Phosphorus valence

The chief differences between nitrogen and phosphorus arise from the availability of d orbitals in the shell with principal quantum number 3. This gives phosphorus valence properties not shared by nitrogen for example, a capacity for 6 co-ordination associated with octahedral hybridisation (p. 105). 

Chlorine substituents at the phosphorus atom cause relatively small changes of the PO bond lengths. Likewise small is the corresponding elongation of the PCI bonds as compared to the elongation which occurs in amino derivatives (Table 17). Changes in the number of chlorine atoms bonded to phosphorus and in the phosphorus valency state influence the PCI bond length more markedly. 

Phosphorus is in group 5A of the periodic table and has five valence elections. It thus needs to share three more electrons to make an octet and therefore bonds to three hydrogen atoms, giving PH3. 

The largest class of molecules to violate the octet rule consists of species in which the central atom is surrounded by more than four pairs of valence electrons. Typical molecules of this type are phosphorus pentachloride, PC15, and sulfur hexafluoride, SF6. The Lewis structures of these molecules are... 

As you can see, the central atoms in these molecules have expanded octets. In PC15, the phosphorus atom is surrounded by 10 valence electrons (5 shared pairs) in SF6, there are 12 valence electrons (6 shared pairs) around the sulfur atom. 

The extra electron pairs in an expanded octet are accommodated by using d orbitals. The phosphorus atom (five valence electrons) in PC15 and the sulfur atom (six valence electrons) in SF6 make use of 3d as well as 3s and 3p orbitals ... 

How many valence electrons has carbon Silicon Phosphorus Hydrogen Write the electron configurations for neutral atoms of each element. 

The electron configuration in the valence orbitals of the sulfur atom (3s 3p4) suggests that it will form two covalent bonds by making use of two half-filled 3p orbitals. This is, in fact, observed in the molecule S8, which is present in the common forms of solid sulfur. The S8 molecules assume the form of a puckered ring, as shown in Figure 20-3. As with the phosphorus, the stability of this crystalline form of sulfur is due to van der Waals forces between discrete molecules. 

This problem clearly did not worry Stoner, who just went ahead and assumed that three quantum numbers could be specified in many-electron atoms. In any case, Stoner s scheme solved certain problems present in Bohr s configurations. For example, Bohr had assigned phosphorus the configuration 2,4,4,41, but this failed to explain the fact that phosphorus shows valencies of three and five. Stoner s configuration for phosphorus was 2,2,2,4,2,2,1, which easily explains the valencies, since it becomes plausible that either the two or the three outermost subshells of electrons form bonds. 

The octet rule accounts for the valences of many of the elements and the structures of many compounds. Carbon, nitrogen, oxygen, and fluorine obey the octet rule rigorously, provided there are enough electrons to go around. However, some compounds have an odd number of electrons. In addition, an atom of phosphorus, sulfur, chlorine, or another nonmetal in Period 3 and subsequent periods can accommodate more than eight electrons in its valence shell. The following two sections show how to recognize exceptions to the octet rule. 

Elements that can expand their valence shells commonly show variable covalence, the ability to form different numbers of covalent bonds. Elements that have variable covalence can form one number of bonds in some compounds and a different number in others. Phosphorus is an example. It reacts directly with a limited supply of chlorine to form the toxic, colorless liquid phosphorus trichloride ... 

In the vapor phase, phosphorus can exist as P2 molecules, which are highly reactive, whereas N2 is relatively inert. Use valence-bond theory to explain this difference. 

The radius of an atom helps to determine how many other atoms can bond to it. The small radii of Period 2 atoms, for instance, are largely responsible for the differences between their properties and those of their congeners. As described in Section 2.10, one reason that small atoms typically have low valences is that so few other atoms can pack around them. Nitrogen, for instance, never forms penta-halides, but phosphorus does. With few exceptions, only Period 2 elements form multiple bonds with themselves or other elements in the same period, because only they are small enough for their p-orbitals to have substantial tt overlap (Fig. 14.6). 

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