Lewis structures valence shells

Skill 21.2 Apply the concepts of Lewis structures, valence-shell electron-pair repulsion, and hybridization to describe molecular geometry and bonding. 

We now turn from the use of quantum mechanics and its description of the atom to an elementary description of molecules. Although most of the discussion of bonding in this book uses the molecular orbital approach to chemical bonding, simpler methods that provide approximate pictures of the overall shapes and polarities of molecules are also very useful. This chapter provides an overview of Lewis dot structures, valence shell electron pair repulsion (VSEPR), and related topics. The molecular orbital descriptions of some of the same molecules are presented in Chapter 5 and later chapters, but the ideas of this chapter provide a starting point for that more modem treatment. General chemistry texts include discussions of most of these topics this chapter provides a review for those who have not used them recently. 

The structure shown is the best (most stable) Lewis structure for methyl nitrite All atoms except hydro gen have eight electrons (shared + unshared) in their valence shell... 

When writing a Lewis structure we restrict a molecule s electrons to certain well defined locations either linking two atoms by a covalent bond or as unshared electrons on a sm gle atom Sometimes more than one Lewis structure can be written for a molecule espe cially those that contain multiple bonds An example often cited m introductory chem istry courses is ozone (O3) Ozone occurs naturally m large quantities m the upper atmosphere where it screens the surface of the earth from much of the sun s ultraviolet rays Were it not for this ozone layer most forms of surface life on earth would be dam aged or even destroyed by the rays of the sun The following Lewis structure for ozone satisfies fhe ocfef rule all fhree oxygens have eighf elecfrons m fheir valence shell... 

Section 1 3 The most common kind of bonding involving carbon is covalent bond ing A covalent bond is the sharing of a pair of electrons between two atoms Lewis structures are written on the basis of the octet rule, which limits second row elements to no more than eight electrons m their valence shells In most of its compounds carbon has four bonds... 

Examine electrostatic potential maps for potassium hydride and hydrogen chloride. How are they similar and how are they different (Focus on whether the molecules are polar or nonpolar (compare dipole moments), and on the electronic character of hydrogen.) Draw the ionic Lewis structure that is most consistent with each electrostatic potential map. Does each atom have a filled valence shell ... 

To check on the validity of a Lewis structure, verify that each atom has an octet or a duplet. As we shall see in Section 2.10, a common exception to this rule arises when the central atom is an atom of an element in Period 3 or higher. Such an atom can accommodate more than eight electrons in its valence shell. Consequently, the most stable Lewis structure may be one in which the central atom has more than eight electrons. 

EXAMPLE 2.7 Writing a Lewis structure with an expanded valence shell... 

The fluoride SF4 forms when a mixture of fluorine and nitrogen gases is passed over a film of sulfur at 275°C in the absence of oxygen and moisture. Write the Lewis structure of sulfur retrafluoride and give the number of electrons in the expanded valence shell. 

The Lewis structures encountered in Chapter 2 are two-dimensional representations of the links between atoms—their connectivity—and except in the simplest cases do not depict the arrangement of atoms in space. The valence-shell electron-pair repulsion model (VSEPR model) extends Lewis s theory of bonding to account for molecular shapes by adding rules that account for bond angles. The model starts from the idea that because electrons repel one another, the shapes of simple molecules correspond to arrangements in which pairs of bonding electrons lie as far apart as possible. Specifically ... 

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... 

Linnett used the concept that an octet of valence shell electrons consists of two sets of four opposite-spin electrons to show that in diatomic and other linear molecules the two tetrahedra are not in general formed into four pairs as we have discussed for F2 and the CC triple bond in C2H2. This idea is the basis of the double-quartet model, which Linnett applied to describe the bonding in a variety of molecules. It is particularly useful for the description of the bonding in radicals, including in particular the oxygen molecule, which has two unpaired electrons and is therefore paramagnetic This unusual property is not explained by the Lewis structure... 

Lewis considered covalent and ionic bonds to be two extremes of the same general type of bond in which an electron pair is shared between two atoms contributing to the valence shell of both the bonded atoms. In other words, in writing his structures Lewis took no account of the polarity of bonds. As we will see much of the subsequent controversy concerning hypervalent molecules has arisen because of attempts to describe polar bonds in terms of Lewis structures. 

In a Lewis structure a shared pair denoted by a bond line counts as contributing to the valence shell of both atoms, so that both atoms acquire an octet of electrons. Once we have introduced the concepts of a polar bond and unequal sharing of a pair of electrons, the meaning of the octet rule becomes less clear. The conventional Lewis structure of CF4 (6) obeys the octet rule, but structures 7 and 8, which would be used to describe the polarity of the bonds, do not. 

The use of resonance structures such as 7 and 8 to describe bond polarity led to a subtle change in the meaning of the octet rule, namely, that an atom obeys the octet rule if it does not have more than eight electrons in its valence shell. As a result, resonance structures such as 7 and 8 are considered to be consistent with the octet rule. However, this is not the sense in which Lewis used the octet rule. According to Lewis, a structure such as 7 would not obey the octet rule because there are only three pairs of electrons in the valence shell of carbon, just as BF3 does not obey the octet rule for the same reason. Clearly the octet rule as defined by Lewis is not valid for hypervalent molecules, which do, indeed, have more than four pairs of shared electrons in the valence shell of the central atom. 

B The Lewis structure for NI3 is similar to that of NH3. The central nitrogen atom is attached to each iodine atom by a single covalent bond. All of the atoms in this structure get a closed valence shell. 

We begin by writing the Lewis structure. The H atoms are terminal atoms. There are three central atoms and (3 x 1) + 4 + 6 + 4 + (3 x 1) = 20 valence electrons, or 10 pairs. A plausible Lewis structure is drawn at right. Each central atom is surrounded by four electron pairs, requiring sp3 hybridization. The valence-shell orbital diagrams for the atoms follow. 

Table 4.15 summarizes optimized bond lengths and NBO Lewis-like structures for 15 saturated normal-valent H MMH compounds (M = W-Pt) as well as corresponding hydrides of hypovalent Ta for comparison. The accuracy of the localized Lewis-like description (as measured by %pf) is found to be reasonably high both for normal-valent and for hypovalent species, typically >98% of the valence-shell density and >99.5% of the total density. 

In summary, we can say that, because of the unique absence of angular and radial nodes in the H-atom valence shell, the hydride oah orbital is uniquely suited to strong n-a donor-acceptor interactions with Lewis bases. In turn, the unique energetic and angular features of nB-aAH interactions (or equivalently, of B H—A <—> B—H+ A covalent-ionic resonance) can be directly associated with the distinctive structural and spectroscopic properties of B - H—A hydrogen bonding. 

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