Issue Date: September 9, 2013
Trying To Explain A Bond
C&EN asked chemists who spend most of their time working with and thinking about chemical bonding to answer this question: What do you think a chemical bond is in reality? The chemists were prompted to consider how they might describe a chemical bond to a nonchemist, to a nonscientist, or to a family member. They were also asked to consider how a chemical bond impacts our daily lives. Add your thoughts on the chemical bond in the comments below.
Anthony J. Arduengo III, University of Alabama, studies the chemistry of compounds with abnormal valency and bonding arrangements and discovered the first isolable N-heterocyclic carbene.
Humankind’s ability to understand and transform matter through chemistry is as much the reason for the success of our species as the opposable thumb and language. Whether this chemistry was mastering combustion (fire) to release chemically stored energy (breaking and making bonds) to provide heat and light at will, or whether one chooses a suitable material (wood or stone, for example) to construct durable housing, tools, and devices to serve the interest of the human community, it’s all chemistry. When it comes to describing a chemical bond, I worry about being constrained by semantics, so I tend to have a rather broad view. I’ve often thought about applying a definition that is energy related—that is, a chemical bond must have a dissociation energy greater that X kcal/mol. But I’m not sure what constructive purpose such a definition would serve, and I can see how it might constrain creative thought. You certainly don’t want to ever get me started talking about valency or aromaticity. I’d as soon discuss the number of angels on a pinhead. Strangely, I’m much more comfortable discussing unusual valency, which seems to imply that I have some preconceived notion of what valence is—I don’t, but I’m happy to accept the usual historical definitions and claim all exceptions as “unusual.” But when talking to a layman, I have no problem in stating that a chemical bond is what holds materials (chemicals) together and gives them their properties.
Polly L. Arnold, University of Edinburgh, focuses on the synthesis of unusual metal complexes of rare earths, actinides, and early transition metals as catalysts.
A bond is the sharing of electrons between two or more atoms, and it doesn’t matter which atoms they belonged to originally or how many electrons are involved. To my mother, who struggles with the difference between atoms and molecules, I refer to bonds as the “glue” between atoms. The variety of glues on the market allows for some simple analogies to be made, ranging from Post-it note-strength hydrogen bonds to the water-soluble glues that hold ionic salts together and superglued polypropylene skeletons.
If we completely understood everything about chemical bonding, we’d do a lot less exploring and have long-list orders for drugs and chelators that we had to make. I’d probably become quite depressed.
Gregory H. Robinson, University of Georgia, is a specialist in the structure and bonding of organometallic compounds, in particular compounds containing multiple bonds between heavier main-group elements.
Chemists have seemingly always been grappling with the definition of a chemical bond. I simply believe that it is a cooperative truce among two atoms. Most of the time the truce concerns a pair of electrons, but on occasion the truce can involve two or three or more pairs of electrons. It’s almost as if the atoms are holding the electrons captive until something better comes along.
Martyn Poliakoff, University of Nottingham, is a pioneer in the field of green chemistry, in particular for chemical applications of supercritical fluids, and a cocreator of the Periodic Table of Videos.
My feeling is that most chemists don’t define bonds. They just know what bonds are when they see them. But many chemists anguish over whether particular atoms in their latest crystal structure are really bonded. The biggest surprise in my research career has been the invention of STM, AFM, and other scanning techniques. We can now actually “see” molecules and their bonds, and they look just as fuzzy as we imagined. The recent Science paper using AFM to show individual molecules before and after they react is something that I was brought up to think was quite impossible. Do the images teach us much about bonding? No, but chemistry is an immensely visual subject and, if we are honest, most of our images are theoretical constructs. So it’s quite reassuring that molecules do really appear to resemble our pictures.
James L. Marshall and Virginia (Jenny) Marshall, University of North Texas. The Marshalls’ special collection of the elements features samples of minerals collected at the sites from which all the natural elements were originally discovered.
