MOLECULES and The Chemical Bond
by
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Sometimes the hardest thing to see is what stares us in the face. Consider the experience that two things cannot be at the same place at the same time. Since all things are made of molecules, and nothing else, one infers that two molecules cannot be at the same place at the same time. And since molecules are composed of very small nuclei and wave-like (i.e., space-filling) electrons, and nothing else, one infers that there’s a limit to how many electrons can be at the same place at the same time. Since the formula of the simplest gas, hydrogen, whose atoms contain a single electron, is H2 (not, e.g., H3 or H4), that limit on the number of electrons that can be at the same place at the same time appears to be two. The simplest model of impenetrability-producing electron pairs is a sphere. Packing of spherical domains about sites of atomic cores yields immediately the directional and saturation character of chemical affinity, molecular shapes, profiles of molecular electron density distributions, and, accordingly, locations of molecules’ nucleophilic and electrophilic sites and, consequently, likely routes of chemical reactions.
“Impenetrability” implies “size”. The two concepts are two sides of the same coin. Atomic cores of different sizes account for occurrence of two kinds of elements, metals (large cores) and nonmetals (small cores), and three types of chemical bonds: covalent (all cores small), ionic (some cores small, some large), and metallic (all cores large). The remainder of the story is episodes: illustrative applications of the cited themes.
Dedication. For students, teachers, and professionals in the pure and applied sciences who might welcome an account of molecular structure that, in Einstein’s words, is as simple as possible but [it’s believed] no simpler and that provides, thereby, in Gibbs’ words, a point of view from which the subject appears in its greatest simplicity.
Introduction. MCB is several books interlaced. It is an introduction to atomic theory, for newcomers to chemical thought; an historical account of the development of classical structural theory, without the use of atomic orbitals (whose origin in quantum physics is shrouded in mathematical complexities); a step-by-step guide for drawing scientifically sound bond diagrams that, through use of rabbit-eared valence strokes for lone pairs and dative bonds, satisfy a powerful Valence Stroke Termination Rule; an introduction to an exclusive orbital model of bonding that embraces from one point of view covalent, ionic, and metallic bonds and that enables one to see, almost magically, at a glance, the saturation and directional character of chemical affinity, molecules’ electron density profiles, and, consequently, their frontier orbitals and manner of participation in curly-arrow reaction mechanisms. Included are special topics simply explained of paradoxical molecular structures, for connoisseurs of the chemical bond; domestication for easy use in valence theory of fundamental principles of quantum physics; and synoptic views for beginners at any level of pH calculations, the Periodic System, and the First and Second Laws of Thermodynamics. The book is, in short, a short textbook of general chemistry, in a new key.
From H2O and CO2 to Tetravalent Carbon. Consider chemistry’s best-known chemical formula: H2O. (Generalizations — i.e. inductions — often begin with the simplest, most transparent particular cases.) If both hydrogen atoms of a water molecule are “combined” in some manner with the oxygen atom, and if hydrogen is taken as having a “combining capacity” or, for short (2 syllables instead of 7), a “valence” of 1, then one might say that oxygen in HOH has a valence of 2 and carbon of CO2, if both oxygen atoms are combined with the carbon atom, a valence of 4.
“Hypervalence”. Sulfur and oxygen are congeners in periodic tables. The formula H2O suggests, therefore, the formula H2S, well known to chemists. The formula SF6, a highly stable but less well-known molecule, together with the nonexistence of OF6, has suggested to some chemists that sulfur in SF6 is, in some sense, “hypervalent”.
Naming a phenomenon does not explain it. Moreover, the name “hypervalent” is a red flag. Like the terms “electron deficient” and “antiaromatic” (to be criticized, later), “hypervalent” suggests that there is something abnormal or unnatural about sulfur in SF6. According to The First Law of Natural Science, however, Nature is always natural.
