Atomic orbitals : Resonance structures

Here we revisit the VB resonance model in formamide by taking all six resonance structures into account. Such a study allows us to compare the individual contributions from resonance structures 1-6 to the resonance effect in formamide. For comparison, the isoelectronic systems vinylamine and formamidine are also investigated to gain insights into the trends of resonance stabilization. A 6-31G(d) basis set is employed in the calculations, and the orbitals in the VB functions are self-consistently determined for each resonance structure, but restricted to be atomic orbitals. The structural weights of the six resonance structures are listed in Table 2. 

All resonance structures must have identical geometries. Otherwise they do not represent the same molecule. For example, the following structure (known as Dewar benzene) is not a resonance form of benzene because it is not planar and has two less ir electrons. Because molecular geometry is linked to hybridization, it follows that hybridization also is unchanged for the atoms in resonance structures. (Note If it is assumed that the central bond in this structure is a bond, then it has the same number of electrons as benzene. However, in order for the p orbitals to overlap, the central carbon atoms would have to be much closer than they are in benzene, and this is yet another reason why Dewar benzene is an isolable compound rather than a resonance form of benzene.)... 

To describe the o bonds in benzene, begin with a Lewis structure and count the electron domains on the carbon atoms. (Either resonance structure will give the same result.) Each C atom has three electron domains around it (two single bonds and one double bond). Recall from Table 9.4 that an atom that has three electron domains is sf hybridized. To obtain the three unpaired electrons necessary on each C atom, one electron from each C atom must be promoted If ran the doubly occupied 2s orbital to an empty 2p orbital ... 

Each of the increased-valence structures of Fig. 13-4 involves an increased-valence bonding unit of type (17) for the six rc-electrons. It summarizes resonance between the standard Lewis structure (18) and the two long-bond Lewis stractures (19) and (20). (Because structure (21) involves three electrons located in the A-atom atomic orbital, this structure cannot be included in the Lewis structure resonance scheme.)... 

Some fundamental structure-stability relationships can be employed to illustrate the use of resonance concepts. The allyl cation is known to be a particularly stable carbocation. This stability can be understood by recognizing that the positive charge is delocalized between two carbon atoms, as represented by the two equivalent resonance structures. The delocalization imposes a structural requirement. The p orbitals on the three contiguous carbon atoms must all be aligned in the same direction to permit electron delocalization. As a result, there is an energy barrier to rotation about the carbon-carbon... 

Molecular orbitals are useful tools for identifying reactive sites in a molecule. For example, the positive charge in allyl cation is delocalized over the two terminal carbon atoms, and both atoms can act as electron acceptors. This is normally shown using two resonance structures, but a more compact way to see this is to look at the shape of the ion s LUMO (the LUMO is a molecule s electron-acceptor orbital). Allyl cation s LUMO appear s as four surfaces. Two surfaces are positioned near- each of the terminal car bon atoms, and they identify allyl cation s electron-acceptor sites. 

Look closely at the acid-base reaction in Figure 2.5, and note how it is shown. Dimethyl ether, the Lewis base, donates an electron pair to a vacant valence orbital of the boron atom in BF3, a Lewis acid. The direction of electron-pair flow from the base to acid is shown using curved arrows, just as the direction of electron flow in going from one resonance structure to another was shown using curved arrows in Section 2.5. A cuived arrow always means that a pair of electrons moves from the atom at the tail of the arrow to the atom at the head of the arrow. We ll use this curved-arrow notation throughout the remainder of this text to indicate electron flow during reactions. 

As the following resonance structures indicate, enamines are electronically similar to enolate ions. Overlap of the nitrogen lone-pair orbital with the double-bond p orbitals leads to an increase in electron density on the a carbon atom, making that carbon nucleophilic. An electrostatic potential map of N,N-6imethyl-aminoethvlene shows this shift of electron density (red) toward the a position. 

As pointed out in Chapter 7, the atomic orbital (valence bond) model regards benzene as a resonance hybrid of the two structures... 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

Never draw a carbon atom with more than fonr bonds. This is a big no-no. Carbon atoms only have fonr orbitals therefore, carbon atoms can form only fonr bonds (bonds are formed when orbitals of one atom overlap with orbitals of another atom). This is true of all second-row elements, and we discuss this in more detail in the chapter on drawing resonance structures. 