A real impediment to the correct understanding of chemical bonding in the early 1800s was the commonplace belief that like atoms could not combine with one another. This was the logical conclusion ensuing from electrolysis experiments of metals, historically credited to Humphry Davy (1778–1829) but in fact pioneered by Jöns Jakob Berzelius (1779–1848). Berzelius noticed that in a battery the “alkalis and earths” were drawn toward the negative pole, and that oxygen, acids, and oxidized substances migrated to the positive pole. He was very impressed by the electrical forces needed to rip apart these reactive metals, and he championed the hypothesis of “electrochemical dualism,” wherein the forces holding atoms together were positive-like (the metals) and negative-like (the nonmetals). Berzelius was thus the first to conceptualize ionic bonding, and he assumed this bonding occurred in all compounds. The idea that elemental hydrogen and chlorine were H2 and Cl2 never occurred to him, and his ideas held sway through the first half of the 19th century.
One of the first to understand “like”-bonding was Jean Baptiste André Dumas (1800–84). In his mid-20s, Dumas was asked to explore the reason why the burning candles in the Tuileries Palace were emitting obnoxious odors. Dumas is best known today for his method of molecular weight determination, currently used in undergraduate chemistry labs. He found that these candles had been bleached by chlorine and that the irritating stench was hydrogen chloride. His curiosity piqued, he followed up with research that allowed him to conclude that either H (positive-like) or Cl (negative-like) could combine with carbon in a similar way, without losing the general physical properties of the compound. That is, the compound still acted like an organic substance. By 1828 he introduced the terms “molécule chimique” (atoms) and “molécule physique” (true molecules). Thus, he formulated the idea that there must be an additional type of bonding, which today we recognize as covalent bonding.
Richard Eisenberg, University of Rochester, studies inorganic and organometallic compounds applied to photochemistry and solar energy conversion and is past editor-in-chief of ACS’s journal Inorganic Chemistry.
In teaching bonding of homonuclear diatomic molecules to first-year students in chemistry, I am always fascinated by the electronic structure of molecular oxygen. It possesses unpaired electrons when a Lewis structure suggests otherwise, and because of those unpaired electrons it can be trapped at low temperature between the poles of a magnet as a blue liquid. This is extraordinary. The paramagnetism makes O2 relatively unreactive. In fact, there is an excited state of O2 that lies close in energy to the ground state, and it can be generated relatively easily. This singlet form of O2 is exceedingly reactive. I tell students that if O2 really existed as a singlet molecule rather than as a triplet with two unpaired electrons, life as we know it would never exist because we would all be turned into inorganic oxide solids. It is the triplet state of O2 that allows life to survive on this planet, and it is molecular orbital theory that allows us to understand why.
Debbie C. Crans, Colorado State University, studies the chemistry and biochemistry of vanadium and other transition metals, fueled by their applications in medicine and their mechanisms of toxicity.
Chemical bonding is the association between atoms facilitated by electrons, and as such produces inherent properties manifested in chemical structure, stability, and reactivity. Structure: When electrons fill up the space between atoms they create a material, which is characterized as having a bond between the two or three components in question. Electrons and bonds are therefore responsible for the shape and volume of molecules. Stability: Bonding can be envisioned as the glue that keeps the different parts of the molecules together. When bond distances are near the ideal, the molecule is stable. Reactivity: Reactivity is a direct consequence of the nature of elements, molecular shape, and bonding. Weak bonds are readily and rapidly broken and can be formed as a result of, for example, reactive forms of elements, undesirable shapes, and long bonds.
