Covalent bonds concept

It was based on these uniqueness and commonalities, my colleague and I submitted a paper entitled Crystal Structure and A Unique martensitic Transition of TiNi to a Journal concerned with metals and alloys for publication in 1965. But, the paper was rejected outright by two anonymous reviewers who could not accept our observation that the Nitinol transition was unique. Obviously the reviews contend that by accepting Nitinol transition being unique, may make all other martensitic transformations garden variety. This may upset the theory of martensitic transition formulated thus far. We then, submitted the paper to the Journal of Applied Physics and was accepted for publication and eventually appeared in print [10]. A few months after the appearance of this article, the editor of the very journal that rejected my paper, asked me to review two papers on Nitinol for the journal. Suddenly, I was an undisputed expert in Nitinol Up to this point I had not really start to apply covalent-bond concept but devoting more time in collecting experimental data [14,15], which may be important in support or non-support of covalent-bond concept. 

Bond - Coordinate Bond . Maybe it doesn t exactly roll off the tongue, but it s hard to avoid this adaptation of the personal introduction used by perhaps our best-known and most enduring screen spy to introduce this section - it serves its purpose to remind us of the endurance and strength of bonds between metals and ligands, which at a basic level we can consider as a covalent bond. Moreover, it isn t just any bond, but a specially-constructed coordinate bond - hence the name of this field, coordination chemistry. Unfortunately, the simple covalent bonding concept does not provide valid interpretations for all of the physical properties of coordination complexes, and more sophisticated theories are required. We shall examine a number of bonding models for coordination complexes in this chapter. 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

The valence-bond concept of orbital hybridization described in the previous four sections is not limited to carbon compounds. Covalent bonds formed by-other elements can also be described using hybrid orbitals. Look, for instance, at the nitrogen atom in methylamine, CH3NH2, an organic derivative of ammonia (NH3) and the substance responsible for the odor of rotting fish. 

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. 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

Because nonmetals do not form monatomic cations, the nature of bonds between atoms of nonmetals puzzled scientists until 1916, when Lewis published his explanation. With brilliant insight, and before anyone knew about quantum mechanics or orbitals, Lewis proposed that a covalent bond is a pair of electrons shared between two atoms (3). The rest of this chapter and the next develop Lewis s vision of the covalent bond. In this chapter, we consider the types, numbers, and properties of bonds that can be formed by sharing pairs of electrons. In Chapter 3, we revisit Lewis s concept and see how to understand it in terms of orbitals. 

The concept of silicates as inorganic polymers was implicit in the ideas developed by W. H. Zacheriasen in the early 1930s. He conceived of silicates as consisting of macromolecular structures held together by covalent bonds but including network-dwelling cations. These cations were not assumed to have a structural role but merely to be present in order to balance the charges on the anionic polymer network. 

These representations are intended to mean that each C-C bond comprises three electrons. This in itself seems not such a difficult concept, but since students are usually indoctrinated to the idea that a covalent bond is two shared electrons, the notion of IV2 bonds between each pair of C atoms may be bewildering for some. 

Scarcely had the covalent chain concept of the structure of high polymers found root when theoretical chemists began to invade the field. In 1930 Kuhn o published the first application of the methods of statistics to a polymer problem he derived formulas expressing the molecular weight distribution in degraded cellulose on the assumption that splitting of interunit bonds occurs at random. 

Frequently, directionality is a property attributed to the covalent bond which supposedly is taken to be the cause of the resulting structures. However, as the success of the valence electron pair repulsion theory shows, there exists no need to assume any orbitals directed a priori. The concept of directed orbitals is based on calculations in which hybridization is used as a mathematical aid. The popular use of hybridization models occasionally has created the false impression that hybridization is some kind of process occurring prior to bond formation and committing stereochemistry. 

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