Electrons, Bonds, and Lewis Structures

As mentioned, atoms are connected to each other by bonds. That is, bonds are the glue that hold atoms together. But what is this mysterious glue and how does it work In order to answer this question, we must focus our attention on electrons. 

The existence of the electron was first proposed in 1874 by George Johnstone Stoney (National University of Ireland), who attempted to explain electrochemistry by suggesting the existence of a particle bearing a unit of charge. Stoney coined the term electron to describe this particle. In 1897, J. J. Thomson (Cambridge University) demonstrated evidence supporting the existence of Stoney s mysterious electron and is credited with discovering the electron. In 1916, 

Gilbert Lewis (University of California, Berkeley) defined a covalent bond as the result of two atoms sharing a pair of electrons. As a simple example, consider the formation of a bond between two hydrogen atoms  

An energy diagram showing the total energy as a function of the internuclear distance between two hydrogen atoms. 

Armed with the idea that a bond represents a pair of shared electrons, Lewis then devised a method for drawing structures. In his drawings, called Lewis structures, the electrons take center stage. We will begin by drawing individual atoms, and then we will draw Lewis structures for small molecules. First, we must review a few simple features of atomic structure  

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Electrons, Bonds, and Lewis Structures (Section 1.3) Molecular Orbital Theory (Section 1.8)... 

Valence-bond theory, 32—34, 42, 46 Valence electrons, 10 and Lewis structures, 20 Valence-shell electron pair repulsion and molecular geometry, 26-29, 45 L-Valine, 1054, 1059... 

By the 1920s, vitalism had been discarded. Chemists were aware of constitutional isomerism and had developed the structural theory of matter. The electron had been discovered and identified as the source of bonding, and Lewis structures were used to keep track of shared and unshared electrons. But the understanding of electrons was about to change dramatically. 

Bond Energy and Enthalpy The Localized Electron Bonding Model Lewis Structures Exceptions to the Octet Rule Resonance... 

Age-old questions concerning the nature of the bonds between atoms in molecules culminated in the remarkable Lewis structure model of G. N. Lewis (1916). The notion that such bonds were formed from directed hybrids was subsequently developed by Linus Pauling (1932), shortly after the discovery of quantum mechanics. Although many theoretical advances have ensued, it is fair to say that the underlying concepts of valence-shell hybridization, shared-electron pair bonds, and Lewis structural dot diagrams continue to dominate chemical thinking and pedagogy to this day. 

Sixteen electrons are required to give each atom an octet (four for S and six for each O). There is a deficiency of two electrons. This means that a single bond in the skeleton must be converted to a double bond. The Lewis structure of S02 is... 

Now consider the alkynes, hydrocarbons with carbon-carbon triple bonds. The Lewis structure of the linear molecule ethyne (acetylene) is H—O C- H. To describe the bonding in a linear molecule, we need a hybridization scheme that produces two equivalent orbitals at 180° from each other this is sp hybridization. Each C atom has one electron in each of its two sp hybrid orbitals and one electron in each of its two perpendicular unhybridized 2p-orbitals (43). The electrons in the sp hybrid orbitals on the two carbon atoms pair and form a carbon—carbon tr-bond. The electrons in the remaining sp hybrid orbitals pair with hydrogen Ls-elec-trons to form two carbon—hydrogen o-bonds. The electrons in the two perpendicular sets of 2/z-orbitals pair with a side-by-side overlap, forming two ir-honds at 90° to each other. As in the N2 molecule, the electron density in the o-bonds forms a cylinder about the C—C bond axis. The resulting bonding pattern is shown in Fig. 3.23. 

C09-0138. In the following reactions, phosphorus forms a bond to a Row 2 element. In one reaction, phosphoms donates two electrons to make the fourth bond, but in the other reaction, phosphorus accepts two electrons to make the fourth bond. Use Lewis structures of starting materials and products to determine in which reaction phosphoms is a donor and in which it acts as an acceptor. 

Main-group organometallic compounds are versatile tools in organic synthesis, but their structures are complicated by the involvement of the multicenter, two-electron bonds and ion-dipole interactions that are involved in aggregate formation (5). Electron deficiency or Lewis acidity of the metallic center and nucleophilicity or basicity of the substituents are important considerations in synthesis. The complexity of the structures and interactions is, however, the origin of much of the unique behavior of these organometallic compounds. 

IN the past twenty years the electronic structures of many organic molecules, particularly benzene and related compounds, have been discussed in toms of the molecular orbital and valence bond methods.1 During the same period the structures of inorganic ions have been inferred from the bond distances f a bond distance shorter than the sum of the conventional radii has been attributed to the resonance of double bonded structures with the single bonded or Lewis structure. 

For many simple compounds having no more than one double bond, the modern picture may be quite adequately represented by the Lewis structures (although the Lewis rules are noncommittal about the shapes of molecules). For compounds such as butadiene, benzene, and nitrous oxide, where there is extensive delocalization of electron density, the Lewis structures are not as suitable as the x-electron structures or, better still, as the streamer structures. Both of the latter type, however, are more difficult to draw and, for more complex molecules, more difficult to visualize they become extremely unwieldy when one attempts to use them to represent the progress of a chemical reaction. 

A Lewis structure cannot be written for NO, for it has an odd number of electrons. Pauling interprets its structure as having a three-electron bond and estimates that the extra electron adds about half as much extra stability to the molecule as would an ordinary covalent bond. Alternatively, the molecular orbital picture describes this molecule as having one a bond and f ir bonds (p. 71). At any rate, nitric oxide is, as it should be, paramagnetic. 

The energy level diagram shown in Fig. 13 along with the six electrons from the Lewis structure shown earlier shows that the lower three BMO s are filled. These MO s are four-center two-electron bonds and are therefore delocalized over the whole molecule. 

The formalism of connectivity indices is an embodiment of graph theory. Connectivity indices are intuitively appealing because each index can be calculated exactly from valence bond (Lewis) diagrams familiar to organic chemists, which depict molecular structure in terms of atoms, inner shell and valence shell electrons, valence shell hybridization, o and k electrons, bonds and lone pairs. The indices can then be correlated with physical or chemical properties of interest. Connectivity indices have, in the past, been very useful in treating molecular systems with well-defined chemical formulae and fixed numbers of atoms [24,25]. 

The octet rule predicts that atoms form enough covalent bonds to surround themselves with eight elechons each. When one atom in a covalently bonded pair donates two electrons to the bond, the Lewis structure can include the formal charge on each atom as a means of keeping track of the valence electrons. There are exceptions to the octet rule, particularly for covalent beryllium compounds, elements in Group 3A, and elements in the third period and beyond in the periodic table. 

Three-centre bonds are also found in compounds, particularly of sulfur, the halogens and the noble gases in which there is formally an expansion of the octet (that is, the central atom is surrounded by more than eight electrons in the Lewis structure). 

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