Visual images of atoms
The last opposition to the existence of atoms vanished in the early 20th century when techniques were developed that portrayed visual representations of atoms. The first such techniques made use of the diffraction of X-rays, where the pattern of interference between rays that are reflected by a crystal can be interpreted in terms of the scattering from individual atoms. More images of atoms were produced in the 1960s by using methods that stripped electrons out of arrays of atoms at the surfaces of solids so that a map of the surface could be made, as well as by using improved techniques in electron microscopy that increased the resolving power of the microscope to nearly the point where individual atoms could be distinguished. The most visually compelling evidence came in the 1980s with the development of scanning tunneling microscopy. In this technique a needle point sharpened to consist of a single atom is moved like a delicate plow just above the surface of a sample, and its position is monitored. The results appear in the form of a visual image of the sample’s surface. The technique has been perfected to a point where it can be used to determine the locations of individual atoms. Of these techniques, electron microscopy comes the closest to an actual “sighting” of an atom, as the image requires the least construction. Images are obtained from X-ray diffraction data only after intense mathematical manipulation. Both field-emission and scanning tunneling microscopy give portrayals of the properties of a surface on an atomic scale and show atomlike features.
Molecular structure
Most chemists were confident that atoms really existed long before these sophisticated techniques provided such irrefutable evidence. In the 19th century, when the compositions of countless compounds were being determined, it was found that in certain cases different compounds have the same chemical composition. Thus, the composition C3H4 was found for two entirely different organic compounds (as judged by both their physical and chemical properties)—namely, propyne and allene. Confident about their analyses, chemists were forced to the conclusion that the two compounds differ in the manner in which their constituent atoms are linked together. In modern terms, the compounds are represented, respectively, as:
(The nature of the links between atoms is the major topic of this article and is discussed in detail below.) Thus, the sense of molecular structure (i.e., the arrangement of atoms in space) entered chemistry and, by implication, supported the view that atoms are real.
About the same time (in the 1860s), a more subtle aspect of structure became apparent—that of the three-dimensional spatial disposition of atoms in molecules. The concept of molecular structure began with the realization that atoms have different neighbours in different compounds even though their overall chemical compositions might be the same (as in the two structures corresponding to the formula C3H4). This is a topological distinction, meaning that the distinction is based on which atom is linked to which atom. The additional distinction introduced is geometric, referring to the spatial disposition of atoms relative to one another. As an example of this kind of distinction, the compound dichloromethane (CH2Cl2) can be considered. The topological structure of this molecule is:
with the hydrogen and chlorine atoms linked to a central carbon atom. It was observed that there is only one such compound. The significance of this is that the molecule cannot be planar, because, if that were the case, two different molecules of formula CH2Cl2 would be found:
The fact that there is only one dichloromethane suggests that its molecules are tetrahedral, for then, in whichever arrangement the four hydrogen and chlorine atoms are linked to the central carbon atom, the molecule is identical (apart from its orientation in space, which is irrelevant):
With observations such as this, the sense of molecular shape entered chemistry and since then has assumed a central and fundamental position.
Nineteenth-century chemists had to infer the shapes of molecules from clever but indirect experimentation. The modern understanding of molecular shape is more direct (if one discounts the computing that intervenes between observation and representation). In particular, X-ray diffraction has provided incomparably detailed images of molecules even as large as those of proteins, which contain thousands of atoms. Scanning tunneling microscopy has provided realistic images that confirm beyond doubt the essential features of molecular geometry.
The importance of the determination (and understanding) of molecular structure cannot be overestimated. At the simplest level, the properties of small molecules (including the ubiquitous and important water molecule, H2O) stem in large measure from their shapes and not merely from their atomic compositions. The oceans, for instance, might not exist if water molecules were linear rather than angular, for the interactions between H2O molecules would not be as strong, and hence it is doubtful whether life would have emerged if water molecules were linear. At the most complex level, that of proteins, geometric structure is essential to biochemical function and thus has a critical role in all living systems.
