Exceptions to the octet rule

The metal consists of positive sodium ions in a sea of valence electrons.Valence (bonding) electrons are free to move throughout the metal crystal (beige area). 

The octet rule stems from the fact that the main-group elements in most cases employ only an ns and three np valence-shell orbitals in bonding, and these orbitals hold eight electrons. Elements of the second period are restricted to these orbitals, but from the third period on, the elements also have unfilled nd orbitals, which may be used in bonding. For example, the valence-sheU configuration of phosphorus is 

Using just these 3s and 3p orbitals, the phosphorus atom can accept only three additional electrons, forming three covalent bonds (as in PF3). However, more bonds can be formed if the empty 3d orbitals of the atom are used. If each of the five electrons of the phosphorus atom is paired with unpaired electrons of fluorine atoms, PF5 can be formed. In this way, phosphorus forms both the trifluoride and the pentafluoride. By contrast, nitrogen (which has no available d orbitals in its valence shell) forms only the trifluoride, NF3. 

Writing Lewis Formuias (Exceptions to the Octet Rule) 

a noble gas, forms a number of compounds. One of these is xenon tetra-fluoride, XeF4, a white, crystalline solid first prepared in 1962. What is the electron-dot formula of the Xep4 molecule  

The octet rule is so simple and useful in introducing the basic concepts of bonding that you might assume it is always obeyed. In Section 8.2, however, we noted its limitation in dealing with ionic compounds of the transition metals. The rule also fails in many situations involving covalent bonding. These exceptions to the octet rule are of three main types  

The octet rule almost always holds for second-period elements. Exceptions to the octet rule fall into three categories  

The octet rule has been our mainstay in writing Lewis structures, and it will continue to be one. Yet at times, we must depart from the octet rule, as we will see in this section. 

The molecule NO has 11 valence electrons, an odd number. If the number of valence electrons in a Lewis structure is odd, there must be an unpaired electron somewhere in the structure. Lewis theory deals with electron pairs and does not tell us where to put the impaired electron it could be on either the N or the O atom. To obtain a structure free of formal charges, however, we will put the unpaired electron on the N atom. 

The presence of unpaired electrons causes odd-electron species to be paramagnetic. NO is paramagnetic. Molecules with an even number of electrons are expected to have all electrons paired and to be diamagnetic. An important exception is seen in the case of O2, which is paramagnetic despite having 12 valence electrons. Lewis theory does not provide a good electronic structure for O2, but the molecular orbital theory that we will consider in the next chapter is much more successful. 

The number of stable odd-electron molecules is quite limited. More common are free radicals, or simply radicals, highly reactive molecular fragments with one or more unpaired electrons. The formulas of free radicals are usually written with a dot to emphasize the presence of an impaired electron, such as in the methyl radical, CH3, and the hydroxyl radical, OH. The Lewis structures of these two free radicals are 

Both of these free radicals are commonly encountered as transitory species in flames. In addition, OH is formed in the atmosphere in trace amounts as a result of photochemical reactions. 

We find the total number of valence electrons for all the atoms. 

We use three pairs to form single bonds between the carbon and three hydrogen atoms. 

This leaves one electron pair, which we nse as a lone pair on the nitrogen atom. 

If necessary, we use multiple bonds to satisfy the octet rule (i.e., give atoms the noble 

A diamond anvil cell used to study materials at very high pressures. 

When writing Lewis structures, do not worry about which electrons come from which atoms in a molecule. The best way to look at a molecule is to regard it as a new entity that uses all the available valence electrons of the atoms to achieve the lowest possible energy. The valence electrons belong to the molecule, rather than to the individual atoms. Simply distribute all valence electrons so that the various rules are satisfied, without regard for the origin of each particular electron. 

The localized electron model is a simple but very successful model, and the rules we have used for Lewis structures apply to most molecules. However, with such a simple model, some exceptions are inevitable. Boron, for example, tends to form compounds in which the boron atom has fewer than eight electrons around it—it does not have a complete octet. Boron trifluoride (BF3), a gas at normal temperatures and pressures, reacts very energetically with molecules such as water and ammonia that have available electron pairs 

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Note that in this structure boron has only 6 electrons around it. The octet rule for boron can be satisfied by drawing a structure with a double bond, such as 

Equivalent Lewis structures contain the same numbers of single and multiple bonds. For example, the resonance structures for 03, 

The concept of resonance is necessary because the LE model postulates that electrons are localized between a given pair of atoms. However, nature doesn t really operate this way. Electrons are actually delocalized—they can move around the entire molecule. The valence electrons in the N03 molecule distribute themselves to provide equivalent N—O bonds. Resonance is necessary to compensate for this defective assumption of the LE model. However, because this model is so useful, we retain the concept of localized electrons and add resonance to accommodate species like NO,-. 

