Chemistry: Chemical Bonds

Introduction

Since the dawn of chemistry, scientists have sought to understand the forces that hold atoms together in a chemical bond. Since molecules (a name that means “little masses”) are invisible to the naked eye, researchers have had to rely on indirect methods to unravel these structures. American physical chemist Gilbert Lewis (1875–1946) endowed molecular theory with a formidable array of concepts that continue to influence how structural chemistry is practiced and taught throughout the world. The consensus on the chemical bond that emerged during the early twentieth century has yielded powerful insights into many areas of science.

Historical Background and Scientific Foundations

An important inference about bonding was that some compounds are “ionic”—they form charged species when dissolved in a liquid (such as water). Throughout the nineteenth century, a small but distinguished group of scientists subscribed to this hypothesis. These ionists (as their opponents often called them) included, among others, the Prussian physicist Rudolf Clausius (1822–1888), and the British scientist Michael Faraday (1791–1867). In addition to his wide-ranging discoveries, Faraday introduced many of the terms that scientists continue to use today: ion (a charged particle), cation (a positively charged species) and anion (one bearing a negative charge). While his studies focused on high-temperature melts, they helped to bring ions into the common vocabulary of science.

During the nineteenth century, speculation and debate centered around the electrical interactions that could stabilize chemical compounds. These concepts acquired a new legitimacy in 1897 when the British physicist Sir Joseph J. Thomson (1856–1940) published the first of a series of experimental articles on the properties of electrons from gases subjected to a high voltage.

His discovery that cathode rays consisted of such particles (for which he eventually measured both the charge and the mass) led to new efforts to incorporate the electron into models for the chemical bond. As a physicist, Thomson relied on familiar principles (like

electrostatic attraction) to explain the existence of molecules. His early theories treated all compounds as ionic and thus could not adequately account for the stability of molecules that did not dissociate into charged particles. In spite of such drawbacks, other scientists did embrace Thomson's ideas and developed variations of his electrostatic model.

The Science

One chemist who made a distinction between ionic and covalent compounds was Lewis, who published a seminal paper in 1916 entitled “The Atom and the Molecule.” Like many scientists who also served as instructors, his research on the chemical bond began in 1902 as an effort to help his students understand the valence of different elements. Later, the advent of quantum mechanics drastically altered the theoretical interpretation of Lewis's model, but it remains part of the modern vocabulary of every chemist.

Lewis was quite unusual in that he spent over a decade developing his bonding theory before publishing it. This was partly because as an instructor at Harvard University he could find no chemist interested in working on such abstract concepts. By 1916, however, he had a large department of chemists at the University of California on whom he could rely for constructive feedback.

Another benefit of this delay was that researchers had a much better count of the number of electrons for each element, thanks to the spectroscopic data compiled by British physicist Henry Moseley (1887–1915). Lewis was able to incorporate this information into molecular pictures, called Lewis structures, that use dots to represent the number of valence electrons contributed by each atom in a covalent bond. Each dot denotes an electron: In water (H₂O), the hydrogen contributes one electron, while the oxygen atom has six in its valence shell. Lewis assumed that these latter atoms also had two electrons in an inner core, which did not play a direct role in the bonding and thus were not shown explicitly. A noteworthy feature is that the oxygen has two pairs of valence electrons not used for bonding.

Lewis's real innovation was his treatment of covalent bonds as shared electron pairs. Therefore, the two dots that hydrogen and oxygen have in common denote a single bond in the water molecule, and the four electrons between the oxygen molecules in O₂ represent the much stronger double bond. He attributed the affinity of hydrogen (with only one electron) for other atoms as a natural tendency to complete its shell, achieving a configuration similar to that of helium (with two electrons). Likewise, oxygen in his scheme will usually form two single bonds or one double bond, so that it will acquire another two electrons, like neon. Lewis thus explained his bonding rules by noting that neon and helium belong to a group of elements known as noble gases, which have filled electron shells and are especially unreactive.

