More than an Octet of Valence Electrons

The third and largest class of exceptions consists of molecules or polyatomic ions in which there are more than eight electrons in the valence shell of an atom. When we draw the Lewis structure for PF5, for example, we are forced to place ten electrons around the central phosphorus atom  

Molecules and ions with more than an octet of electrons around the central atom are often called hypervalent. Other examples of hypervalent species are SF4, AsFg, and ICl4. The corresponding molecules with a second-period atom as the central atom, such as NCI5 and OF4, do not exist. 

Hypervalent molecules are formed only for central atoms from period 3 and below in the periodic table. The principal reason for their formation is the relatively larger size of the central atom. For example, a P atom is large enough that five F (or even five Cl) atoms can be bonded to it without being too crowded. By contrast, an N atom is too small to accommodate five atoms bonded to it. Because size is a factor, hypervalent molecules occur most often when the central atom is bonded to the smallest and most electronegative atoms— F, Cl, and O. 

The notion of a valence shell containing more than an octet of electrons is also consistent with the presence of unfilled nd orbitals in atoms from period 3 and below. By comparison, elements of the second period have only the 2s and 2p valence orbitals available for bonding. Detailed analyses of the bonding in molecules such as PF5 and SF, suggest that the presence of unfilled 3d orbitals in P and S has a relatively minor impact on the formation of hypervalent molecules, and the general current belief is that the increased size of third-period atoms is the more important factor. 

Lewis Structure for an Ion with More than an Octet of Electrons 

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These examples reveal that formal charges appear on an atom that does not have its usual covalence and does not have more than an octet of valence electrons. Formal charges always occur in a molecule or ion that can conceived to be formed as a result of coordinate covalent bonding. 

No. The elements in Groups 4 through 7 do attain the octet, but the elements in Groups 2 and 3 have less than an octet. (The elements in the third and higher periods, such as Si, S, and P, may achieve more than an octet of valence electrons.)... 

Write the Lewis formula for each of the following molecules or ions. Which ones contain at least one atom with a share in more than an octet of valence electrons (a) PFg ... 

Molecules and polyatomic ions in which an atom has more than an octet of valence electrons... 

The standard group valence is 4. For Si, Ge, and Sn, however, coordination numbers exceeding 4 occur regularly. The central atoms in such species are thus hypervalent, that is, they have more than an octet of valence electrons. For example, 2,2-bipyridyl reacts with triphenylchlorosilane to yield a relatively stable pentacoordinate silicon complex. Observe that, although the Si carries a -1 formal charge, the complex as a whole is cationic ... 

For molecules with more than an octet of valence-shell electrons on the central atom, the employment of correlation diagrams has been less systematic. Instead, some individual cases have been treated to see what distortions from assumed idealized geometries might be expected. For example, the T-shaped C1F3 molecule has been treated as indicated in Fig. 4-6, where the results suggest that the angle should be 10° less than 90°, in semi-quantitative agreement with observation. 

The concept of an octet of electrons is one of the foundations of chemical bonding. In fact, C, N, and O, the three elements that occur most frequently in organic and biological molecules, rarely stray from the pattern of octets. Nevertheless, an octet of electrons does not guarantee that an inner atom is in its most stable configuration. In particular, elements that occupy the third and higher rows of the periodic table and have more than four valence electrons may be most stable with more than an octet of electrons. Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons. In the third row, phosphoms, with five valence electrons, can form as many as five bonds. Sulfur, with six valence electrons, can form six bonds, and chlorine, with seven valence electrons, can form as many as seven bonds. 

Elements beyond the second row of the periodic table can form bonds to more than four ligands and can be associated with more than an octet of electrons. These features are possible for two reasons. First, elements with > 2 have atomic radii that are large enough to bond to 5, 6, or even more ligands. Second, elements with > 2 have d orbitals whose energies are close to the energies of the valence p orbitals. An orbital overlap description of the bonding in these species relies on the participation of d orbitals of the inner atom. 