Bonding is defined differently in the life sciences and even within each field of chemistry. At a general level, a stick-and-marbles model set can be used to envision molecules and their properties. However, that description does not properly describe the wave nature of electrons and their probabilistic location. Organic chemists use the simple hybridization explanations to describe and understand bonding to tetrahedral carbon. Physical chemists use mathematical and statistical equations to explain the electronic properties of materials. Each description and approach to bonding has strengths and limitations. Organic chemists are simplifying systems and as a result can work with complex molecules. Inorganic chemists concern themselves with a large range of different elements and embrace the relativistic aspects, but as a consequence of the diversity lack the well-developed framework to make predictions that organic chemists have. Physical chemists embrace the mathematical and electronic details but ignore facts such as shape and 3-D occupancy of space, and as a result address mainly electronic and statistical properties. Yet, any discovery exploring these parameters is critical for the future progress of chemistry.
Leo Manzer, an expert in catalysis, is a retired DuPont chief scientist (DuPont Fellow) and is currently head of the chemistry consulting firm Catalytic Insights.
When I think of developments that have impacted society and involved chemical bonding, I think of the development of catalysts in the petroleum industry to convert crude oil to transportation fuels such as gasoline, diesel, and jet fuel. This has largely involved breaking and making C–C and C–H bonds. Without the ability to selectively crack or break the chemical bonds in the viscous oil that comes out of the ground to meet the stringent needs of engine producers, we might still be traveling on coal-fired trains and riding horses or bicycles. In addition, during this bond-breaking process, contaminants such as nitrogen and sulfur are reduced to extremely low levels so that the combustion products don’t contaminate the environment as they did in the early days. Still, new challenges are being tackled by chemists as they learn how to catalytically convert renewable feedstocks such as sugars and wood chips into renewable fuels and chemicals. This involves new technologies for the conversion of C–O and C–OH bonds to C–C bonds.
I also think of the amazing use of chlorofluorocarbons (CFCs) to provide refrigeration to prevent food spoilage; air-conditioning to cool office buildings, homes, and cars; highly insulating foams for energy efficiency; and solvents for circuit board cleaning that helped the computer industry grow. Early on, CFCs were manufactured by carefully and safely creating chlorine and fluorine bonds to carbon. After more than 50 years of production, it was recognized that these CFCs were depleting the ozone layer. Industrial scientists quickly learned how to make new fluorocarbons without chlorine. The ability of chemists to identify, prepare, and selectively eliminate C–Cl bonds on a large scale allowed society to continue operating with little disruption.
These technologies were truly major advances in catalysis and bond manipulation.
Alexander I. Boldyrev, Utah State University, studies theoretical and computational chemistry of new compounds, and is coorganizer of the International Conference on Chemical Bonding.
Chemical bonding is at the heart of our chemical language, and it is extremely important for teaching. When I introduce myself to a nonchemical audience I frequently hear a comment that “I hated chemistry. I really did not understand it.” Our science indeed is very complicated, and a big part of the problem is how we teach chemical bonding. We teach the “golden chemical bonding model” based on Lewis structures. That is a simple concept, and it’s easy to teach. Then we have aromaticity, which is a quite fuzzy concept. People are still arguing about how to recognize aromaticity. Then, we teach valence bond theory, where bonds start to jump from one part of a molecule to another. That confuses freshman students, especially those who were not previously exposed to chemistry. Finally, we teach molecular orbital theory. Now, chemical bonds completely disappear from the picture and we have orbitals instead. I am not surprised that many students, even if they passed general chemistry classes, are still confused about chemical bonding. I personally believe we need to develop a comprehensive chemical bonding theory that will be able to describe most of chemistry. That will help to teach our beautiful science. It will help build financial support of our science and our standing in society.
Akira Sekiguchi, University of Tsukuba, is a specialist in organosilicon chemistry and multiple bonding in main-group elements.
