Page 45 - Advanced Organic Chemistry Part A - Structure and Mechanisms, 5th ed (2007) - Carey _ Sundberg

24 Heteroatoms with unshared electron pairs can also interact with adjacent ∗
bonds. For example, oxygen and nitrogen substituents substantially weaken an adjacent
CHAPTER 1 (geminal) C−H bond.
Chemical Bonding
and Molecular Structure – –
H H H H
+
+
R O C R O C R N CH R N C
2
2
This interaction is readily apparent in spectroscopic properties of amines. The C−H
bond that is antiperiplanar to a nitrogen unshared electron pair is lengthened and
34
weakened. Absorptions for C−H bonds that are anti to nitrogen nonbonded pairs are
shifted in both IR and NMR spectra. The C−H vibration is at higher frequency (lower
1
bond energy) and the H signal is at higher field (increased electron density), as implied
by the resonance structures. There is a stereoelectronic component in hyperconjugation.
The optimal alignment is for the C−H bond that donates electrons to be aligned
with the orbital. The heteroatom bond-weakening effect is at a maximum when the
∗
electron pair is antiperiplanar to the C−H bond, since this is the optimal alignment
∗
for the overlap of the n and orbitals (see Topic 1.2 for further discussion).
H
N C

Population of σ* orbital

weakens anti C H bond

1.1.9. Covalent and van der Waals Radii of Atoms

Covalent and van der Waals radii are other fundamental properties of atoms in
molecules that are influenced by nuclear charge and electron distribution. A glance
at a molecular model or graphic suggests that most atoms have several different
dimensions. There is the distance between each bound atom and also a dimension in
any direction in which the atom in not bonded to another atom. The former distance,
divided between the two bonded atoms, is called the covalent radius. The nonbonded
dimension of an atom or group in a molecule is called the van der Waals radius. This
is the distance at which nonbonded atoms begin to experience mutual repulsion. Just
short of this distance, the interatomic forces are weakly attractive and are referred to
as dispersion or London forces and are attributed to mutual polarization of atoms.
There are several definitions and values assigned to covalent radii. Pauling created
an early scale using bond lengths in simple homonuclear compounds as the starting
point. An extended version of this scale is listed as “covalent” in Table 1.5. A related,
but more comprehensive, approach is to examine structural data to determine covalent
radii that best correlate with observed bond distances. This approach was developed
by Slater. 35 An extensive tabulation of bond lengths derived from structural data was
published in 1987. 36 These values are labeled “structural” in Table 1.5. A set of
34 A. Pross, L. Radom, and N. V. Riggs, J. Am. Chem. Soc., 102, 2253 (1980).
35 J. C. Slater, J. Chem. Phys., 39, 3199 (1964); J. C. Slater, Quantum Theory of Molecules and Solids,
McGraw-Hill, New York, 1965, Vol. 2.
36
F. H. Allen, O. Kennard, D. G. Watson, L. Brammer, A. G. Opren and R. Taylor, J. Chem. Soc., Perkin
Trans. 2, Supplement S1–S19 (1987).