In principle, the sum of the two covalent radii should equal the covalent bond length between two atoms. This relationship does not hold exactly because the size of an atom is not constant but depends on its chemical environment. In particular, polar covalent bonds tend to be shorter than would be expected on the basis of the sum of covalent radii. Tabulated values of covalent radii are either average or idealized values, which nevertheless show a certain transferability between different situations.
Covalent radii are measured by X-ray diffraction (more rarely, neutron diffraction on molecular crystals). Rotational spectroscopy can also give extremely accurate values of bond lengths. One method takes the covalent radius to be half the single bond length in the element, e.g. d(H–H, in H2) = 74.14 pm so rcov(H) = 37.07 pm: in practice, it is usual to obtain an average value from a variety of covalent compounds, although the difference is usually small. Sanderson has published a recent set of non-polar covalent radii for the main-group elements, but the availabilty of large collections of bond lengths, which are more transferable, from the Cambridge Crystallographic Database has rendered covalent radii obsolete in many situations.
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