![]() ![]() ![]() When calculating the formal charges on structures (a) and (b), we see that the S atom in (a) has a formal charge of +2, whereas the S atom in (b) has a formal charge of 0. ![]() In fact, experimental data show that the S-to-O bonds in the SO 4 2− ion are intermediate in length between single and double bonds, as expected for a system whose resonance structures all contain two S–O single bonds and two S=O double bonds. We can draw five other resonance structures equivalent to (b) that vary only in the arrangement of the single and double bonds. An alternative structure (b) can be written with S=O double bonds, making the sulfur again six-coordinate. We know that sulfur can accommodate more than eight electrons by using its empty valence d orbitals, just as in SF 6. If we use a single pair of electrons to connect the sulfur and each oxygen, we obtain the four-coordinate Lewis structure (a). Sulfate, for example, has a total of 32 valence electrons. In addition to its use as an electrical insulator, it is used as the coolant in some nuclear power plants, and it is the pressurizing gas in “unpressurized” tennis balls.Īn expanded valence shell is often written for oxoanions of the heavier p-block elements, such as sulfate (SO 4 2−) and phosphate (PO 4 3−). In fact, SF 6 is so inert that it has many commercial applications. Some species with expanded valences, such as PF 5, are highly reactive, whereas others, such as SF 6, are very unreactive. ![]() There is no correlation between the stability of a molecule or an ion and whether or not it has an expanded valence shell. Whether or not such compounds really do use d orbitals in bonding is controversial, but this model explains why compounds exist with more than an octet of electrons around an atom. Thus species such as SF 6 are often called expanded-valence molecules A compound with more than an octet of electrons around an atom. Sulfur has an 3 s 23 p 43 d 0 electron configuration, so in principle it could accommodate more than eight valence electrons by using one or more d orbitals. To accommodate more than eight electrons, sulfur must be using not only the ns and np valence orbitals but additional orbitals as well. The octet rule is based on the fact that each valence orbital (typically, one ns and three np orbitals) can accommodate only two electrons. Molecules such as NO, NO 2, and ClO 2 require a more sophisticated treatment of bonding, which will be developed in Chapter 9 "Molecular Geometry and Covalent Bonding Models". With 5 + 6 = 11 valence electrons, there is no way to draw a Lewis structure that gives each atom an octet of electrons. Some important examples are nitric oxide (NO), whose biochemical importance was described in earlier chapters nitrogen dioxide (NO 2), an oxidizing agent in rocket propulsion and chlorine dioxide (ClO 2), which is used in water purification plants. There are, however, a few molecules containing only p-block elements that have an odd number of electrons. Bonding in these compounds will be discussed in Chapter 23 "The ". Molecules or ions containing d-block elements frequently contain an odd number of electrons, and their bonding cannot adequately be described using the simple approach we have developed so far. Because most molecules or ions that consist of s- and p-block elements contain even numbers of electrons, their bonding can be described using a model that assigns every electron to either a bonding pair or a lone pair. ![]()
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