To date, there are many classes of chemical compounds that do not conform to the standard definitions of covalent and ionic bonds. These are the so-called nonclassical compounds. Among them are odd electron bonds such as radical species, hypervalent bonds in molecules with an expanded octet, electron-deficient bonds such as three-center two-electron bonds commonly found in boranes, singlet biradicaloid bonds in the highly strained cluster hydrocarbons propellanes, trans-bent multiple bonds between the heavier main-group elements, a covalent form of the ionic bond in pyramidal shaped hydrocarbons, and more. However, even given the number of these nonclassical compounds and nontrivial bonding situations in them, I still favor the general definition of the chemical bond as the attractive interaction of electrons provided by the participating atoms. This is just like a handshake between two atoms: each atom stretches its “arm” (electron) toward the other, and when these two “arms” (electrons) meet, then the “handshake” (chemical bond) takes place.
Kendall N. Houk, University of California, Los Angeles, solves problems in organic and bioorganic chemistry using theoretical and computational methods.
A chemical bond is what holds atoms together in molecules. Bonds arise from the electrostatic forces between positively charged atomic nuclei and negatively charged electrons (the positions of which in space are determined by quantum mechanics). Pretty simple, except for the parenthetical phrase!
Cathleen M. Crudden, Queen’s University, Kingston, Ontario, focuses on chiral catalysis in organic synthesis and served as 2012 president of the Canadian Society for Chemistry.
The accurate depiction of bonding arrangements in molecules is critical to chemists and chemistry. Take valency, which is something that chemists feel we understand. Carbon likes four bonds, nitrogen three, oxygen two. Simple. Yet, for those of us who teach second-year organic chemistry and see errant pentavalent carbons on a regular basis, molecules that don’t fit our typical idea of valency can be challenging! But there’s no doubt that some such odd molecules do exist. When atoms don’t follow the normal rules, we like to think we understand the consequences. For example, carbon with only two bonds is not a happy creature. These so-called carbenes tend to dimerize, cyclopropanate, insert into C–H bonds, or generally react with just about anything at hand. Until recently, the idea of creating a class of carbenes that can be stored, crystallized, or even distilled would have been largely unthinkable. However, placing two heteroatoms such as nitrogen on the carbene carbon and then confining these two heteroatoms in a ring made the unthinkable a reality. N-heterocyclic carbenes and their derivatives are not only more stable than typical carbenes, they are exceptionally valuable ligands for transition metals, serve as organocatalysts, and have increasing applications in materials chemistry. Thus it seems that taming divalent carbon has been not only a fun endeavor but a useful one as well.
But if asked to explain bonding to a nonscientist, I would probably equate chemical bonds to molecular glue. Carbon-hydrogen bonds would be like superglue, very strong and difficult to break with ordinary methods. The significant amount of energy stored in carbon-hydrogen bonds is one of the things that makes hydrocarbons (perhaps unfortunately) such spectacular fuels because this energy is released when they are burned. Bonds that are significantly weaker, such as hydrogen bonds, are also important. Although individually weak, when a multitude of them act in concert, the effect is dramatic. Consider the fact that water, with only three atoms and a molecular weight of 18, boils at 100 °C. Compare this with methane that has no hydrogen bonding ability, which boils at –164 °C, not so far from liquid nitrogen at –195 °C. It is hydrogen bonding that is responsible for this dramatic change in properties, so tea and coffee drinkers should thank the humble hydrogen bond for the fact that water has the perfect boiling point!
Arnold L. Rheingold, University of California, San Diego, is one of the world’s most prolific crystallographers.
A chemical bond forms when two or more atoms in close proximity achieve a lower overall energy either by creating new orbitals encompassing multiple nuclei or by the transfer of one or more electrons from one atom to another. If asked to state this in a way that my grandmother might understand it, I’d say: “You and grandpa have been together for more than 50 years. Over that time the two of you have functioned better as a team than you would have separately, and during that time you have shared just about everything. The two of you are clearly happier together than you would have been apart. On the other hand, I have known many couples that are just as happy by clearly defining separate roles for themselves by dividing responsibilities. In the world of chemistry, you and grandpa formed a share-all covalent bond, while other couples, for whom functions are kept separate, form an ionic bond.”