One of Lewis's most compelling achievements was that his 1916 publication solved several puzzles that had plagued chemists for generations, including the structure of ammonium chloride. Some chemists believed that NH₄Cl really consisted of two molecules: an ammonia part (NH₃) in close proximity to HCl. Other scientists believed that nitrogen in this molecule must have five “ionic” bonds to H^- and Cl^- anions.

Lewis cleverly solved this conundrum by suggesting that nitrogen may form four covalent bonds to hydrogen, but also an ionic bond between NH₄⁺ and Cl^-. Implicit in this interpretation is the central concept that a base can be defined as an electron-pair donor, and that an acid can be an electron-pair acceptor. However, a similar structure was proposed a decade earlier by Swiss chemist Alfred Werner (1866–1919), whose work was less widely read or discussed.

GILBERT NEWTON LEWIS (1875–1946)

Gilbert Newton Lewis (1875–1946) was born in the town of West Newton, Massachusetts, and was raised by progressive, independent parents. At the age of seven, his family moved to Nebraska, and Lewis was homeschooled until he attended high school in the city of Lincoln. His parents and teachers recognized his remarkable intelligence, and at the age of 16 he was admitted to the University of Nebraska.

Seeking a more cosmopolitan environment, he transferred to Harvard University, from which he received his bachelor's degree in chemistry in 1896. After a year of teaching at Phillips Academy in New Hampshire, Lewis returned to Harvard to begin research for his doctorate. His faculty advisor was Theodore Richards (1868–1928), a well-known American chemist who received a Nobel prize in 1914 for his precise atomic weight measurements.

Lewis earned his doctoral degree in 1899, then spent a year in the German laboratories of chemists Walther Nernst (1864–1941) in Göttingen and Russian-born Wilhelm Ostwald (1853–1932), then in Leipzig. By the time Lewis arrived, each of these scientists was an established figure in the emerging field of physical chemistry. The American student followed an educational route that was common at the end of the nineteenth and the beginning of the twentieth centuries: Aspiring scientists from the United States performed an apprenticeship at more established universities in Germany, then the world leader in science.

In 1901 Lewis returned to Harvard as an instructor in chemistry, only to leave after three years to work in a government laboratory in the Philippines. In 1905 he accepted an offer to join the faculty of the Massachusetts Institute of Technology, then in a period of expansion under the direction of American analytical and organic chemist Albert Noyes (1857–1941). He remained at MIT for seven years, focusing his research on the energies of chemical reactions.

Lewis was well-known enough by 1912 to be hired by the University of California in Berkeley to head its chemistry department. Under his leadership, he encouraged instructors to learn about recent developments in their discipline and to teach those concepts in the classroom. He also organized weekly meetings at which students and faculty members presented their research and discussed its significance. Such sessions were common in Europe, but rare within any American university at the time.

Lewis remained at the University of California for the rest of his career, continuing the research he had begun at MIT and Harvard. In later years he shifted his focus increasingly to the study of spectroscopy, developing new methods to characterize light. He also played an important role as the author of books that trained a generation of chemists in his concepts and models. He suffered a heart attack and died in 1946 while working in his Berkeley laboratory.

Among the early advocates of Lewis's model were his fellow chemists at the University of California. Wendell Latimer (1893–1955) published an important discussion in 1920, in which he analyzed the unusual behavior of water. Its most remarkable property is that water remains in the liquid state over such a wide range of temperatures: It melts at 32°F (0°C) and boils at 212°F (100°C) at normal atmospheric pressure and was therefore chosen as the basis of the Celsius (centigrade) scale. By contrast, the related compound hydrogen sulfide (H₂S) is liquid only over a 73°F (23°C) range and boils at -76°F (-60°C).

Latimer suggested that water's strange behavior might be due to the electron pairs on the oxygen, which could be attracted to the hydrogens from neighboring molecules. Our modern name for this interaction is hydrogen bonding, and it is widely believed to explain many of water's properties. Such bonds are much weaker than a covalent bond, but each water molecule can act as both a donor and acceptor. The high boiling point thus reflects the cumulative energy (in the form of heat) needed to transform liquid water into a gas, where such hydrogen bonds no longer exist.

