Considering the tremendous variety in properties of elements and compounds in the periodic system, it is asking a great deal to expect a rule as simple as Lewis’ octet theory to be able to predict all formulas or to account for all molecular structures involving covalent bonds. Lewis’ theory concentrates on resemblances to noble-gas ns2np6 valence octets. Therefore it is most successful in accounting for formulas of compounds of the representative elements, whose distinguishing electrons are also s and p electrons. The octet rule is much less useful in dealing with compounds of the transition elements or inner transition elements, most of which involve some participation of d or f orbitals in bonding.
Even among the representative elements there are some exceptions to the Lewis theory. These fall mainly into three categories:
- Some stable molecules simply do not have enough electrons to achieve octets around all atoms. This usually occurs in compounds containing Be or B.
- Elements in the third period and below can accommodate more than an octet of electrons. Although elements such as Si, P, S, Cl, Br, and I obey the octet rule in many cases, under other circumstances they form more bonds than the rule allows.
- Free Radicals
Electron Deficient Species
Good examples of the first type of exception are provided by BeCl2 and BCl3. Beryllium dichloride, BeCl2, is a covalent rather than an ionic substance. Solid BeCl2 has a relatively Complex structure at room temperature, but when it is heated to 750°C, a vapor which consists of separate BeCl2 molecules is obtained. Since Cl atoms do not readily form multiple bonds, we expect the Be atom to be joined to each Cl atom by a single bond. The structure is
Instead of an octet the valence shell of Be contains only two electron pairs. Similar arguments can be applied to boron trichloride, BCl3, which is a stable gas at room temperature. We are forced to write its structure as
in which the valence shell of boron has only three pairs of electrons. Molecules such as BeCl2 and BCl3 are referred to as electron deficient because some atoms do not have complete octets. Electron-deficient molecules typically react with species containing lone pairs, acquiring octets by formation of coordinate covalent bonds. Thus BeCl2 reacts with Cl– ions to form BeCl4–;
BCl3 reacts with NH3 in the following way:
Species with Expanded Octets
Examples of molecules with more than an octet of electrons are phosphorus pentafluoride (PF5) and sulfur hexafluoride (SF6). Phosphorus pentafluoride is a gas at room temperature. It consists of PF5 molecules in which each fluorine atom is bonded to the phosphorus atom. Since each bond corresponds to a shared pair of electrons, the Lewis structure is
Instead of an octet the phosphorus atom has 10 electrons in its valence shell. Sulfur hexafluoride (also a gas) consists of SF6 molecules. Its structure is
Here the sulfur atom has six electron pairs in its valence shell.
An atom like phosphorus or sulfur which has more than an octet is said to have expanded its valence shell. This can only occur when the valence shell has enough orbitals to accommodate the extra electrons. For example, in the case of phosphorus, the valence shell has a principal quantum number n = 3. An octet would be 3s23p6. However, the 3d subshell is also available, and some of the 3d orbitals may also be involved in bonding. This permits the extra pair of electrons to occupy the valence (n = 3) shell of phosphorus in PF5.
Expansion of the valence shell is impossible for an atom in the second period because there is no such thing as a 2d orbital. The valence (n = 2) shell of nitrogen, for example, consists of the 2s and 2d subshells only. Thus nitrogen can form NF3 (in which nitrogen has an octet) but not NF5. Phosphorus, on the other hand, forms both PF3 and PF5, the latter involving expansion of the valence shell to include part of the 3d subshell.
The majority of molecules or complex ions discussed in general chemistry courses are demonstrated to have pairs of electrons. However, there are a few stable molecules which contain an odd number of electrons. These molecules, called "free radicals", contain at least one unpaired electron, a clear violation of the octet rule. Free radicals play many important roles a wide range of applied chemistry fields, including biology, medicine, and astrochemistry.
Three well-known examples of such molecules are nitrogen (II) oxide, nitrogen(IV) oxide, and chlorine dioxide. The most plausible Lewis structures for these molecules are
Free radicals are usually more reactive than the average molecule in which all electrons are paired. In particular they tend to combine with other molecules so that their unpaired electron finds a partner of opposite spin. Since most molecules have all electrons paired, such reactions usually produce a new free radical. This is one reason why automobile emissions which cause even small concentrations of NO and NO2 to be present in the air can be a serious pollution problem. When one of these free radicals reacts with other automobile emissions, the problem does not go away. Instead a different free radical is produced which is just as reactive as the one which was consumed. To make matters worse, when sunlight interacts with NO2, it produces two free radicals for each one destroyed:
In this way a bad problem is made very much worse.
A fourth very interesting example of a free radical is oxygen gas. The Lewis structure for Oxygen usually hides the fact that it is a "diradical", containing two unpaired electrons. This is often cited as a serious flaw in Lewis bond theory, and was a major impetus for development of molecular orbital theory. We know oxygen is a diradical because of its paramagnetic character, which is easily demonstrated by attraction of oxygen to an external magnet.