According to the receptor theory of drug action, the drug must interact by combining or binding with the macromolecular tissue element at the site of action in order to initiate the series of events that produce its characteristic biologic effect. Therefore, there must be some forces that not only attract the drug to its receptor, but also hold it in combination with the receptor long enough to initiate the chain of events leading to the effect. These forces are the chemical bonds that hold two atoms, groups of atoms or molecules together with sufficient stability that the combination may be considered an independent molecular species. Since these forces underlie all interactions between drugs and the tissue elements of a living system, we shall briefly consider each of the four types of bonds that may be formed.
Let us first recall some features of the structure of the atom. The internal structure of the atom consists of a nucleus that has a positive electric charge and accounts for most of the mass of the atom. The nucleus is surrounded by electrons in sufficient number that their total negative charge is equal to the positive charge of the nucleus. This makes the atom electrically neutral. The simplest atom is hydrogen, which contains a single positively charged proton in its nucleus and a single negatively charged electron moving about the nucleus. Atoms larger than hydrogen also contain neutrons in their nucleus, and these, as the name implies, carry no charge. For example, carbon has six protons and six neutrons in its nucleus and six surrounding electrons; oxygen has eight protons and eight neutrons with eight surrounding electrons. The extranuclear electrons are the seat of chemical reactivity, and this reactivity is dependent on the configuration of the external electrons.
The external electrons move about the nucleus in groups and subgroups, referred to respectively as shells and subshells, each shell and subshell having a definite number of electrons that may be situated in it. The number of shells and subshells is determined by the total number of external electrons of the atom. The simplest atoms, hydrogen and helium, have only one external shell; two is the maximum number of electrons that can be accommodated in this shell. Successive shells may contain eight or more electrons, but there can be no more than eight electrons in the outermost shell of an atom before the next shell is started. In other words, in shells that can contain more than eight electrons, only eight enter initially, then a new shell is started and the incomplete shell is left to be filled in later. This electronic configuration of eight electrons in the outermost shell (two in the case of helium) corresponds to the structure of the inert, noble or rare gases. The chemical inactivity of these gases indicates that their configuration must be a highly stable arrangement of electrons.
All atoms try to reach chemical stability and to attain the configuration of a rare gas.
They do this by giving up or taking on electrons. For example, the sodium atom, with eleven external electrons, has two closed shells, the first with two electrons corresponding to helium, the second with eight electrons corresponding to neon and the remaining electron in the third outermost shell. When the sodium atom gives up this electron to reach the configuration of neon, the resulting particle is positively charged because it now has one more proton in its nucleus than it has electrons in its surrounding shells. Chlorine, with seventeen electrons, has seven of these in its outermost shell. When chlorine accepts an electron to attain the configuration of the inert gas argon, it also loses its neutrality and becomes negatively charged. Charged particles are called ions. Thus, chemical reactivity is associated primarily with the electrons in the outermost shell.
THE IONIC BOND. The atoms of metallic elements, such as sodium, tend to give up their electrons easily, whereas the nonmetallic atoms, such as chlorine, tend to add electrons. These natural tendencies come into play when a metallic atom and a nonmetallic atom approach one another to form a stable molecule or crystal. In the electronic formulation of the interaction, the symbol of the element, i.e. Na for sodium, Cl for chlorine, represents the kernel of the atom, standing for the nucleus and the closed shells of electrons. The electrons of the outermost shell, the valence shell, are shown by dots, thus:
An electron is transferred from the Na to the Cl to yield a positively charged sodium cation and a negatively charged chloride anion, each with the configuration of a rare gas.
These ions are stable and retain their electronic configuration essentially independently of each other in solution and even in the solid, crystalline form. However, they are held together to form a molecule by the electrostatic attraction between them. The bond formed between atoms involving the outright transfer of one or more electrons from one atom to the other is called the ionic bond. The ionic bond is defined, then, as the electro static attraction between oppositely charged ions. The strength of this bond depends on the distance between the two ions and diminishes as the square of the distance between them.
THE COVALENT BOND. Whereas electrostatic attraction can explain the binding found in a simple salt like sodium chloride, or in a more complicated molecule such as that formed by heparin with a protein of the coagulation process, this type of binding can hardly account for the formation of molecules such as the gases hydrogen, H2, or methane, CH4, in which no ions can be detected. To explain this latter type of binding, G.N.Lewis proposed in 1916 that not only can a bond between atoms arise from outright transfer of electrons, but a rare gas configuration can also be attained from the sharing of a pair of electrons by the two bonded atoms. Thus we can write electronic structures such as
in which a pair of electrons held jointly by two atoms is doing double duty and is effective in completing a stable electronic configuration for each atom. Each hydrogen
can claim the shared pair and thereby attain the structure of helium. And the carbon, by being able to share in its own as well as in the four acquired electrons of hydrogen, has the configuration of the rare gas neon. The bond formed when two atoms share a pair of electrons is known as a covalent bond. It is about twenty times stronger than the ionic bond and is responsible for the chemical stability of organic molecules.
Covalent bonding resulting from electron sharing also accounts for the formation of double and triple bonds in molecules. Moreover, a covalent bond can result from the sharing of electrons supplied by one atom only, the resulting bond being called a coordinate covalent bond. The atom contributing the electron pair is called the donor atom and in biologic systems is usually nitrogen, oxygen or sulfur. Consider the formation of ammonium ion (NH4+) from ammonia (NH3) and hydrogen ion:
The ammonia molecule has four electron pairs, of which only three are shared; thus each hydrogen has attained a stable helium configuration, and the nitrogen atom with five electrons in its outermost shell has attained its complete octet. When a hydrogen ion approaches the ammonia, the nitrogen allows the hydrogen proton to share with it the free pair of electrons. But the positive charge associated with the hydrogen ion is retained by the ammonium ion complex, since there has been no net gain or loss of electrons. We shall see that this coordinate covalent bond formation is important in the ionization of drugs and in certain interactions with receptors.
THE HYDROGEN BOND. The hydrogen atom, with only a single electron, can form only one covalent or one ionic bond with another atom. The hydrogen nucleus, however, being a bare proton, is strongly electropositive. When hydrogen is bound by an ionic or covalent bond to a strongly electronegative atom, the hydrogen may further coordinate two more electrons donated by another strongly electronegative atom, such as oxygen (O), nitrogen (N) or fluorine (F). This second bond of hydrogen is referred to as the hydrogen bond, and it forms a bridge between two strongly electronegative groups.
Hydrogen bonding may form this bridge between different molecules or may lead to the association between like molecules, as in acetic acid:
The dotted lines represent hydrogen bonds; the solid lines are covalent bonds.
Although the hydrogen bond is ionic in character, its strength is less than that of a true ionic bond. A single hydrogen bond confers little stability on the association of two molecules, but several such hydrogen bonds can stabilize an interaction significantly. For example, in a compound like water, which can form hydrogen bonds readily, the hydrogen bonding does not stop at two molecules but may extend throughout the entire mass:
This association between molecules of water in the liquid state makes it more resistant to disruption by heat. This is why water has a higher boiling point than organic solvents such as chloroform, CHCl3, which do not exist as associated liquids. Chloroform cannot form hydrogen bonds, since the hydrogen is attached by a covalent bond to carbon, which is not a strongly electronegative atom.
VAN DER WAALS FORCES. These are very weak attractive forces between any two neutral atoms or atomic groupings; they operate only at close range. Since the force of attraction is inversely proportional to the seventh power of the distance between the atoms, these forces decrease rapidly with a slight increase in interatomic distance.