Marcetta Y. Darensbourg, Texas A&M University, carries out the synthesis of transition-metal catalysts as mimics for natural hydrogenase enzymes for producing hydrogen.
As only he can, the great professor Harry Gray of Caltech sometimes describes how inorganic chemists account for inexplicable structure and bonding mysteries: “If the Jahn-Teller effect doesn’t work, we go for π-backbonding!” While the former [geometrical distortion of a molecule] is rarely appropriate in my research, the π-backbonding idea definitely applies.
The beautiful simplicity of dative bonding in classical coordination chemistry had to make way for electron delocalization in the huge class of metal carbonyls developed by German chemists in the 1930s and onward. How does it work? Compounds with metals in impressively low oxidation states, nickel(0), cobalt(–1), and iron(–2) in Ni(CO)4, Co(CO)4– and, Fe(CO)42–, follow a pattern of 18-electron counts about the metals. The CO ligands donate two electrons each from a lone pair “forward” to the metal to form a σ-bond, and combined with the d-electrons on the metal, it usually equals 18 electrons. Additionally there are empty antibonding π orbitals on CO that “back”-accept those d-electrons from the electron-rich metals, resulting in decreasing the CO bond order and stabilizing the metal-carbon bond. This push-pull effect is also seen in any metal-ligand system where empty π antibonding orbitals of (most usually) carbon-based ligands, such as olefins or cyclic aromatics, are an energetic match of the filled metal’s d-orbital set.
Synthetic organic chemists and the chemical industry have capitalized on the metal’s way of binding and activating CO and olefins in this manner. But what’s amazing for my research, an organoiron unit, Fe(CN)2(CO), as in the piano-stool organometallic complex, (C5H5)Fe(CO)(CN)2−, is completely analogous to one found in biology. The active sites of hydrogenase enzymes, natural biocatalysts that facilitate production or use of hydrogen, contain iron-carbonyl bonds replete with this π backbonding. As a synthetic inorganic chemist, I can use my infrared spectrometer to contrast my synthetic analogs with biological moieties. Harry is correct––it works!
Douglass W. Stephan, University of Toronto, conducts research in inorganic main-group and organometallic chemistry and is the discoverer of frustrated Lewis pair reactive molecules.
Chemical bonds are the glue that bind atoms together into the ensembles that are molecules. Understanding the nature of bonds, how they influence the properties of molecules, and developing methods to judiciously reconfigure bonds among atoms are the goals of chemistry. These studies can lead to dramatic new technologies that give us such things as pharmaceuticals, flat-screen TVs, and plastics. However, the study of chemical bonds provides insights well beyond the pragmatic. Everything we make, everything we do, and everything we are requires the breaking and making of chemical bonds to supply the energy and yield the fruits of our labors. While chemistry is often focused on real-world goals, in a broader philosophical sense the study of chemical bonds is the study of all that is our world.
Robin D. Rogers, University of Alabama, Tuscaloosa, is director of the Center for Green Manufacturing and editor of the ACS journal Crystal Growth & Design.
Being somewhat of a contrarian, my thoughts on chemical bonding tend toward developments that would more accurately be said to challenge what a bond is, rather than to strictly define it. The development of supramolecular chemistry and the rising importance of coordination chemistry, coupled with hydrogen bonding and halogen bonding, have led to a new way of thinking about how to make molecules and design materials with defined macroscopic properties that goes beyond the classic covalent bond we all learned. This has led to rather heated arguments on exactly what constitutes a chemical bond. Is a bond to be defined by sharing of electrons? If so, how many electrons, and to what degree of sharing? How does this work when virtually every interaction between atoms or molecules involves some sharing of electron density in a continuum of states between absolutely no sharing to absolutely equal sharing? Perhaps it is time to refrain from restrictive nomenclature that limits our imagination by constraining our expectations. Combining all of our advancements in experimental and computational science, let’s try following the electron density and trying to relate properties not just with “ionic” or “covalent” bonds, but with the actual degree of electron density shared.
