Lewis theory noble gases

Clearly the explanation of the chemical bond given by Kossel cannot apply to homonuclear molecules such as CI2. Almost simultaneously with the publication of Kossel s theory, Lewis published a theory that could account for such molecules. Like Kossel, Lewis was impressed with the lack of reactivity of the noble gases. But he was also impressed by the observation that the vast majority of molecules have an even number of electrons, which led him to suggest that in molecules, electrons are usually present in pairs. In particular, he proposed that in a molecule such as CI2 the two atoms are held together by sharing a pair of electrons because in this way each atom can obtain a noble gas electron arrangement, as in the following examples ... 

Almost simultaneous with the publication of Kossel s paper there appeared a rival electronic theory. The American chemist Lewis introduced the idea of the covalent electron-pair bond. Like Kossel, he was impressed by the apparent stability of the noble gas configuration. He was also impressed by the fact that, apart from many compounds of the transition elements, most compounds when rendered as molecules have even numbers of electrons, suggesting that electrons are usually found in pairs. Lewis devised the familiar representations of molecules and polyatomic ions (Lewis structures, or Lewis diagrams) in which electrons are shown as dots (or as noughts and crosses) to show how atoms can attain noble gas configurations by the sharing of electrons in pairs, as opposed to complete transfer as in Kossel s theory. It was soon apparent from the earliest X-ray studies that Kossel s theory was more appropriate... 

In 1915, G. N. Lewis proposed several new theories describing how atoms bond together to form molecules. One of these theories states that a filled shell of electrons is especially stable, and atoms transfer or share electrons in such a way as to attain a filled shell of electrons. A filled shell of electrons is simply the electron configuration of a noble gas, such as He, Ne, or Ar. This principle has come to be called the octet rule... 

The electrons in the outer shell, or valence shell, of an atom are the electrons involved in bonding. In most of our discussion of covalent bonding, we will focus attention on these electrons. Valence shell electrons are those that were not present in the preceding noble gas, ignoring filled sets of d and / orbitals. Lewis formulas show the number of valence shell electrons in a polyatomic molecule or ion (Sections 7-4 through 7-7). We will write Lewis formulas for each molecule or polyatomic ion we discuss. The theories introduced in this chapter apply equally well to polyatomic molecules and to ions. 

Lewis assumed that the pair of electrons was one of the stable arrangements, so that apart from the formation of noble gas electron arrangements, atoms having unpaired electrons would form bonds. This he calls the rule of two, to distinguish it from the rule of eight. Both rules are explicit in his theory. 

Atoms can also achieve noble-gas electron arrangements by the transfer of electrons from one atom to another. This process is also an integral part of Lewis theory. The process produces opposite... 

The first quantitative theory of chemical bonding was developed for the hydrogen molecule by Heitler and London in 1927, and was based on the Lewis theory of valence in which two atoms shared electrons in such a way that each achieved a noble gas structure. The theory was later extended to other, more complex molecules, and became known as valence bond theory. In this approach, the overlap of atomic orbitals on neighbouring atoms is considered to lead to the formation of localized bonds, each of which can accommodate two electrons with paired spins. The theory has been responsible for introducing such important concepts as hybridization and resonance into the theory of the chemical bond, but applications of the theory have been limited by difficulties in generating computer programs that can deal efficiently with anything other than the simplest of molecules. 

Lewis no doubt derived some of his ideas from Parson, but then ideas are rarely conceived in a vacuum. A month before Lewis s publication, Walther Kossel of Germany had published a paper that assumed that atoms gained or lost electrons to achieve the same number of electrons as a noble gas atom, but this work was unknown to Lewis while he was preparing his manuscript. And in Lewis s hands the theory did more than explain ions or ionic bond formation. It became a strong rationalization for the nonpolar bond. 

Lewis theory was a pre-quantum theory and based on the simple idea that stable electron structures were those where the atoms could get a noble gas configuration by sharing or transferring electrons. The ideas of covalent and ionic bonds came from this model. Resonance structures were introduced because no unique bonds could be defined for some molecules, for example, benzene and some oxyanions, such as sulfate(vi) and nitrate(v), and were therefore an extension of Lewis ideas to such cases. 

Lewis summarized much of his theory of chemical bonding with the octet rule. According to the octet rule, atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration. This rule enables us to predict many of the formulas for compounds consisting of specific elements. The octet rule holds for nearly all the compounds made up of second period elements and is therefore especially important in the smdy of organic compounds, which contain mostly C, N, and O atoms. 

The hydrides of Groups IV-VII consist of EH. molecules, in which x is the number of electrons needed to bring the electronic configuration of E up to that of the subsequent noble gas. This is easily explained by simple Lewis theory the EH. molecules contain E—H single bonds, and the x bonds add jc electrons to the outer electronic configuration of E to give it a noble gas configuration. 

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