Iodine bonding atomic radius

Scientists have developed a variety of methods for measuring the distances separating nuclei in molecules. From observations of these distances in many molecules, each element can be assigned a bonding atomic radius. For example, in the I2 molecule, the distance separating the iodine nuclei is observed to be 2.66 A. We can define the bonding atomic radius of iodine on this basis to... 

The Lewis dot formalism shows any halogen in a molecule surrounded by three electron lone pairs. An unfortunate consequence of this perspective is that it is natural to assume that these electrons are equivalent and symmetrically distributed (i.e., that the iodine is sp3 hybridized). Even simple quantum mechanical calculations, however, show that this is not the case [148]. Consider the diiodine molecule in the gas phase (Fig. 3). There is a region directly opposite the I-I sigma bond where the nucleus is poorly shielded by the atoms electron cloud. Allen described this as polar flattening , where the effective atomic radius is shorter at this point than it is perpendicular to the I-I bond [149]. Politzer and coworkers simply call it a sigma hole [150,151]. This area of positive electrostatic potential also coincides with the LUMO of the molecule (Fig. 4). 

There is an ill-defined boundary between molecular and polymeric covalent substances. It is often possible to recognise discrete molecules in a solid-state structure, but closer scrutiny may reveal intermolecular attractions which are rather stronger than would be consistent with Van der Waals interactions. For example, in crystalline iodine each I atom has as its nearest neighbour another I atom at a distance of 272 pm, a little longer than the I-I distance in the gas-phase molecule (267 pm). However, each I atom has two next-nearest neighbours at 350 and 397 pm. The Van der Waals radius of the I atom is about 215 pm at 430 pm, the optimum balance is struck between the London attraction between two I atoms and their mutual repulsion, in the absence of any other source of bonding. There is therefore some reason to believe that the intermolecular interaction amounts to a degree of polymerisation, and the structure can be viewed as a two-dimensional layer lattice. The shortest I-I distance between layers is 427 pm, consistent with the Van der Waals radius. Elemental iodine behaves in most respects - in its volatility and solubility, for example - as a molecular solid, but it does exhibit incipient metallic properties. 

By way of contrast, in iodocyclohexane, iodine has little preference as to whether it is axial or equatorial. The reason is that although iodine is a large atom, the C-I bond length of ca. 0.195 nm is much longer than the axial C-H bonds (ca. 0.11 nm), and although the van der Waals radius of iodine, ca. 0.22 nm, is large, 1,3-diaxial interactions that involve iodine are not significant. 

Neutron diffraction studies of crystalline PH4Br and PHJ indicate the presence of weak P-H -halogen bonding. Each iodine atom in the iodide structure is surrounded by fonr H atoms at 3.35 A (corresponding to the van der Waal radius sum) and four H atoms at a closer distance of 2.87 A. The latter may involve weak bent H bonds with P/H/1 < 180°. 

How will the lengths of the H—X and C—X bonds change from fluorine to iodine They will increase as the covalent radius of the halogen atom increases. 

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