Internal structure of atoms
The concept of atoms thus emerged from the meticulous measurement of mass and volume, which in the earliest days of chemistry were the only quantitative probes of matter available. The reality of atoms was established by the explanatory power of the model on the one hand and by ever more direct images of microscopic entities on the other. As the atomic model of matter became more firmly established, attention turned to the existence of molecules, which are specific assemblages of atoms. As molecules were examined, it was discovered that they have characteristic links between atoms and that the atoms are positioned in three-dimensional arrangements that are characteristic of the compound and of the constituent atoms.
Discovery of the electron
The questions raised by this fund of knowledge remained unanswered until the internal structure of atoms began to be unraveled at the end of the 19th century. The classic view, proposed by John Dalton, that atoms are irreducible, unchangeable entities, virtually eliminated the prospect of understanding their properties, for it implied the absence of internal structure. The mutability of atoms, and hence the first glimmerings of an understanding of their constitution and their properties, came with the discovery of the electron as a universal constituent of matter. The electron was the first subatomic particle to be discovered and in due course proved to be the most important one for the explanation of the chemical bond. This importance stems in large part from the ease with which electrons can be removed from one atom and transferred to another. As will be seen below, this transferability of electrons is the key to bond formation, and all theories of the chemical bond focus on the redistribution of an atom’s electrons when it links to another atom.
More will be said about the essential features of the arrangement of electrons in atoms in the following section. The key to understanding the structure of the periodic table and hence the pattern of bonding between atoms was the realization that electrons are arranged in shells that surround a central positively charged nucleus. Each shell can contain a characteristic maximum number of electrons. The outermost shell contains the electrons that are involved in bond formation, for they are the least tightly bound to the nucleus and thus can be removed most readily. This shell is called the valence shell. The most important feature of the valence shell is that for the noble gases it is complete (in the sense explained below) with its full complement of electrons (i.e., eight, excepting the case of helium). Thus, the formation of chemical bonds appears to be related to the incompleteness of the valence shell.
Contributions of Lewis
The role of the valence shell in bond formation was expounded by the American chemist Gilbert N. Lewis about 1916. Important independent studies were made by German physicist Walther Kossel. (Lewis’s ideas about bonding are sometimes called Kossel-Lewis theory.) .Later contributions followed from American chemist Irving Langmuir. First, Lewis proposed that ionic bonds are formed by the complete transfer of electrons from the valence shell of one atom into the valence shell of another atom and that the transfer proceeds until the valence shells of both have reached the electronic composition characteristic of the nearest noble gas atom in the periodic table. Thus, sodium has one electron in its valence shell, and its loss results in a singly charged cation, Na+, with a neonlike arrangement of electrons. Chlorine, on the other hand, has a valence shell that needs one more electron to achieve the closed shell characteristic of its noble gas neighbour, argon, and so readily forms the singly charged anion Cl−. Thus, it is easy to comprehend the formation of sodium chloride as a collection of Na+ ions and Cl− ions.
Lewis proposed that a covalent bond consists of two electrons that are shared between atoms rather than being fully donated by one atom to another. He had no means of knowing why a pair of electrons should be so important (that understanding would come only later with the introduction of quantum mechanics), but his insight rationalized a great body of chemical facts. As in the formation of ionic bonds, Lewis emphasized the importance of the nearest-noble-gas valence shell and proposed that, as in the formation of ionic bonds, electron sharing continues until each atom possesses a noble gas configuration.
In summary, Lewis’s ideas are expressed by his celebrated octet rule, which states that electron transfer or electron sharing proceeds until an atom has acquired an octet of electrons (i.e., the eight electrons characteristic of the valence shell of a noble gas atom). When complete transfer occurs, the bonding is ionic. When electrons are merely shared, the bonding is covalent, and each shared electron pair constitutes one chemical bond.
Such is the basis of the theory of chemical bonding that is still widely held. There is much to explain and more to understand, however, and there are many important exceptions to Lewis’s ideas, which cannot as a consequence provide a complete explanation of bonding. The following sections step back from this historical account and put Lewis’s important ideas in a broader context that will show more of their power. At the same time the more-advanced treatment of bonding will transcend Lewis’s ideas and account for features of bonding that his views could not embrace.