Describe the electron arrangement in the nitrite anion (NO -), using the LE model. 

We will follow the usual procedure for obtaining the Lewis structure for the NO - ion. 

In N02 there are 5 + 2(6) + 1 = 18 valence electrons. Indicating the single bonds gives the structure 

Tie octet rule is an important guide to understanding how most compounds are formed. 

However, there are a number of cases in which the octet rule does not apply. Answer the following questions about exceptions to the octet rule. 

Does BeF2 obey the octet rule Explain. 

Not all substances conform to the octet rule. Two common oxides of nitrogen, NO and NO2, have an odd number of electrons. It is therefore impossible to write Lewis diagrams for these compounds in which each atom is surrounded by eight valence electrons. Their Lewis diagrams are drawn with one single electron in a space where an electron pair would usually appear  

Electrons first appear in c/-orbitals in the third principal energy level (Section 11.3). 

All of the bonds we have encountered thus far have been formed by the valence electrons, the highest-energy s- and p-orbital electrons in the atom. Atoms of elements in the third period and higher can have more than four electron pairs surrounding them. This is accomplished by tf-orbitals forming covalent bonds. Phosphorus pentafluoride, PF5, places five electron-pair bonds around the phosphorus atom six pairs surround sulfur in SEgi 

Two other substances for which satisfactory octet-rule diagrams can be drawn, but which are contradicted experimentally, are the fluorides of beryllium and boron. We 

Lewis theory is often correct in its predictions, but exceptions exist. For example, if we try to write the Lewis structure for NO, which has 11 electrons, the best we can do is  

The nitrogen atom does not have an octet, so this is not a great Lewis structure. However, NO exists in nature. Why As with any simple theory, Lewis theory is not sophisticated enough to be correct every time. It is impossible to write good Lewis structures for molecules with odd numbers of electrons, yet some of these molecules exist in nature. In such cases, we simply write the best Lewis structure that we can. Another significant exception to the octet rule is boron, which tends to form compounds with only six electrons around B, rather than eight. For example, BF3 and BH3—both of which exist in nature— lack an octet for B. 

A third type of exception to the octet rule is also common. A number of molecules, such as Sp6 and PCI5, have more than eight electrons around a central atom in their Lewis structures. 

These are often referred to as expanded octets. Expanded octets can form for period 3 elements and beyond. Beyond mentioning them, we do not cover expanded octets in this book. In spite of these exceptions, Lewis theory remains a powerful and simple way to understand chemical bonding. 

Which two species have the same number of lone electron pairs in their Lewis structures  

As mentioned earlier, the octet rule applies mainly to the second-period elements. Exceptions to the octet rule fall into three categories characterized by an incomplete octet, an odd number of electrons, or more than eight valence electrons around the central atom. 

In some compounds, the number of electrons surrounding the central atom in a stable molecule is fewer than eight. Consider, for example, beryllium, which is a Group 2A (and a second-period) element. The electron configuration of beryllium is I/2/ it has two valence electrons in the 2s orbital. In the gas phase, beryllium hydride (BeH2) exists as discrete molecules. The Lewis structure of BeH2 is 

As you can see, only four electrons surround the Be atom, and there is no way to satisfy the octet rule for beryllium in this molecule. 

Elements in Group 3A, particularly boron and aluminum, also tend to form compounds in which they are surrounded by fewer than eight electrons. Take boron as an example. Because its electron configuration is Is ls lp, it has a total of three valence electrons. Boron reacts with the halogens to form a class of compounds having the general formula BX3, where X is a halogen atom. Thus, in boron trifluoride there are only six electrons around the boron atom  

The following resonance structures all contain a double bond between B and F and satisfy the octet rule for boron  

Beryllium, unlike the other Group 2A elements, forms mostly covalent compounds of which BeHj is an example. 

Elements in Group 3A, particularly boron and aliuninum, also tend to form compounds in which they are surrounded by fewer than eight electrons. Take boron as an example. Because its electron configuration is it has a total of three valence 

In this stable compound boron has an octet of electrons. 

It is characteristic of boron to form molecules in which the boron atom is electron-deficient. On the other hand, carbon, nitrogen, oxygen, and fluorine do obey the octet rule. 

Some atoms appear to exceed the octet rule. This behavior is observed only for those elements in Period 3 of the periodic table and beyond. To see how this arises, we will consider the Lewis structure for sulfur hexafluoride (SFg). The sum of the valence electrons for SFg is 

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Recent studies indicate that double bonding may be important in BF3. However, the boron atom in BF3 certainly behaves as if it is electron-deficient, as indicated by the 

Atoms form bonds to make their electronic structures similar to those of noble gases. All noble gases except He have an electron structure ending with ns2 np6. Most atoms complete their valence shell with eight electrons (an octet) to become stable. However, some exceptions occur. 

