Editing Covalent radius (section)

{{short description|Measure of the size of an atom that forms part of one covalent bond}}
{{Atomic radius}}
[[File:Radis_d'un_àtom.png | thumb | right]]
The '''covalent radius''', ''r''<sub>cov</sub>, is a measure of the size of an [[atom]] that forms part of one [[covalent bond]]. It is usually measured either in [[picometre]]s (pm) or [[angstrom]]s (Å), with 1&nbsp;Å&nbsp;= 100&nbsp;pm.

In principle, the sum of the two covalent radii should equal the covalent [[bond length]] between two atoms, ''R''(AB) = ''r''(A) + ''r''(B). Moreover, different radii can be introduced for single, double and triple bonds (r<sub>1</sub>, r<sub>2</sub> and r<sub>3</sub> below), in a purely operational sense. These relationships are certainly not exact because the size of an atom is not constant but depends on its chemical environment. For [[heteroatom]]ic A–B bonds, ionic terms may enter. Often the [[polar covalent bond]]s are shorter than would be expected based on the sum of covalent radii. Tabulated values of covalent radii are either average or idealized values, which nevertheless show a certain [[transferability (chemistry)|transferability]] between different situations, which makes them useful.

The bond lengths ''R''(AB) are measured by [[X-ray diffraction]] (more rarely, [[neutron diffraction]] on [[molecular crystal]]s). [[Rotational spectroscopy]] can also give extremely accurate values of bond lengths. For [[homonuclear]] A–A bonds, [[Linus Pauling]] took the covalent radius to be half the single-bond length in the element, e.g. ''R''(H&ndash;H, in H<sub>2</sub>)&nbsp;= 74.14&nbsp;pm so ''r''<sub>cov</sub>(H)&nbsp;= 37.07&nbsp;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,<ref>{{ cite journal |doi=10.1021/ja00346a026 |author=Sanderson, R. T. |year=1983| title=Electronegativity and Bond Energy|journal=Journal of the American Chemical Society| volume=105|pages=2259–2261 |issue=8 }}</ref> but the availability of large collections of bond lengths, which are more [[Transferability (chemistry)|transferable]], from the [[Cambridge Crystallographic Database]]<ref>{{ cite journal|author1=Allen, F. H. |author2=Kennard, O. |author3=Watson, D. G. |author4=Brammer, L. |author5=Orpen, A. G. |author6=Taylor, R. |year=1987|title=Table of Bond Lengths Determined by X-Ray and Neutron Diffraction|journal=J. Chem. Soc., Perkin Trans. 2| doi=10.1039/P298700000S1|pages= S1–S19| issue=12}}</ref><ref>{{cite journal|last1=Orpen|first1=A. Guy|last2=Brammer|first2=Lee|last3=Allen|first3=Frank H.|last4=Kennard|first4=Olga|last5=Watson|first5=David G.|last6=Taylor|first6=Robin|title=Supplement. Tables of bond lengths determined by X-ray and neutron diffraction. Part 2. Organometallic compounds and co-ordination complexes of the d- and f-block metals|journal=Journal of the Chemical Society, Dalton Transactions|pages=S1|year=1989|doi=10.1039/DT98900000S1|issue=12}}</ref> has rendered covalent radii obsolete in many situations.