Chemical Bonds The Octet Rule

The first explanations of the nature of chemical bonds were advanced by G. N. Lewis (of the University of Califomia, Berkeley) and W. Kossel (of the University of Munich) in 1916. Two major types of chemical bonds were proposed  

Ionic (or electrovalent) bonds are formed by the transfer of one or more electrons from one atom to another to create ions. 

Covalent bonds result when atoms share electrons. 

The concepts and explanations that arise from the original propositions of Lewis and Kossel are satisfactory for explanations of many of the problems we deal with in organic chemistry today. For this reason we shall review these two types of bonds in more modem 

Atoms may gain or lose electrons and form charged particles called ions. 

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There is a well defined relationship between the chemical structure of a Zintl phase and its electronic structure. For the majority of these compounds, AX homopolar X-X contacts are present and can be explained as two-electron, two-center bonds. The octet rule [5] is fijlfiUed for the A and for the X atoms. This is provided by a formal charge transfer of the valence electrons fiom A to X leading to... 

During some chemical interactions, the octet rule is satisfied when electrons are transferred from one atom to another. As a result of the transfers, neutral atoms acquire net positive or negative electrical charges and become attracted to one another. These charged atoms are called simple ions, and the attractive force between oppositely charged atoms constitutes an ionic bond. A second type of interaction that also satisfies the octet rule is discussed in Section 4.6. 

Lewis structure (Section 1 3) A chemical formula in which electrons are represented by dots Two dots (or a line) be tween two atoms represent a covalent bond in a Lewis structure Unshared electrons are explicitly shown and sta ble Lewis structures are those in which the octet rule is sat isfied... 

Lewis s interest in chemical bonding and structure dated from 1902. In attempting to explain "valence" to a class at Harvard, he devised an atomic model to rationalize the octet rule. His model was deficient in many respects for one thing, Lewis visualized cubic atoms with electrons located at the corners. Perhaps this explains why his ideas of atomic structure were not published until 1916. In that year, Lewis conceived of the... 

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. 

Atoms of these elements have empty J-orbitals in the valence shell. Another factor—possibly the main factor—in determining whether more atoms than allowed by the octet rule can bond to a central atom is the size of that atom. A P atom is big enough for as many as six Cl atoms to fit comfortably around it, and PC15 is a common laboratory chemical. An N atom, though, is too small, and NC15 is unknown. A compound that contains an atom with more atoms attached to it than is permitted by the octet rule is called a hypcrvalent compound. This name leaves open the question of whether the additional bonds are due to valence-shell expansion or simply to the size of the central atom. 

There are several important chemical species that consist of four atoms and have a total of 24 valence-shell electrons. Some of the most common isoelectronic species of this type are C032-, N03 , S03, and P() j (known as the metaphosphate ion). Because four atoms would require a total of 32 electrons for each to have an octet, we conclude that eight electrons must be shared in four bonds. With four bonds to the central atom, there can be no unshared pairs on that atom if the octet rule is to be obeyed. Therefore, we can draw the structure for CO, 2 showing one double C=0 bond and two single C-O bonds as... 

Quantum chemists have developed considerable experience over the years in inventing new molecules by quantum chemical methods, which in some cases have been subsequently characterized by experimentalists (see, for example, Refs. 3 and 4). The general philosophy is to explore the Periodic Table and to attempt to understand the analogies between the behavior of different elements. It is known that for first row atoms chemical bonding usually follows the octet rule. In transition metals, this rule is replaced by the 18-electron rule. Upon going to lanthanides and actinides, the valence f shells are expected to play a role. In lanthanide chemistry, the 4f shell is contracted and usually does not directly participate in the chemical bonding. In actinide chemistry, on the other hand, the 5f shell is more diffuse and participates actively in the bonding. 

The octet rule is one of the cornerstones of chemical bonding theory. While the vast majority of molecules conform, apparent exceptions occur for molecules incorporating second-row (and heavier) main-group elements. Apparent refers to the fact that molecules such as dimethylsulfoxide and dimethylsulfone may either be represented in terms of structures with ten and twelve valence electrons, respectively, surrounding sulfur, or as zwitterions with the normal complement of eight valence electrons (see also discussions in Chapters 5 and 16). 

There are many exceptions to the octet rule—after all, it s called the octet rule, not the octet law—but it is nevertheless useful for making predictions and for providing insights about chemical bonding. 

The covalence of hydrogen is always one because it cannot form more than one chemical bond. The covalence of oxygen is almost always two and occasionally one. The covalence of carbon is four in almost all of its stable compounds—there may be single, double, or triple bonds involved, but the total number of bonds is four. Although the octet rule is not a rigid rule for chemical bonding, it is obeyed for C, N, O, and F in almost all their compounds. The octet is exceeded commonly for elements in the third and higher periods. 

The electronic structure of nitric oxide, NO (Figure 10.1), is not as straightforward as some compounds we have discussed earlier in this book. It appears to have an odd electron left over and does not obey the octet rule for atoms in chemical bonding. 

We suggest that the octet rule be demoted in favour of the democracy principle almost all valence electrons can participate in chemical bonding if provided with sufficient energetic incentives. Simple concepts of atomic size and of electronegativity differences prove to be of particular utility in qualitative descriptions. We find no evidence for the utilization of d functions as valence orbitals, or to support notions of p -d back-bonding. 

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