Three models are used to describe covalent bonding. The Lewis model is based on the experimental observation that atoms that contain the same number of electrons as the inert gases are especially stable. For example, chlorine, a member of Group VII, easily attracts an electron to form the chloride ion, Cl-. The chloride ion contains a total of 18 electrons, the same number as the argon atom, the third member of the inert gases (see Figure 9, the periodic chart). Likewise, the oxygen atom, which contains 8 electrons frequently gains two more electrons to produce the oxide ion, O2-. The oxide ion has the same number of electrons as neon, the second member of the inert gas group. If you look carefully at the periodic table in Figure 9 you will realize that there are eight elements in Periods (rows) two and three (if you ignore the transition metals, there are also eight elements in Periods 4, 5, and 6). The first element of each period has one electron in the outermost value of n, the principal quantum number. Thus, all of the elements in a given Group have the same number of outermost electrons. These outermost electrons are called valence electrons. Hence, all of the elements of Group IV have four valence electrons; all of the elements of Group VII have seven valence electrons. For example, fluorine has seven electrons in its outermost quantum level (n=2); chlorine has seven electrons in its outermost quantum level where n=3, bromine has seven electrons in its outermost quantum level where n=4, and so on.
Lewis concluded that most atoms have the tendency to attract eight valence electrons. If they form ionic compounds, then the atoms lose or gain enough electrons to obtain a valence shell of eight electrons. But, if they form covalent compounds, they can only attract and share a total of eight electrons. One of the simplest examples of a diatomic molecule is diatomic fluorine, the elemental form of fluorine. Because each fluorine atom has seven electrons, the two fluorine atoms in the molecule must have a total of 14 electrons. According to Lewis, these electrons must be positioned so that each fluorine atom "sees" eight electrons around it. In electron-dot (or Lewis ) formula (a) below, the seven electrons from one fluorine are designated by dots and the seven from the other fluorine are indicated by x's. Some authors prefer to use this notation, but in fact, after the molecule is formed there is no way to distinguish between the electrons. Consequently, most authors prefer to use dots for all of the electrons (b).
After the student begins to understand how to draw the electron-dot formulas, the electrons that are shared by the atoms (those between the symbols) are shown as a line or bond. Thus, in (c) the bond between the two fluorines represents two electrons.
For some molecules, each atom can only be surrounded by eight electrons if double or triple bonds are present. For example, for diatomic oxygen (the molecule that enters your bloodstream even as you read this),
formula (a) above has eight electrons in the valence shell of the right-hand oxygen, but only six electrons in the valence shell of the left hand oxygen. The electrons between the atoms, where the light blue and dark blue circles overlap, are counted in the valence shell of both atoms. Thus, electron-dot formula (a) does not obey the Lewis octet rule. In formula (b) two sets of electrons are counted in the valence shell of both atoms, and therefore both atoms have a total of eight electrons. Finally, in (c) the electron-dot formula (b) is given its usual form using two lines (bonds) to represent the four electrons between the atoms. It is also important to be aware of two other aspects of electron-dot formulas: a) the electrons are always placed in pairs, and b) the total number of electrons must add up to the total number of valence electrons in all of the atoms in the molecule or ion.
Write an electron dot formula for diatomic nitrogen, N2.
Each nitrogen has five valence electron (nitrogen is in group V on the periodic chart) and therefore there are a total of ten electrons that must be distributed around the two atoms in such a way as to give each atom eight electrons. The six electrons between the nitrogens (each line represents two electrons) count in the valence shell of both nitrogens.
As a final example of electron-dot formulas, let us consider the bonding in the nitrite ion, NO2- (sodium nitrite is used as a preservative in meats). The electron dot formula below distributes the 18 electrons (five from nitrogen, six from each of the oxygens, and one additional because of the negative charge on the ion) in a way that gives each atom an octet of electrons:
If we were to predict the structure of this ion from its electron dot formula we would say that one nitrogen-oxygen linkage has a double bond while the other has a single bond. The number of bonds is related to the length of the linkage: the greater the number of bonds, the shorter the linkage. Hence, we would conclude that the nitrite ion has an unsymmetrical structure--that one nitrogen-oxygen linkage is shorter than the other (see below).
We can determine whether this prediction is correct by determining the structure of the ion with x-ray diffraction. The experimentally-determined structure is shown below and is symmetrical; that is, each nitrogen-oxygen bond has the same length.
As good scientists we now wonder how we can bring our electron-dot model into line with the experimental fact. In other words, we try to modify our model to make it fit the facts. Linus Pauling, one of the great American chemists, suggested that this can be accomplished by simply writing another electron-dot formula and then combining the two, a process he referred to as resonance hybridization. This process is sometimes likened to a description of a rhinoceros as a hybrid of a unicorn and a dragon. Neither of the two things that are hybridized are real, but the hybrid is a real creature. The analogy is not quite accurate in that the resulting resonance hybrid of the two nitrite ions formulas is still not an accurate description of the real nitrite ion, but it is as close as we can get with the electron-dot model. The two resonance contributors are shown below with the double-headed arrow that is used only to designate resonance [we will find that chemists use a variety of arrows and we must be careful to chose the right one for the job.]
The hybridization of the two structures is similar to the process of fusing or melting them together: the resulting melt contains equal nitrogen-bond lengths and gives us a description that is consistent with the x-ray data.
Finally, we must mention the concept of formal charge. When the electrons around the oxygen with the single bond in the formula
are counted up, it turns out that there are eight as there should be if the octet rule is obeyed. But for the purpose of determining how much extra electron density the oxygen has, one of the electrons in the nitrogen-oxygen bond should be assigned to the nitrogen. This means that the oxygen has seven electrons around it, one more than the six valence electrons that it has if it is electrically neutral. Thus, the oxygen is assigned a formal charge of minus one. This is not a real charge; it is just the result of an arbitrary way of assigning electrons to atoms. It does, however, give some indication of which atoms contain more or less electron-density than they would if they were isolated atoms. As shown below, formal charge is always circled.