Anastassia N. Alexandrova, University of California, Los Angeles, conducts theoretical and computational chemistry of proteins, enzymes, catalytic surfaces, and clusters and is coorganizer of the International Conference on Chemical Bonding.
In defining a chemical bond, I tend to be conservative and refer to Pauling’s book “The Nature of the Chemical Bond.” There is a chemical bond between two atoms or groups of atoms in the case that the forces acting between them are such as to lead to the formation of an aggregate with sufficient stability to make it convenient for the chemist to consider it as an independent molecular species. This definition is both inclusive of various types of bonding (if you are a connoisseur) and simple for a nonscientist to grasp. The foundation of every chemical or physical phenomenon is electronic structure. But do we always have to run an electronic structure calculation or take a spectrum to gain a chemical insight? I hope not. For chemists, electronic structure traditionally and fruitfully translates into the qualitative language of chemical bonding. This language, though rooted in the solution of the Schrödinger equation, is simple enough to enable intuition and fast thinking about chemical systems. Whether you think of efficient heterogeneous catalysts, organic solar cells, or fluorescent labeling of biomolecules, consider bonding—it will help to guide and accelerate your progress. The theory of chemical bonding is undergoing rapid development, and it should be appreciated as a simple qualitative tool for the design of structures and properties of materials and molecules across the discipline.
Sason Shaik, Hebrew University of Jerusalem, carries out theoretical and computational chemistry with a focus on reactions of metalloenzymes and new chemical bonding concepts.
Given that every sticky interaction is called a chemical bond today, it is impossible to give an effective and productive description of “the chemical bond.” If, however, we focus on the most common bond that holds molecules together, this would be the electron-pair bond, which lowers the energy of the bonded fragments by 20 to 141 kcal/mol per bond. Accepting that the electron-pair bond leads to formation of molecules, it can be explained to nonscientists and to grandmothers, using the simple imagery of a game of Lego. Take two H atoms. Each one has an electron. If the two electrons have opposite spin around their axes, then when the H’s approach one another they click to make a bond. This click bond obeys magic numbers, which allow you to construct and conceptualize an infinite number of molecules. This is how I teach chemistry to humanities and social sciences students in my course at Hebrew University.
Considering how the chemical bond impacts our daily life and our existence, imagine if water molecules did not obey the rules of bonding and were linear instead of bent. Water would then most likely be a gas. Where would we then be? Take graphite and diamond. They have different bonding patterns, and as a result, graphite is cheap and ugly, while diamond is beautifully scintillating, has a high value, and evokes emotions such as falling in love. Think about the retinal molecule in our eyes. Because of its bonding in cis and trans structures, we can “see.” The eternal homochirality of living matter is a result of bonding that makes proteins and sugars persistently chiral. Think about DNA and RNA: Our genetic code is a chemical code of architecture and dynamics of weak bonds (hydrogen bonds). Matter without the chemical bond as we know it would be an atomic soup!
Gabriel Merino, Center for Research & Advanced Studies of the National Polytechnic Institute, in Mérida, Mexico, conducts computational chemistry to understand and predict new molecules.
The chemical bond is a fuzzy concept, explained by limited models, that has historically been full of intense debate and controversy. It is really complicated to understand that chemical bonding is a concept, and a chemical bond is not a real object. In this regard, the teaching of chemical bonding promotes many misconceptions when this point is not clear. If, from the beginning, we teach that chemistry is based on models, perhaps it will be simpler to understand why sometimes two models provide conflicting views for the same chemical problem. In my opinion, one way to understand chemical bonding is forcing bonding to extreme situations. Molecules under pressure, delocalized systems, nonclassical molecules, transition states, and many other challenging systems are strong motivations to many chemists all around the globe to continue our quest in better understanding this difficult, challenging, controversial, but fascinating cornerstone concept of chemistry: the chemical bond.
