1.1     Preview

The picture of atoms that all scientists had in their collective mind’s eye at the beginning of the past century was little different from that of the ancient Greek philosopher Democritus (~460–370 B.C.), who envisioned small, indivisible particles as the constituents of matter. These particles were called atoms, from the Greek word for indivisible. The English chemist John Dalton (1766–1844) had the idea that different atoms might have different characteristic masses, but he did not abandon the notion of a solid, uniform atom. That picture did not begin to change dramatically until 1897, when another English physicist, J. J. Thomson (1856–1940), discovered the negatively charged elementary particle called the electron. Thomson postulated a spongelike atom with the negatively charged electrons embedded within a positively charged material, rather like raisins in a pudding. In 1909, Ernest Rutherford’s (1871–1937) discovery that an atom was mostly empty space demolished the pudding picture and led to his celebrated planetary model of the atom, in which electrons were seen as orbiting a compact, positively charged nucleus, the core of positively charged protons and neutral neutrons at the center of the atom.

It was Niels Bohr who made perhaps the most important modification of the planetary model. He made the brilliant and largely intuitive² suggestion that electrons were required to occupy only certain orbits. Because an electron’s energy depends on the distance of its orbit from the positively charged nucleus, Bohr’s suggestion amounted to saying that electrons in atoms can have only certain energies. An electron might have energy x or energy 2x but nothing in between. Whenever a property is restricted to certain values in this way, we say the property is quantized. Although this notion may seem strange, there are similar phenomena in the everyday world. You cannot create just any tone by blowing across the mouth of a bottle, for example. Say you start by blowing gently and creating a given tone. The tone you hear depends on the size and shape of the particular bottle, but if you gradually increase how hard you blow, which gradually increases the energy you are supplying, the tone does not change smoothly. Instead, you hear the first tone unchanged over a certain range of energy input then a sudden change in tone when just the right “quantum” of energy has been provided.

In the 1920s and 1930s, several mathematical descriptions emerged from the need to understand Bohr’s quantum model of the atom. It became clear that we must take a probabilistic view of the subatomic world. Werner Heisenberg discovered that it is not possible to determine simultaneously both the position and momentum (mass times speed) of an electron.³ Thus, we can determine where an electron is at any given time only in terms of probability. We can say, for example, that there is a 90% probability of finding the electron in a certain volume of space, but we cannot say that at a given instant the electron is at a particular point in space.

The further elaboration of this picture of the atom has given us the conceptual basis for all modern chemistry: the idea of the orbital. Loosely speaking, an orbital describes the region of space surrounding an atomic nucleus that may be occupied by either an electron or a pair of electrons of a certain energy. Both the combining of atoms to form molecules and the diverse chemical reactions these molecules undergo involve, at a fundamental level, the interactions of electrons in orbitals. This notion will appear throughout this book; it is the most important unifying principle of organic chemistry. In atoms, we deal with atomic orbitals, and in molecules, we deal with molecular orbitals.

Various graphic conventions are used in this book to represent atoms and molecules (letters for atoms, dots for electrons not involved in bonding, and lines for electrons in bonds), but it is important to keep in mind from the outset that the model that most closely approximates our current understanding of reality at the atomic and molecular level is the cloudy, indeterminate—you might even say poetic—image of the orbital.⁴

There is great conceptual overlap between the concept of an orbital and the notion you probably encountered in general chemistry of shells of electrons surrounding the atomic nucleus. For example, you are accustomed to thinking of the noble gas elements as having filled shells of electrons: two electrons for helium in the first shell, two electrons in the first shell and eight in the second shell for neon, and so on. In the noble gases, the outermost, or valence, shells are filled. We will speak of those valence shells as valence orbitals. We shall say much more about orbitals in a moment, especially about their shapes, but the point to “get” here is the move from the old word shell to the new word orbital.

 

²Some people’s intuitions are better able than others’ to cope with the unknown!

³This idea is extraordinarily profound—and troubling. The Heisenberg uncertainty principle (which states that the product of the uncertainty in position times the uncertainty in momentum is a constant) seems to limit fundamentally our access to knowledge. For an exquisite exposition of the human consequences of the uncertainty principle, see Jacob Bronowski, The Ascent of Man, Chapter 11 (Little Brown, New York, 1973).

⁴In the wonderful quote that opens this chapter, Niels Bohr points out that once we transcend the visible world, all that is possible is modeling or image making. To us, what is even more marvelous about the quote is the simple word too: It was obvious to Bohr that scientists speak in images, and he was pointing out to Heisenberg that there was another group of people out there in the world who did the same thing—poets.