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The octet rule accounts for the valences of many of the elements and the structures of many compounds. Carbon, nitrogen, oxygen, and fluorine obey the octet rule rigorously, provided there are enough electrons to go around. However, some compounds have an odd number of electrons. In addition, an atom of phosphorus, sulfur, chlorine, or another nonmetal in Period 3 and subsequent periods can accommodate more than eight electrons in its valence shell. The following two sections show how to recognize exceptions to the octet rule. 

Lewis recognized that certain molecules such a PCI5 and SF6 are exceptions to the octet rule because their Lewis structures indicate that the central atom has more than eight electrons in its valence shell 10 for the P atom in PCI5 and the S atom in SF4, and 12 for the S atom in SFg (Figure 1.17). Such molecules are called hypervalent because the valence of the central atom is greater than its principal valence. To write a Lewis structure for such molecules, the Lewis symbol for the hypervalent atom must be modified to show the correct number of unpaired electrons. For the molecules in Figure 1.17 we would need to write the Lewis symbols as follows ... 

There are also molecules that are exceptions to the octet rule because one of the atoms has fewer, rather than more than, eight electrons in its valence shell in the Lewis structure (Figure 1.19). These molecules are formed by the elements on the left-hand side of the periodic table that have only one, two, or three electrons in their valence shells and cannot therefore attain an octet by using each of their electrons to form a covalent bond. The molecules LiF, BeCl2, BF3, and AIC13 would be examples. However, as we have seen and as we will discuss in detail in Chapters 8 and 9, these molecules are predominately ionic. In terms of a fully ionic model, each atom has a completed shell, and the anions obey the octet rule. Only if they are regarded as covalent can they be considered to be exceptions to the octet rule. Covalent descriptions of the bonding in BF3 and related molecules have therefore... 

Figure 1.19 Some examples of molecules that are exceptions to the octet rule because the central atom has fewer than eight electrons in its valence shell.

Not only molecules with LLPCN > 4, but all molecules of the elements in period 3 and beyond in their higher valence states, including most of their numerous oxides, oxoacids, and related molecules such as SO3 and (H0)2S04 should be regarded as hypervalent if AO bonds are described as double bonds (1). However, Lewis did not regard these molecules as exceptions to the octet rule because he wrote the Lewis structures of these molecules with single bonds and the appropriate formal charges (2). 

The same principle applies for BeCl2, BeH2, BCl3 etc. Beryllium and boron compounds are exceptions to the octet rule. 

In a compound that is an exception to the octet rule, there is usually only one atom, other than hydrogen, that is an exception. There are few nonhydrogen compounds with more than one exception present. 

Even though we have an exception, we can still complete the Lewis structure. We need to draw a bond from each of the fluorine atoms to the central xenon. This gives us 4 bonds and uses 8 electrons. Each fluorine atom needs to complete its octet. The bond accounts for 2 electrons, so we need 6 more electrons (3 pairs) for each. Therefore, we add 3 separate pairs to each of the fluorine atoms. Six electrons per fluorine times 4 fluorine atoms accounts for 24 electrons. Our Lewis structure now contains 8 + 24 = 32 electrons. The number of available electrons (A) is 36, so we still need to add 36 - 32 = 4 electrons. These 4 electrons will give us 2 pairs. The xenon atom will get these pairs and become an exception to the octet rule. The actual placement of the pairs is not important as long as it is obvious that they are with the central atoms and not one of the fluorine atoms. The final Lewis structure is ... 

Elements in the first two rows on the periodic table will never exceed an octet. If they are exceptions to the octet rule, they will have less than an octet. 

As we have mentioned previously, there are many exceptions to the octet rule. In these cases, the N — A = S rule does not apply, as illustrated by the following example. 

This polyatomic ion has six bonds around the central Si atom, an obvious exception to the octet rule, so the central atom needs an expanded valence shell. 

The exceptions to the octet rule described in the previous section, the xenon compounds and the tri-iodide ion, are dealt with by the VSEPR and valence bond theories by assuming that the lowest energy available d orbitals participate in the bonding. This occurs for all main group compounds in which the central atom forms more than four formal covalent bonds, and is collectively known as hypervalence, resulting from the expansion of the valence shell This is referred to in later sections of the book, and the molecular orbital approach is compared with the valence bond theory to show that d orbital participation is unnecessary in some cases. It is essential to note that d orbital participation in bonding of the central atom is dependent upon the symmetry properties of individual compounds and the d orbitals. 

There are many exceptions to the octet rule—after all, it s called the octet rule, not the octet law—but it is nevertheless useful for making predictions and for providing insights about chemical bonding. 

The octet rule states that atoms tend to gain, lose, or share electrons in order to acquire eight valence electrons. Are there exceptions to the octet rule ... 

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