Estrutura e Ligação em Moléculas Orgânicas

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12 Chapter 1 STRUCTURE AND BONDING IN ORGANIC MOLECULES trigonal planar structure because the lone pair, being attached to only one atom, exerts more repulsion than do the bonded pairs, closing the bond angle down a little bit. Now let us examine nitrogen dioxide. The N now bears a single nonbonding electron rather than a lone pair. One electron exerts less repulsion than do two, so we can predict that the bond angle should be more open in nitrogen dioxide than it is in nitrite. You do not have enough information to predict how much more open the angle will be. In actual fact it is 134°. The fact that it is larger than 120° means that the two bonding pairs exert more repulsion than does the single nonbonded electron. You'll no doubt be thrilled to find out that nitrogen dioxide is a significant component of urban atmos- pheric smog. Smoggy air derives much of its distinctive character from this toxic, smelly, brownish gas. (e) Now we compare two new dioxides, SO₂ and CIO₂, with the one we've already done, NO₂. Lewis structures and resonance forms first: + :0=Cl=0: The structures on the extreme right both have expanded valence shells (more than octets), which is OK for third-row atoms. Based on VSEPR both SO₂ and CIO₂ will be bent structures, because of the lone pair on S and the lone pair + extra single unshared electron on The actual angle in SO₂ is 129°, and that in CIO₂ is 116°, the difference being due to the extra repulsion of the third nonbonding electron on despite the fact that it is smelly, toxic, and tends to blow up, is actually a major industrial chemical used to bleach wood pulp in the manufacture of paper. Prudently, it is prepared just before it is used. eliminating the need to store the stuff. 34. (a) The molecular orbitals are obtained as follows: (Antibonding) Is Is (Bonding) Therefore, the resulting electronic configurations are H₂, (σ)², with two bonding electrons vs. (σ)¹, with one bonding electron. So H₂ possesses the stronger bond. (b) Same as Exercise 1-14. (c) and (d) We prepare an orbital diagram similarly. How do we begin? First, out of all the various types of orbitals presented in the chapter, let's see which ones we do and do not have to consider. In multiply-bonded molecules with many atoms, such as ethene and ethyne (Figure 1-21, textbook page 35), we needed to invoke hybrid orbitals, because we had to use them to explain geometry. In diatomic molecules such as O₂ and N₂, however, there is no "geometry" to explain, so orbital hybridization has no purpose, and we can just use simple atomic orbitals. That's good-it makes life simpler. We note also that the Is and 2s orbitals in 0 and N are completely filled. In cases such as these, it is customary to ignore the orbitals, because their overlap will produce no net bonding (just like between two atoms of He)-another welcome simplification. We're down to considering for bonding just the three 2p orbitals on each atom, because they are the only ones that are partly filled. Referring again to Figure 1-21, we can visualize end-to-end overlap (σ bonding) between the Px orbitals (one on each atom), which happen to point toward each other, and side-by- side overlap (π bonding) between the remaining p orbitals (two on each atom, and Our complete molecular orbital diagram will therefore include three sets of orbital interactions, each of which is shown separately at the left on page 13, and then the three combined at the right. The