From SF6 to Rabbit Eared Valence Strokes. Sometimes hypotheticals are interesting. Suppose that before chemists knew of H2S, SF6 was well known, together with SO3, SO2F2, SOF4, and good old sulfuric acid, “H-Two-S-O-Four”, as (HO)2SO2. (The group OH is univalent, to judge from the formula HOH and the existence of HOF.) Sulfur’s standard valence would have been 6. Then, in this revision of history, H2S and SF2 arrive on the scene. Would chemists have considered sulfur in those compounds to be “hypovalent”? Historically speaking, they haven’t done so. Is there a way to look at S in SF6 and SF2 that makes the terms “hypervalent”, or “hypovalent”, unnecessary?
One obtains SF2 from SF6 by mentally stripping off 4 fluorine atoms from SF6.
One’s left on the right with 4 dangling affinities.
But, as mentioned, SF2 is a stable molecule!
To save appearances, assume that –
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About the Book
MOLECULES AND THE CHEMICAL BOND & Other Leading Chemical Concepts Simplified This highly original book by a noted chemist and chemical educator may change the way newcomers to chemical thought learn and the way its connoisseurs think about - • Atomic Theory • The Mole Concept and Avogadro’s Constant • The Gas Laws • Solving Problems in Chemical Stoichiometry • The Saturation and Directional Character of Chemical Affinity • The Pauli Exclusion Principle • Linnett’s Double Spin Set Theory • Pauling’s Rules of Crystal Chemistry • The Octet Rule • Lewis Structures for O2, NO, CO, SO2 and SO3 • Construction of Bond Diagrams • VSEPR Theory • Dative Bonding • Multicenter Bonding • Bonding in Metals • pH Calculations • The Periodic Table • The Energy Function and the First Law of Thermodynamics • The Entropy Function and the Second Law of Thermodynamics • How an Inductive Science Advances Dedicated to students, teachers, and professionals in the pure and applied sciences who might welcome an account of molecular structure that, in Einstein's words is as simple as possible but [it's believed] no simpler and that provides, thereby, in Gibbs' words, a point of view from which the subject appears in its greatest simplicity, MCB is several books interlaced. It is a novel account of evidence for atoms; an historical account of the development classical structural theory of molecules; a simple, step-by-step guide on how to draw scientifically sound bond diagrams; an exclusive orbital model of bonding that embraces from one point of view covalent, ionic, and metallic bonding; philosophical justifications for uses of molecular models; explanations for a number of previously unexplained molecular features; domestication for easy use in valence theory of fundamental principles of quantum physics; and, withal, a short textbook of general chemistry in a new key. Principally MCB is a highly visual account of a chemical mechanics of the Pauli Exclusion Principle, in the form of the story of a stroke, a stick, and a sphere and what happens if one takes chemists' seemingly unsophisticated cartoons of molecules and their corresponding tinker-toy-like ball-and-stick models seriously. One theme runs through the book: the nature of the inductive sciences, illustrated by the union of facts and ideas with creation of concepts and models, principles and rules that, jointly, comprise what is called in MCB "Conceptual Valence Bond Theory". The book has been described "as a pedagogical hierarchy of progressively more sophisticated treatments of an easily visualizable model of the chemical bond." In the words of the author's daughter (a chemist) - "This book is the culmination of my Father’s insights into the molecules he has literally breathed, consumed, and digested, for the past 84 years. It is his intimate knowledge about the elements, learned from a lifetime of reading, experimenting, and teaching that makes this book different. Dad truly loves (and believes in!) molecules, and that single tenet comes across on every page. Flat valence stroke diagrams are inflated to three dimensional valence sphere models whose geometries correlate with calculations and provide, with ease, explanations for reaction mechanisms, multicenter bonds, and molecular geometries considered “exceptions” or “unexpected”. Describing molecules as “hypervalent” or “electron deficient” suggests something abnormal or unnatural, and is misleading since “nature is always natural”. Concepts such as the gas laws and the energy function are presented from an historical perspective, and with algebraic rigor, eliminating inconsistencies that bug you as a chemistry student, but you