The phosphate ion, P04-, together with many oxoanions in which the notional content of the valence shell of the central atom is in excess of eight, may be described as exhibiting hypervalence, i.e. the central atom is surrounded by more than an octet of electrons in its valence shell. Such octets normally occupy the four molecular orbitals (MOs) that have contributions from the ns and np atomic orbitals of the central atom, where n is the highest value of the principal quantum... 

Note that the nitrogen atom in the ammonium ion (NH4+) has more than the usual number of bonds—four instead of three—but that it still has an octet of valence electrons. Nitrogen, oxygen, phosphorus, and sulfur form coordinate... 

Compounds of period 2 / V I elements have a maximum of four covalent bonds and an octet of valence electrons elements beyond period 2 may have more than four bonds and more than eight valence electrons. 

The hypervalent bond is found in molecules and ions in which an atom is formally surrounded by more than an octet of electrons and contains a linear three-center arrangement of atoms. One example is PF5, which has a trigonal bipyramidal (tbp) geometry. The traditional valence bond (VB) approach... 

So far our discussion of hybridization has extended only to period 2 elements, specifically carbon, nitrogen, and oxygen. The elements of period 3 and beyond introduce a new consideration because in many of their compounds these elements have more than an octet of electrons in the valence shell, as we saw in Section 9.2. 

The valence-bond model we developed for period 2 elements works well for compounds of period 3 elements so long as we have no more than an octet of electrons in... 

Of course, most molecules are more complicated than H2 or F2, and will contain more than one pair of shared electrons. If two oxygen atoms share a single pair of electrons, for example, then neither atom has an octet of valence electrons. However, if two pairs of electrons are shared, then both oxygen atoms conform to the octet rule. For diatomic nitrogen, three pairs of electrons must be shared to... 

Exceptions to the octet rule include a relatively few cases in which there are fewer than 8 electrons in the valence-shell orbitals. By contrast, the structures of many compounds of third row and heavier elements can best be described by assuming that the valence-shell orbitals hold more than an octet of electrons. 

Elements of the second period have only the 2s and 2p valence orbitals available for bonding. Because these orbitals can hold a maximum of eight electrons, we never find more than an octet of electrons around elements from the second period. Elements from the third period and beyond, however, have ns, np and unfilled nd orbitals that can be used in bonding. Eor example, the orbital diagram for the valence shell of a phosphorus atom is as follows ... 

Section 8.7 The octet rule is not obeyed in all cases. The exceptions occur when (a) a molecule has an odd number of electrons, (b) it is not possible to complete an octet aroimd an atom without forcing an unfavorable distribution of electrons, or (c) a large atom is surrounded by so many small electronegative atoms that it must have more than an octet of electrons aroimd it. In this last case we envision using the unfilled d orbitals of the large atom to "expand" the valence shell of the atom. Expanded octets are observed for atoms in the third row and beyond in the periodic table, for which low-energy d orbitals are available. 

As the next examples show, the provisional stmcture may contain one or more inner atoms with less than octets of valence electrons. These provisional stmctures must be optimized in order to reach the most stable molecular configuration. To optimize the electron distribution about an inner atom, move electrons from adjacent outer atoms to make double or triple bonds until the octet is complete. Examples and illustrate this procedure. 

The use of resonance structures such as 7 and 8 to describe bond polarity led to a subtle change in the meaning of the octet rule, namely, that an atom obeys the octet rule if it does not have more than eight electrons in its valence shell. As a result, resonance structures such as 7 and 8 are considered to be consistent with the octet rule. However, this is not the sense in which Lewis used the octet rule. According to Lewis, a structure such as 7 would not obey the octet rule because there are only three pairs of electrons in the valence shell of carbon, just as BF3 does not obey the octet rule for the same reason. Clearly the octet rule as defined by Lewis is not valid for hypervalent molecules, which do, indeed, have more than four pairs of shared electrons in the valence shell of the central atom. 

The total number of valence electrons that must appear in the Lewis structure is 24, from (2 x 7)(2Crs) + 4(C) + 6(0). Structures (b) and (c) can be rejected because they each show only 22 electrons. Furthermore, in (b), O has 4 rather than 2 bonds, and. in (c) one Cl has 2 bonds. In (a). C and O do not have their normal covalences. In (d), O has 10 electrons, though it cannot have more than an octet. 

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