Inspection of the second resonance structure reveals that this nitrogen atom is actually sp hybridized, not sp. It might look like it is sp hybridized in the first resonance structure, but it isn t. Here is the general rule a lone pair that participates in resonance must occupy ap orbital. In other words, the nitrogen atom in the compound above is sp hybridized. And as a result, this nitrogen atom is trigonal planar rather than trigonal pyramidal. 

A clue to the nature of the third itt MO can be found in the placement of electrons in the two resonance structures for ozone, which are shown with color highlights in Figure 10-36a. Notice that in one resonance structure, the left outer atom has three lone pairs and a single bond, while the right outer atom has two lone pairs and a double bond. In the other resonance structure, the third lone pair is on the right outer atom, with the double bond to the left outer atom. The double bond appears in different positions in the two stmctures, and one of the lone pairs also appears in different positions. These variations signal delocalized orbitals. 

Triatomic species can be linear, like CO2, or bent, like O3. The principles of orbital overlap do not depend on the identity of the atoms involved, so all second-row triatomic species with 16 valence electrons have the same bonding scheme as CO2 and are linear. For example, dinitrogen oxide (N2 O) has 16 valence electrons, so it has an orbital configuration identical to that of CO2. Each molecule is linear with an inner atom whose steric number is 2. As in CO2, the bonding framework of N2 O can be represented with sp hybrid orbitals. Both molecules have two perpendicular sets of three tt molecular orbitals. The resonance structures of N2 O, described... 

Follow the four-step procedure for the composite model of bonding. Use localized bonds and hybrid orbitals to describe the bonding framework and the inner atom lone pairs. Next, analyze the system, paying particular attention to resonance structures or conjugated double bonds. Finally, make sure the bonding inventory accounts for all the valence electrons and all the valence orbitals. 

Resonance structures result from a phenomenon known as electron delocalization. The electron pairs in the three double bonds in a benzene ring are delocalized. These are electrons that belong to no particular atom or bond. As a consequence, no ordinary double bonds exist in a benzene ring. The electrons are in an orbital that extends across adjacent atoms. This smear of electrons is usually represented as a circle within the ring. 

It is difficult to give a localized orbital description of the bonding in a period 3 hypervalent molecule that is based only on the central atom 3s and 3p orbitals and the ligand orbitals, that is, a description that is consistent with the octet rule. One attempt to do this postulated a new type of bond called a three-center, four-electron (3c,4e) bond. We discuss this type of bond in Box 9.2, where we show that it is not a particularly useful concept. Pauling introduced another way to describe the bonding in these molecules, namely, in terms of resonance structures such as 3 and 4 in which there are only four covalent bonds. The implication of this description is that since there are only four cova-... 

In the nonbonding orbital two electrons are delocalized over the two fluorine atoms and do not contribute to the bonding, which is due only to the two electrons in the bonding orbital. This type of 3c, 4e bond is often denoted by a dashed line, as shown in Figure 2. Each P—F bond is effectively a half-bond, so this description of the bonding is roughly equivalent to the two resonance structures 1 and 2 ... 

The double bond that is shown in each of the two structures just shown is not localized as is reflected by the two resonance structures. However, the two single bonds and the unshared pair are localized as a result of the hybrid orbitals in which they reside. The hybrid orbital type is sp2, which accounts for the bond angle being 119.5°. There is one p orbital not used in the hybridization that is perpendicular to the plane of the molecule, which allows for the tv bonding to the two oxygen atoms simultaneously. The n bond is described as being delocalized, and this can be shown as follows ... 

It is readily apparent that the three a bonds are capable of holding the six bonding electrons in the a t and e molecular orbitals. The possibility of some 7r-bonding is seen in the molecular orbital diagram as a result of the availability of the a2" orbital, and in fact there is some experimental evidence for this type of interaction. The sum of the covalent radii of boron and fluorine atoms is about 152 pm (1.52 A), but the experimental B-F bond distance in BF3 is about 129.5 pm (1.295 A). Part of this "bond shortening" may be due to partial double bonds resulting from the 7r-bonding. A way to show this is by means of the three resonance structures of the valence bond type that can be shown as... 

---