Lewis theory octet

According to Lewis s octet rule, each atom should be surrounded by four pairs of electrons, either shared or free pairs. Lewis derived stmctures for halogen molecules, the ammonium ion, and oxy acids, inexplicable according to previous valence theories. He viewed polar bonds as unequally shared electron pairs. Because the complete transfer of electrons was only an extreme case of polarity, he abandoned his earlier dualistic view the polar theory was just a special case of his more general theory. 

Recall from Chapter 5 that when nonmetals bond with other nonmetals, a molecular compoimd results. Molecular compounds contain covalent bonds in which electrons are shared between atoms rather than transferred. In Lewis theory, we represent covalent bonding by allowing neighboring atoms to share some of their valence electrons in order to attain octets (or duets for hydrogen). For example, hydrogen and oxygen have the Lewis structures ... 

In Lewis theory, atoms can share more than one electron pair to attain an octet. For example, we know from Chapter 5 that oxygen exists as the diatomic molecule, O2. The Lewis structure of an oxygen atom is ... 

The nitrogen atom does not have an octet, so this is not a great Lewis structure. However, NO exists in nature. Why As with any simple theory, Lewis theory is not sophisticated enough to be correct every time. It is impossible to write good Lewis structures for molecules with odd numbers of electrons, yet some of these molecules exist in nature. In such cases, we simply write the best Lewis structure that we can. Another significant exception to the octet rule is boron, which tends to form compounds with only six electrons around B, rather than eight. For example, BF3 and BH3—both of which exist in nature— lack an octet for B. 

These are often referred to as expanded octets. Expanded octets can form for period 3 elements and beyond. Beyond mentioning them, we do not cover expanded octets in this book. In spite of these exceptions, Lewis theory remains a powerful and simple way to understand chemical bonding. 

In Lewis theory, what is an octet What is a duet What is a chemical bond ... 

Lewis theory A simple theory for chemical bonding involving diagrams showing bonds between atoms as lines or dots. In this theory, atoms bond together to obtain stable octets (8 valence electrons). 

Magnesiiun has lost all its valence electrons and appears without any dots, while the fluorine atoms both acquire octets. When we look at magnesium fluoride in nature, we indeed find that it is composed of two fluoride ions to every magne-siiun ion, just as Lewis theory predicts. 

In covalent bonds, atoms share their electrons. We represent covalent bonding in Lewis theory by letting atoms share their dots such that some dots count for the octet of more than one atom. For example, we learned in Chapter 3 that elemental chlorine exists as the diatomic (two atom) molecule CI2. Lewis theory explains why. Consider the Lewis structure of chlorine ... 

Because Lewis theory requires all atoms to have octets, we might expect all molecules to be tetrahedral. However, because central atoms often have lone pairs as well as bonding pairs, and because doubie and tripie bonds force electrons to group together, a number of different geometries are aiso possible. For example, the Lewis structure for water has four eiectron groups on the central atom, two bonding pairs and two lone pairs ... 

We represent chemical bonding with Lewis theory. In this model, valence electrons are represented as dots surrounding the chemical symbol for the element. Compounds are formed by allowing the electrons to be transferred between atoms (ionic bonding) or shared between atoms (covalent bonding) so that all atoms acquire an octet. 

Many chemical models are known to be incorrect, for example, the Lewis theory of bonding, or an incomplete description of chemical phenomena, for example, the octet rule (Chapter 4). For the purposes of teaching chemistry it is often preferable to use simple approximate bonding models that give correct predictions for the majority of cases, than to use a more accurate but complicated model, such as the quantum mechanical model (Chapter 12), which is based on the electron s wave properties. 

For a number of electron-excess molecules that involve atoms of first-row elements (in particular, nitrogen atoms), an older type of increased-valence structure is sometimes used to represent their electronic structures. Since the 1860s and until the introduction of the Lewis-Langmuir octet theory, nitrogen atoms were often represented with valencies of 3 or 5 in valence-bond structures. For example, valence-bond structures for N2O, MCjNO, and N2O4 were written as structures (l)-(3). 

Using structure (1) for N2O, with an N-N triple bond and an N-0 double bond, we would be let to predict that the bond-lengths are similar to the 1.10 A and 1.21 A for N2 and CH3N=0. The experimental lengths of 1.13 A and 1.19 A confirm this expectation (Section 2-3 (b)), and on this basis structure (1) is a suitable valence-bond structure for N2O. But the Lewis theory, with electron-pair bonds, does not permit a valence of five for first-row atoms, provided that only the 2s and three 2p orbitals of these atoms are valence orbitals for bonding. Therefore, in the Lewis theoiy, the quinquevalent structures are replaced by octet stractures such as structures (4)-(6) for N2O, and (7) and (8) (together with equivalent resonance forms) for MesNO and N2O4. 

Lewis Theory An Overview—Lewis symbol represents the valence electrons of an atom by using dots placed around the chemical symbol. A Lewis structure is a combination of Lewis symbols used to represent chemical bonding. Normally, all the electrons in a Lewis structure are paired, and each atom in the structure acquires an octet—that is, there are eight electrons in the valence shell. In Lewis theory, chemical bonds are classified as ionic bonds, which are formed by electron transfer between atoms, or covalent bonds, which are formed by electrons shared between atoms. Most bonds, however, have partial ionic and partial covalent characteristics. 

The simplest hydrocarbon observed imder normal laboratory conditions is methane, CH4. This is a stable, unreactive molecule with a molecular formula consistent with the octet rule of the Lewis theory. To obtain this molecular formula by the valence bond method, we need an orbital diagram for carbon in which there are four unpaired electrons so that orbital overlap leads to four C—H bonds. To get such a diagram, imagine that one of the 2s electrons in a ground-state C atom absorbs energy and is promoted to the empty 2p orbital. The resulting electron configuration is that of an excited state. 

The apparent inertness of the noble gases gave them a key position in the electronic theories of valency as developed by G. N. Lewis (1916) and W. Kossel (1916) and the attainment of a stable octet was regarded as a prime criterion for bond formation between atoms (p. 21). Their monatomic, non-polar nature makes them the most nearly perfect gases known, and has led to continuous interest in their physical properties. 

Irving Langmuir, "The Structure of Molecules," BAAS Rep. Edinburgh. 1921 (1922) 468469. G. N. Lewis. "The Atom and the Molecule," 762785 Irving Langmuir, "The Structure of Atoms and the Octet Theory of Valence." Proc. NAS 5 (1919) 252259, and JACS "The Arrangement of Electrons in Atoms and Molecules,"... 

Niels Bohr s 1913 hydrogen atom paper demonstrates the traditional interest of some physicists in placing the facts and laws of chemistry within a broader framework of foundational principles laid out by physicists. During the course of the next two decades, a number of physicists who became known as quantum physicists developed physical theories and mathematical techniques that they claimed would create a mathematical and theoretical chemistry. However, few of them had much chemical knowledge beyond a general understanding of the periodic table of the elements and familiarity with the Lewis-Langmuir theory of the electron duplet and octet. 

Draw Lewis structures for the following molecules and ions, and use VSEPR theory to predict the molecular shape. Indicate the examples in which the central atom has an expanded octet. 

G. N. Lewis s first sketch of the octet theory of valence electrons. 

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