can’t really put your finger on why. From the correct placement of helium above beryllium in the periodic table, to pointing out the problems with omitting nucleus-electron attractions in the popular Valence Shell Electron Pair Repulsion theory (where “correct conclusions regarding molecular shapes support an incorrect conceptual model”), this book challenges many of the generally accepted models taught in general chemistry, shaking up the establishment, and demanding rigor in the presentation of the subject." Libby likes specifically - • The warm-up, regarding evidence for atoms in everyday phenomena • The chemical pun: mole = mol e (with mol ≈ 600 sextillion) • The name “population” (of entities) for “n” (IUPAC's "amount") • The relations NA(vogadro) = N/n = 1/e = 12/doz. = mol/mole ≈ 600 sextillion/mole • Capture of the periodic table without the use of atomic orbitals • The Mule/Mare/Jackass Model of the Superposition Principle • The simple account of chemical affinity’s saturation and directional character • The tetrahedral model of multiple bonds • The rabbit-ear representation of lone pairs and dative bonds • The Valence Stroke Termination Rule • The Analogy between structural theories of organic and inorganic chemistry • Novel uses of Linnett’s Double Spin-Set Theory • The critique of VSEPR Theory • The model of lone pairs in the valence shells of large atomic cores • Induction of the hard-electron-pair model from matter’s impenetrability • The atomic-orbital-free account of the electronic structure of matter • The geometrical explanation of the Octet Rule, in terms of “rattling” • Discussion of the structure of the methane cation • The ease with which Valence Sphere Models handle multicenter bonding • The electride ion model of metals • Why some gaseous alkaline earth metal halides are nonlinear • The explanation for the nonlinearity of Ph3P=C=PPh3 • Details of hydrogen bonding to molecules’ nucleophilic sites • The essay on The Nature of the Double Bond • Connections between valence bond theory and Quantum Physics • The step-by-step construction of the Energy and Entropy Functions • The author’s “chemical haiku” style of writing The book is based on what is sometimes the hardest thing to see: what stares us in the face. Consider the experience that two things cannot be at the same place at the same time. Since all things are made of molecules (which may be monatomic), and nothing else, one infers that two molecules cannot be at the same place at the same time. And since molecules are composed of very small nuclei and wave-like (i.e., space-filling) electrons, and nothing else, one infers that there’s a limit to how many electrons can be at the same place at the same time. Since the formula of the simplest gas, hydrogen, whose atoms contain a single electron, is H2 (not, e.g., H3 or H4), that limit on the number of electrons that can be at the same place at the same time appears to be two. The simplest model of impenetrability-producing electron pairs is a sphere. Packing of spherical domains about sites of atomic cores yields immediately the directional and saturation character of chemical affinity, molecular shapes, approximate profiles of molecular electron density distributions, and, accordingly, locations of molecules’ nucleophilic and electrophilic sites, and, consequently, the nature of intra- and inter-molecular electron pair donor-acceptor interactions. The occurrence of atomic cores of two sizes, large and small, gives rise to two type of elements, metals and nonmetals, and three types of chemical bonds: covalent (all electron-sharing cores small), ionic (some cores small, some large); and metallic (all cores large). In "Molecules and the Chemical Bond" the author takes readers on leisurely walks in the big garden that is valence theory, so that they may pick for themselves bouquets to their liking.
About the Author
Following service in the U.S Navy during WWII, an Oberlin A.B., and a Berkeley Ph.D. in physical chemistry, the author held faculty positions at the Universities of Connecticut, Minnesota, Pittsburgh, and North Carolina State University. Since the 1960s he has been the world’s principal investigator of exclusive orbital models of molecules. A recipient of several of his nation’s leading awards in chemical education, he has serve as a chair of the American Society’s Division of Chemical Education, its Committee of Professional Training, and its Presidential Committee on Periodic Table Column Labels. He is a fellow of the American Association for the Advancement of Science and a member of a family of chemists.