Bonding considerations

   Analogies between groups 14 and 15 are seen if we consider certain bonding aspects. Table 1.1 lists some covalent bond enthalpy terms for group 15 elements. Data for most single bonds follow trends reminiscent of those in group 14

 These observations, together with the absence of stable P-containing analogues of N₂, NO, HCN, [N₃]^- and [NO₂]‏  indicate that strong (p–p)π-bonding is important only for the first member of group 15.  It can be argued that differences between the chemistries of nitrogen and the heavier group 15 elements (e.g. existence of PF₅, AsF₅, SbF₅ and BiF₅, but not NF₅) arise from the fact that an N atom is simply too small to accommodate five atoms around it. Historically, the differences have been attributed to the availability of d-orbitals on P, As, Sb and Bi, but not on N. However, even in the presence of electronegative atoms which would lower the energy of the d-orbitals, it is now considered that these orbitals play no significant role in hypervalent compounds of the group 15 (and later) elements. As we saw in Chapter 4, it is possible to account for the bonding  in hypervalent molecules of the p-block elements in terms of a valence set of ns and np orbitals, and we should be cautious about using sp3d and sp³d² hybridization schemes to describe trigonal bipyramidal and octahedral species of p-block elements. Although we shall show molecular structures of compounds in which P, As, Sb and Bi are in oxidation states of ‏5 (e.g. PCl₅, [PO₄]^3-, [SbF₆]^-), the representation of a line between two atoms does not necessarily mean the presence of a localized twocentre two electron bond. Similarly, the representation of a double line between two atoms does not necessarily imply that the interaction comprises covalent σ- and π- contributions. For example, while it is often convenient to draw structures for Me₃PO and PF₅ as:

it is more realistic to show the role that charge-separated species play when one is discussing the electronic distribution in ions or molecules, i.e.

  Furthermore, PF₅ should really be represented by a series of  resonance structures to provide a description that accounts for the equivalence of the two axial P^_F bonds and the equivalence of the three equatorial P^_F bonds. When we wish to focus on the structure of a molecule rather than on its bonding, charge-separated representations are not always the best option because they often obscure the observed geometry. This problem is readily seen by looking at the charge-separated representation of PF₅, in which the trigonal bipyramidal structure of PF₅ is not immediately apparent. The largest difference between groups 14 and 15 lies in the relative strengths of the N≡N (in N₂) and N^_N (in N₂H₄) bonds compared with those of C≡C and C^_C bonds (Tables 1.1).

    There is some uncertainty about a value for the N=N bond enthalpy term because of difficulty in choosing a reference compound, but the approximate value given in Table 1.1 is seen to be more than twice that of the N^_N bond, whereas the C=C bond is significantly less than twice as strong as the C^_C bond.    While N₂ is thermodynamically stable with respect to oligomerization  to species containing N^_N bonds, HC≡CH is thermodynamically unstable with respect to species with C^_C bonds. Similarly, the dimerization of P₂ to tetrahedral P₄ is thermodynamically favourable. The σ- and π-contributions that contribute to the very high strength of the N≡N bond (which makes many nitrogen compounds endothermic and most of the others only slightly exothermic).

   However, the particular weakness of the N^_N single bond calls for comment. The O^_O (146 kJ mol^-1 in H₂O₂) and F^_F (159 kJ mol^-1 in F₂) bonds are also very weak, much weaker than S^_S or Cl^_Cl bonds. In N₂H₄, H₂O₂ and F₂, the N, O or F atoms carry lone pairs, and it is believed that the N^_N, O^_O and F^_F bonds are weakened by repulsion between lone pairs on adjacent atoms (Figure 1.2).

Lone pairs on larger atoms (e.g. in Cl₂) are further apart and experience less mutual repulsion. Each N  atom in N₂ also has a nonbonding lone pair, but they are directed away from each other. Table 1.1 illustrates that N^_O, N^_F and N^_Cl are also rather weak and, again, interactions between lone pairs of electrons can be used to rationalize these data. However, when N is singly bonded to an atom with no lone pairs (e.g. H), the bond is strong. In pursuing such arguments, we must remember that in a heteronuclear bond, extra energy contributions may be attributed to partial ionic character (see Section 1.15).  Another important difference between N and the later group 15 elements is the ability of N to take part in strong hydrogen bonding (see Sections 9.6 and 14.5). This arises from the much higher electronegativity of N (x^P = 3.0) compared with values for the later elements (x^P values: P, 2.2; As, 2.2; Sb, 2.1; Bi, 2.0). The ability of the first row element to participate in hydrogen bonding is also seen in group 16 (e.g. O^_H….O and N^_H…..O interactions) and group 17 (e.g. O^_H…..F, N^_H…..F interactions). For carbon, the first member of group 14, weak hydrogen bonds (e.g. C^_H……O interactions) are important in the solid state structures of small molecules and biological systems.