^{Development of Chemical Bonding Theory}

By the mid-1800s, the new science of chemistry was developing rapidly and chemists had begun to probe the forces holding compounds together. In 1858, August Kekulé and Archibald Couper independently proposed that, in all organic compounds, carbon is tetravalent—it always forms four bonds when it joins other elements to form stable compounds. Furthermore, said Kekulé, carbon atoms can bond to one another to form extended chains of linked atoms. In 1865, Kekulé provided another major advance when he suggested that carbon chains can double back on themselves to form rings of atoms.

    Although Kekulé and Couper were correct in describing the tetravalent nature of carbon, chemistry was still viewed in a two-dimensional way until 1874. In that year, Jacobus van ’t Hoff and Joseph Le Bel added a third dimension to our ideas about organic compounds when they proposed that the four bonds of carbon are not oriented randomly but have specific spatial directions.

  Van ’t Hoff went even further and suggested that the four atoms to which carbon is bonded sit at the corners of a regular tetrahedron, with carbon in the center.

A representation of a tetrahedral carbon atom is shown in Figure 1 Note the conventions used to show three-dimensionality: solid lines represent bonds in the plane of the page, the heavy wedged line represents a bond coming out of the page toward the viewer, and the dashed line represents a bond receding back behind the page, away from the viewer. These representations will be used throughout the text.

Figure 1 A representation of a tetrahedral carbon atom. The solid lines represent bonds in the plane of the paper, the heavy wedged line represents a bond coming out of the plane of the page, and the dashed line represents a bond going back behind the plane of the page.

Why, though, do atoms bond together, and how can bonds be described electronically? The why question is relatively easy to answer: atoms bond together because the compound that results is more stable and lower in energy than the separate atoms. Energy—usually as heat—always flows out of the chemical system when a bond forms. Conversely, energy must be put into the chemical system to break a bond. Making bonds always releases energy, and breaking bonds always absorbs energy. The how question is more difficult. To answer it, we need to know more about the electronic properties of atoms.

  We know through observation that eight electrons (an electron octet) in an atom’s outermost shell, or valence shell, impart special stability to the noblegas elements in group 8A of the periodic table: Ne (2 + 8); Ar (2 + 8 + 8); Kr (2+ 8 + 18 +8). We also know that the chemistry of main-group elements is governed by their tendency to take on the electron configuration of the nearest noble gas. The alkali metals in group 1A, for example, achieve a noble-gas configuration by losing the single s electron from their valence shell to form a cation, while the halogens in group 7A achieve a noble-gas configuration by gaining a p electron to fill their valence shell and form an anion. The resultant ions are held together in compounds like Na⁺ Cl^- by an electrostatic attraction that we call an ionic bond. But how do elements closer to the middle of the periodic table form bonds? Look at methane, CH₄, the main constituent of natural gas, for example. The bonding in methane is not ionic because it would take too much energy for carbon (1s² 2s² 2p²) either to gain or lose four electrons to achieve a noble-gas configuration. As a result, carbon bonds to other atoms, not by gaining or losing electrons, but by sharing them. Such a shared-electron bond, first proposed in 1916 by G. N. Lewis, is called a covalent bond. The neutral collection of atoms held together by covalent bonds is called a molecule.

A simple way of indicating the covalent bonds in molecules is to use what are called Lewis structures, or electron-dot structures, in which the valenceshell electrons of an atom are represented as dots. Thus, hydrogen has one dot representing its 1s electron, carbon has four dots (2s2 2p2), oxygen has six dots (2s² 2p⁴), and so on. A stable molecule results whenever a noble-gas configuration is achieved for all the atoms—eight dots (an octet) for main-group atoms or two dots for hydrogen. Simpler still is the use of Kekulé structures, or linebond structures, in which a two-electron covalent bond is indicated as a line drawn between atoms.

The number of covalent bonds an atom forms depends on how many additional valence electrons it needs to reach a noble-gas configuration. Hydrogen has one valence electron (1s) and needs one more to reach the helium configuration (1s²), so it forms one bond. Carbon has four valence electrons (2s² 2p²) and needs four more to reach the neon configuration (2s² 2p⁶), so it forms four bonds. Nitrogen has five valence electrons (2s² 2p³), needs three more, and forms three bonds; oxygen has six valence electrons (2s² 2p⁴), needs two more, and forms two bonds; and the halogens have seven valence electrons, need one more, and form one bond.

  Valence electrons that are not used for bonding are called lone-pair electrons, or nonbonding electrons. The nitrogen atom in ammonia, NH₃, for instance, shares six valence electrons in three covalent bonds and has its remaining two valence electrons in a nonbonding lone pair. As a time-saving shorthand, nonbonding electrons are often omitted when drawing line-bond structures, but you still have to keep them in mind since they’re often crucial in chemical reactions.