Classical valence rules

Despite the completely different approach to chemical interaction, which has been followed here, the conventional standard symbols which are used to define the connectivity in covalent molecules, can also be applied, without modification, to distinguish between interactions of different order. However, each linkage pictured by formulae such as H3C-CH3, H2C=CH2, HC=CH, represents the sharing of a single pair of electrons with location unspecified. The number of connecting fines only counts bond order and may be established from the classical valence rules, e.g. v(C,N,0,F)=(4,3,2,l). Special symbols are used for non-integral bond orders, as in the symbol for benzene ... 

In a sense, Kekuld s concept of molecular compounds" was a revival of Berzelius s dualisdc theory whereby secondary compounds" (in Kekuld s terminology, atomic compounds") containing a small excess of electrical charge could still combine with other secondary compounds containing a small excess of opposite charge to form tertiary compounds" (in Kekuld s terminology, molecular compounds"). At most, Kekuld s artificial division of compounds into atomic compounds," which obeyed the rules of classical valence theory, and into molecular compounds," which did not obey these rules, had some limited value as a formal classification. However, in no way did it explain the nature or operation of the forces involved in the formation of molecular... 

Aromaticity is no doubt one of the most interesting features of small boron clusters. Possessing planar structures, these boron clusters were characterized as highly aromatic systems, on the basis of various indices such as nucleus-independent chemical shift (NICS), electron localization function (ELF), resonance energy (RE), the presence of ring current induced by an external magnetic field, and the classical Hilckel rule of (4IV -]- 2) valence electrons, etc... [8-10, 27]. In recent reports [20, 26], we found that all small boron clusters B with n <20 have an aromatic character, irrespective of their numbers of valence electrons. 

In 1923. Lewis published a classic book (later reprinted by Dover Publications) titled Valence and the Structure of Atoms and Molecules. Here, in Lewis s characteristically lucid style, we find many of the basic principles of covalent bonding discussed in this chapter. Included are electron-dot structures, the octet rule, and the concept of electronegativity. Here too is the Lewis definition of acids and bases (Chapter 15). That same year, Lewis published with Merle Randall a text called Thermodynamics and the Free Energy of Chemical Substances. Today, a revised edition of that text is still used in graduate courses in chemistry. 

If one organic compound has dominated the historical literature of the last few years, that compound must be benzene. Most probably, this is because its structure in some respects marks a transition from the most austere form of classical organic chemistry, in which carbon was tetravalent and tetrahedral, to a continuing series of changes from oscillating molecules, through partial valencies to MO descriptions, and Huckel s rules of aromaticity. It is the case par excellence of a single substance whose history intersects all major streams of chemical theory - except perhaps the periodic law - and which also has enormous industrial and economic importance. 

The most definite and indisputable evidence for the existence of free radicals is obtained from spectroscopy. Physicists are able to interpret band spectra without ambiguity on the basis of such units as OH, CN, BeCl, SiO, CH2, C2 and others. There is a whole host of radicals of this type which can explain quantitatively all the lines of a complex band spectrum and there is no other way to explain them. Moreover calculations based on quantum mechanics show that many of these free radicals, which violate all rules of the classical theories of valence, are stable and do not necessarily decompose at ordinary or even at moderately high temperatures. The difficulty in finding them is not that they are too unstable but rather that they are so very reactive that they combine immedi-... 

The classical rules of valency do not apply for complex ions. To explain the particularities of chemical bonding in complex ions, various theories have been developed. As early as 1893, A. Werner suggested that, apart from normal valencies, elements possess secondary valencies which are used when complex ions are formed. He attributed directions to these secondary valencies, and thereby could explain the existence of stereoisomers, which were prepared in great numbers at that time. Later G. N. Lewis (1916), when describing his theory of chemical bonds based on the formation of electron pairs, explained the formation of complexes by the donation of a whole electron pair by an atom of the ligand to the central atom. This so-called dative bond is sometimes denoted by an arrow, showing the direction of donation of electrons. In the structural formula of the tetramminecuprate(II) ion... 

The valence shell of electrons in an atom is the outermost shell of electrons of the uncombined atom. The electrons in that shell are called valence electrons. If all the electrons are removed from that shell, the next inner shell becomes the new outermost shell. For example, the sodium atom has 2 electrons in its first shell, 8 electrons in its second shell (the maximum), and its last electron in its third shell. The valence shell is the third shell. If the 1 electron is removed from the third shell, the second shell becomes the outeamost shell, containing 8 electrons. The valence shell is still the (now empty) third shell. The number of electrons in the valence shell of an uncombined main group atom is equal to the classical periodic group number of the element (Figure 5.7). The exceptions to this rule are that hehum has 2 valence electrons and the other noble gases have 8 valence electrons. 

For xenon fluorides and oxides, for example, the same models can be apphed as for interhalogen and halogen oxy species. Furthermore, the very successful valence shell electron pair repulsion (VSEPR) rules for molecule and ion shapes are as effective for noble gas compounds and their relatives as for classical octet compounds. 

G. N. Lewis (above) conceived the octet rule while lecturing to a class of general chemistry students in 1902. He was also one of the two authors of a now classic work on thermodynamic, Lewis and Randall, Thermodynamics and the Free Energy of Chemical Substances (1923). (right) This is his original sketch. From G. N. Lewis, Valence, Dover Publications, Inc., New York, 1966. 

We have used 12 electrons to form the S—F bonds, which leaves 36 electrons. Since fluorine always follows the octet rule, we complete the six fluorine octets to give the structure on the right above. This structure uses all 48 valence electrons for SF6, but sulfur has 12 electrons around it that is, sulfur exceeds the octet rule. How can this happen There are several ways to approach this situation. The classical explanation for molecules like SFg involves using the empty 3d orbitals on the third-period elements. Recall that the second-row elements have only 2s and 2p valence orbitals, whereas the third-row elements have 3s, 3p, and 3d orbitals. The 3s and 3p orbitals fill with electrons in going from sodium to argon, but the 3d orbitals remain empty. For example, the valence-orbital diagram for a sulfur atom is... 

Apart from a few general rules, the alloying behaviour of metals is rather empirical. The classical rules of Hume-Rothery [220] explain this behaviour reasonably well. Such factors as size, electronegativity, valency, electron concentration, free energy, formation of intermediate phases and isomorphism are found to influence the alloying tendency of metals. However, size and electronegativity are the two most important factors, and they profoundly influence the solubility of the solute atoms and greatly affect the crystal structures of the alloys. 

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