Roald Hoffmann, Cornell University, was a 1981 Nobel Laureate in Chemistry for his theories concerning the course of chemical reactions.
I think that any rigorous definition of a chemical bond is bound to be impoverishing, leaving one with the comfortable feeling, “yes (no), I have (do not have) a bond,” but little else. And yet the concept of a chemical bond, so essential to chemistry and with a venerable history, has life, generating controversy and incredible interest. My advice is this: Push the concept to its limits. Be aware of the different experimental and theoretical measures out there. Accept that at the limits a bond will be a bond by some criteria, maybe not others. Respect chemical tradition, relax, and instead of wringing your hands about how terrible it is that this concept cannot be unambiguously defined, have fun with the fuzzy richness of the idea.
Josef Michl, University of Colorado, conducts physical organic chemistry studies of organic and organometallic compounds, including molecular rotors and molecular circuits.
What I told my grandmother is that atoms are sticky. Once two of them stick together in a line, or three in a triangle, they require an effort to pull apart and are said to be connected through a bond. This effort can be small when a bond is weak and the atoms far apart (van der Waals attraction), or a little bigger (hydrogen bonds), or larger still (coordination), or huge when the atoms are really close (covalent, ionic, and multicenter). We indicate a bond between two atoms with a line or several lines, and can display the many bonds present in straight and branched chains, rings, and cages. It is harder to draw three-center bonds (sometimes we use dashed lines).
Why are atoms sticky? They contain heavy positive nuclei and light negative electrons flying about, unable to leave due to electrostatic attraction. When two or three atoms join, their electrons are attracted to each nucleus and repelled by each other. They modify their motion, lowering the total energy even though the nuclei repel. Pushing the atoms closer than the optimal distance would modify the electron motion differently, increase nuclear repulsion, and augment the energy again.
By now, my grandmother’s curiosity waned and I did not get a chance to tell her about quantum mechanics and about what the world would look like if atoms were not sticky.
Dean J. Tantillo, University of California, Davis, is a theoretical organic chemist studying the mechanisms of cascade polycyclization reactions and the design of new catalysts.
To me a bond is simply the attraction between atoms. This attraction may be covalent or ionic, or be labeled with a more specific name like sigma bond, pi bond, delta bond, conjugation, hyperconjugation, percaudal interaction, salt bridge, hydrogen bond, halogen bond, dative bond, charge shift bond, and on and on and on. I tend to think of bonding, rather than bonds. How much bonding? How large is the favorable interaction energy? What are the origins of the attraction? These concepts I digest and describe in terms chemists are generally comfortable thinking about—charges attracting, orbitals with appropriate numbers of electrons overlapping, and so forth. In short, it’s all about continua, in terms of strength, length, and relative contributions of different sources of attraction.
Joel S. Miller, University of Utah, studies multicenter carbon bonding as well as organic-based magnetic materials.
I think a chemical bond is a stabilizing, or attractive, interaction between atoms that significantly alters the properties of the atoms, leading to an independent or new species with new properties. This definition does not state how stabilizing or significant “significantly” has to be. One could limit the definition to something that can be put into a bottle, but justifiably others would argue against that limitation. In a more general way, a chemical bond is a construct, frequently a pictogram, used by chemists to understand the stronger attractive interactions among atoms that enable the organization and understanding of the structure, properties, reactivities, and interrelations for the growing myriad of substances. The concept of chemical bonding impacts our daily lives because it is a language enabling chemists to communicate among each other and facilitate the design and synthesis of improved substances that have benefited mankind. A chemical bond enables us to glean the order and complexity of how the basic building blocks, atoms, interact to form all substances. This language has enabled the broad enterprise of chemistry to rapidly develop and flourish into the central science.
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For the chemists, I would revert to Linus Pauling's book, The Nature of the Chemical Bond, page 6. Linus Pauling is our Saint Nicholas - - the patron saint of chemists.
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