Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction



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Reading Materials for Chemistry English

Periodic Table

Digested from



Chemical & Chemical Engineering News

(80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. 8.

Edited by X. Lu

Introduction

The Periodic Table is nature’s rosetta stone. To the uninitiated, it’s just 100-plus numbered boxes, each containing one or two letters, arranged with an odd, skewed symmetry. To chemists, however, the periodic table reveals the organizing principles of matter, which is to say, the organizing principles of chemistry. At a fundamental level, all of chemistry is contained in the periodic table.

That’s not to say, of course, that all of chemistry is obvious from the periodic table. Far from it. But the structure of the table reflects the electronic structure of the elements, and hence their chemical properties and behavior. Perhaps it would be more appropriate to say that all of chemistry starts with the periodic table. What has always struck me as remarkable about Dmitry Mendeleyev’s discovery in 1869 of a way to arrange the elements known at that time into a meaningful and predictive periodic table is that he accomplished it without any knowledge of the structure of atoms. There are no atomic numbers on Mendeleyev’s periodic table, only atomic weights and the groupings of elements based on their known properties. More than 30 years would pass before J. J. Thomson (who discovered the electron) suggested that the electronic configuration of atoms might account for the periodicity of the elements, and more than 40 years would pass before atomic numbers were recognized as the basis for ordering the elements.

Indeed, as John Emsley notes in his invaluable book, “Nature’s Building Blocks: An A–Z Guide to the Elements,” Mendeleyev never accepted that electrons came from atoms because he was convinced that atoms were indivisible.

No matter. Mendeleyev’s brilliant insight propelled chemistry into the 20th century. New elements were discovered that filled in the holes Mendeleyev left in his table, and their atomic weights and chemical properties corresponded with remarkable accuracy to Mendeleyev’s predictions. And as the revolution in chemistry and physics unfolded in the early decades of the 20th century, most of the discoveries about the structure of atoms, their properties, and how they interact with each other made perfect sense in light of the periodic table.

The periodic table is so central to chemistry that it seemed natural to devote a special issue to it and the elements that compose it as we celebrate C&EN’s 80th anniversary. The 89 essays are delightfully varied. We hope they will give you a new perspective on and appreciation of the building blocks of our science.

HYDROGEN

JUDITH KLINMAN, UNIVERSITY OF CALIFORNIA, BERKELEY




To be the number one--not to mention the most abundant and lightest--element on the periodic table is a weighty responsibility. But hydrogen does not disappoint. Its role in the universe is indisputable, but for me its attraction lies in its two long-lived isotopes: one stable, deuterium (D), and one radioactive, tritium (T). In fact, I can't imagine the emergence of modern chemistry and biochemistry without the benefit of these isotopes.

My first exploration of deuterium was as a graduate student with Edward Thornton, when we used solvent D2O to examine mechanisms of carbonyl reactivity in the condensed phase. But biology beckoned, and the application of hydrogen isotopes to biochemical processes proved irresistible. Working subsequently with Irwin Rose as a postdoctoral researcher, I used enzymes to prepare methyl groups that were chiral by virtue of the presence of H, D, and T. The key to unlocking the cryptic stereochemistry for the production and utilization of chiral methyl groups by metabolic enzymes lay with the difference in reaction rate between H and D.






FUEL FOR THOUGHT

The vast majority of matter in the universe-- and the primary fuel for stars--is hydrogen.


Theories of kinetic isotope effects had been advanced in the preceding years. These attributed isotopic differences in rate to differences in the energy of the ground-state stretching modes of the labeled bonds, which were recognized as arising solely from the altered masses of the isotopes. Eureka--with the expectation that the potential energy surface of the labeled bond would be unaffected--biochemists had found the long-awaited nonperturbing probe with which to analyze the nature of catalysis in enzymatic reactions.

A decade and a half of exciting activity followed. Suddenly, it was possible to dissect enzyme reactions into their individual steps, learning in the process how enzymes alter the reaction-barrier height together with the internal thermodynamics for the interconversion of bound reactants and products.

Around the mid-1980s, scientists began to realize that all was not as it had seemed--experimental anomalies indicated breakdowns in classical theories of kinetic hydrogen isotope effects. This was met with a healthy blend of curiosity and resistance. Both experimentalists and theorists then opened up a new vista in the chemistry and biochemistry of hydrogen.

Returning to our roots, my laboratory called upon the three isotopes of hydrogen--H, D, and T--to examine the origin of the kinetic anomalies in the hydride-transfer reaction catalyzed by alcohol dehydrogenase. From a comparison of H/T and D/T kinetic isotope effects, it was possible to demonstrate that the entire methylene carbon at the reactive position of the alcohol substrate was showing a pattern that had been predicted for hydrogen tunneling. Clearly, the wave property of the hydrogen atom was dominating its reactivity at room temperature.

Soon, an arsenal of experimental tools was developed for investigating tunneling in enzyme-catalyzed reactions. In every system studied, the quantum property of the reacting hydrogen was apparent under physiological temperatures. This extended to the functionally high temperatures of a thermophilic enzyme, where tunneling could be demonstrated above 50 °C.

Suddenly, the dominant theories in physical organic chemistry for the interpretation of the reactivity of hydrogen required reexamination. Initial ideas focused on the addition of a tunneling correction to classical theory, while an alternative view took its lead from the decades of work on electron tunneling. In the latter, the barrier to reaction is dominated by the heavy-atom reorganization energy, a prerequisite for effective wave-function overlap of the tunneling particle. While these divergent views were initially difficult to differentiate, data began to emerge that were incompatible with a tunnel correction, necessitating new theoretical frameworks for the interpretation of hydrogen reactivity in condensed phase.