Modern Cultural Connections

A fundamental defect of Lewis's approach was that he could not explain why a pair of atoms should form a shared-electron bond. Based on the laws of physics as they were known in 1916, the electrons should really fly apart, repelled by its electrical charges. Once again, he was painfully conscious of his dilemma, and stated that such rules needed to be “suspended” for his model to work. Other physicists criticized the implication that the electrons had to remain in one place for a bond to form, since they thought that it was more likely that the electrons would be in motion. In the end, a revolution in physics was necessary before a new justification of the shared pair would emerge.

During the late 1920s, the way scientists viewed electrons underwent dramatic change with the emergence of quantum mechanics. According to the interpretation

of German physicist Max Born (1882–1970), the electron was no longer a particle fixed in space and time, but a wave for which you could only calculate the likelihood (or probability) of where it might be found. American chemist Robert Mulliken (1896–1986) thus introduced the term orbital to describe such regions of space.

Another important conclusion from quantum theory is that electrons only exist in certain energy states. In 1927 a pair of German physicists published the first such treatment of the hydrogen molecule. While Fritz London (1900–1954) and Walter Heitler (1904–1981) explicitly included a stabilizing attraction between the electrons and the nuclei, they also viewed the bond as an exchange between the atoms. Furthermore, they showed that the two electrons in hydrogen (H₂) reside within a true bonding level, which is lower in energy than the free atoms. The study by Heitler and London marked the first time scientists provided a compelling rationale for the pairing of electrons in the Lewis model.

This strategy developed by Heitler and London is known as valence bond theory. As its name implies, one characteristic was that this approach ignored the inner-shell (core) electrons and thus underestimated repulsive effects. However, valence bond theory can also be viewed as a segmented treatment, in which molecules are formed one bond at a time.

Other physicists and chemists adopted a more holistic approach, in which the bonds arise naturally from the interaction of the orbitals from each atom within a molecule. This strategy was usually called molecular orbital theory, and it was first applied in 1927 to the hydrogen molecule by the American physicist Edward Condon (1902–1974). Refinements were introduced by the German physicist Friedrich Hund (1896–1997) and by Mulliken.

The advocates of molecular orbital theory, especially Mulliken, rightly claimed that they achieved more accurate predictions of molecular structure (such as the bond angles in NH₃ and H₂O). On the other hand, such calculations were extremely tedious (when everything was done with a pencil and paper), and the resulting image of bonding was much more abstract. Inevitably, this research was slowly accepted by the community of chemists, but most still embraced the graphic nature of the valence bond approach.

Almost a century after Lewis introduced the shared electron pair, this image remains a useful concept for every chemist. While scientists have frequently debated whether chemical architecture is best viewed through the lens of valence bond theory or molecular orbital methods, these approaches are now seen as complementary. Some daunting challenges remain, especially in applying these tools to the three-dimensional structure of proteins. Our understanding of the chemical bond has been greatly aided by theoretical modeling, and the next generation of scientists undoubtedly will apply such methods in exciting ways.

See Also Chemistry: Molecular Structure and Stereochemistry; Physics: Spectroscopy; Physics: The Quantum Hypothesis.

bibliography

Books

Stranges, Anthony N. Electrons and Valence: Development of the Theory, 1900–1925. College Station, TX: Texas A&M University Press, 1982.

Periodicals

Hoffmann, R., S. Shaik, and P.C. Hiberty. “A Conversation on VB vs. MO Theory: A Never-Ending Rivalry?” Accounts of Chemical Research 36, no. 10 (September 5, 2003): 750–756.

Kohler, R.E., Jr. “The Origin of G.N. Lewis's Theory of the Shared Pair Bond.” Historical Studies in the Physical Sciences 3 (1971): 342–376.

Latimer, W.M., and W.H. Rodebush. “Polarity and Ionization from the Standpoint of the Lewis Theory of Valence.” Journal of the American Chemical Society 42 (1920): 1,419–1,433.

Lewis, G.N. “The Atom and the Molecule.” Journal of the American Chemical Society 38 (1916): 762–785.

William J. Hagan