius than silver in comparable compounds, a phenomenon frequently referred to as relativistic contraction (c/. lanthanide contraction). 

I The small increase in covalent radius for copper and zinc has I no simple explanation. 

After rising at copper and zinc, the curve of metallic radii approaches those of the normal covalent radii and tetrahedral covalent radii (which themselves differ for arsenic, selenium, and bromine because of the difference in character of the bond orbitals, which approximate p orbitals for normal covalent bonds and sp3 orbitals for tetrahedral bonds). The bond orbitals for gallium are expected to be composed of 0.22 d orbital, one s orbital, and 2.22 p orbitals, and hence to be only slightly stronger than tetrahedral bonds, as is indicated by the fact that R(l) is smaller than the tetrahedral radius. 

Identical Ni-S and Cu-S bond distances are observed in these two structures, and the expected reduction of the ionic radius of copper is not reflected in a shorter Cu-S bond. Two possible reasons for this result are (1) a concomitant increase of the ionic radii of the sulfur donors offsets the effect of a shorter Cu(III) radius, and (2) there exists significant covalency in the Cu—S bonds. 

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