One important distinction that enters into the quantum behavior of the electron and hydrogen is the huge mass differential of these particles. By virtue of its much smaller wavelength, movement of hydrogen between reacting centers is expected to be exquisitely sensitive to small changes in donor/acceptor distance. In the context of enzyme-catalyzed reactions, this has led to the realization that motions within a protein backbone must be invoked in the context of the hydrogen-transfer process and, furthermore, that optimal catalysis may require coupling of protein motions that are both proximal and remote from the enzyme active site. In the past several years, investigators have turned toward identifying the structural context for these modes, together with their amplitudes and time constants.

This recent history of hydrogen illustrates beautifully the interplay between advances in enzyme catalysis and solution chemistry. To a researcher at the interface of chemistry and biology, the power of interdisciplinary research to advance our knowledge is extremely gratifying.





Judith Klinman is a professor and former chair of chemistry and professor of molecular and cell biology at the University of California, Berkeley. She has received the ACS Repligen Award for her work on the chemistry of biological systems.

HYDROGEN AT A GLANCE

Name:From the Greekhydro genes,water forming.

Atomic mass: 1.01

History: Produced by scientists for years, but first recognized as an element by Henry Cavendish in 1766.

Occurrence: The most abundant element in the universe. Hydrogen as water is essential to life and is present in all organic compounds.

Appearance: Colorless, odorless gas.

Behavior: Flammable and explosive.

Uses: Used in NH3 production and to hydrogenate fats. H2 was once used in lighter-than-air balloons.

LITHIUM

PAUL V. R. SCHLEYER, UNIVERSITY OF GEORGIA




The third chemical element, lithium, was discovered in 1817 in a rocklike ("lithos") mineral, petalite, by J. August Arfvedson in J. J. Berzelius' laboratory in Stockholm. Uses, while currently ranging from batteries to antidepression treatments, developed very slowly. Unlike its close neighbors, hydrogen and helium, lithium has not been detected in interstellar space. Currently, adequate deposits exist on Earth in certain minerals and in brines. Lithium also is a trace element in the human body.

A century passed before Wilhelm Schlenk prepared the first alkyl lithium derivatives from organomercury precursors. Henry Gilman and Georg Wittig made important contributions to the development of the synthetic capabilities of lithium derivatives, which eventually replaced Grignard reagents as the primary anionic reactive intermediates. Preparative uses blossomed after both discovered independently that many compounds could be lithiated directly by hydrogen-metal exchange. This bypassed the often cumbersome preparations of Grignard reagents from magnesium metal and halogen derivatives. The latter had to be synthesized separately from the parent substrate.

With the readily available n-butyl lithium (used commercially as an anionic polymerization catalyst) typically serving as the reagent, hydrogen-lithium exchange often takes place specifically at positions in the vicinity of functional groups, where it is difficult to introduce a halogen. Such "directed metallations" are favored mechanistically as a result of the lowering of the metallation transition-state energy through lithium complexation with the functional group. Polar solvents and accelerators often, but not always, speed up metallation substantially and may even influence the site of substitution. Empirical trials and experience, rather than understanding, long guided the synthetic chemist.

One of my early mechanistic papers trying to unravel such mysteries computationally met the scorn of a referee: "Who cares what is in the pot, as long as it works." This comment equated lithium chemists with the witches in Shakespeare's "Macbeth." But it was my involvement with the highly peculiar structures of organolithium compounds that initiated my interest in lithium in the 1970s. At that time, alkyl lithium compounds were believed to be covalent. The evidence seemed clear. In contrast to insoluble ionic sodium and potassium alkyls with limited preparative utility, lithium alkyls are soluble in hydrocarbons and, when pure, either are liquids (n-butyl lithium) or have low melting points (ethyl lithium). The X-ray structures show tetramers with hexacoordinate carbons! Andrew Streitwieser pointed out later that a simple ionic cluster model with interpenetrating anionic and cationic tetrahedra (as in rock salt) could rationalize such structures. The hydrocarbon exterior of (RLi)4 tetramers explains their alkane solubility.

John Pople and I joined forces in the early 1970s. Our research groups employed Pople's newly developed Gaussian 70 ab initio program to computionally discover, for example, doubly bridged dilithioacetylene and a dilithiocyclopropane with a planar tetracoordinate carbon. Octahedral CLi6, with six equivalent bonds to carbon, is one example of "hypermetallation" out of many. These predictions were verified, at least in part, experimentally. Such rule-breaking structures illustrate the interplay of ionic and covalent bonding with some Li-Li interactions; the octet rule is not violated.

After my move to the University of Erlangen-Nuremberg in 1976, my coworkers added to the relatively small number of X-ray structures. Each new result revealed some new feature, which required a detailed computational study to understand. This complexity is still true. Lithium compounds are "self-assembling molecules" par excellence. They can aggregate in a variety of ways and bind not only to polar solvents (the "accelerators" mentioned above) and to benzene, but also to the substrates before reaction. The "witches brew" also can be clarified computationally, as can the reaction pathways and transition states. Attacking the planar tetracoordinate carbon problem again, we took advantage of computer modeling to design promising chelated (internally solvated) organolithium compounds. Their subsequent synthesis and X-ray analysis verified our prediction.



A chemical world based on electrostatic interactions, rather than just covalent bonds, is disclosed by the geometries, the bonding, and the course of reactions of lithium compounds. A goal of science is not just to understand what has happened, but to predict the outcome quantitatively.


TIES THAT BIND

Lithium compounds can display "hypermetallation" and an interplay of ionic and covalent bonding.


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