Ozone delocalized bonding

We begin our exploration of delocalized bonds with ozone, O3. As described in Chapter 7, ozone in the upper stratosphere protects plants and animals from hazardous ultraviolet radiation. Ozone has 18 valence electrons and a Lewis stmcture that appears in Figure 10-36a. Experimental measurements show that ozone is a bent molecule with a bond angle of 118°. 

Attempting to write electron-dot formulas leads you to recognize that delocalized bonding exists in many molecules. Whenever you can write several plausible electron-dot formulas—which often differ merely in the allocation of single and double bonds to the same kinds of atoms (as in ozone)—you can expect delocalized bonding... 

By using delocalized bonding schemes, we can avoid writing two or more contributing structures to a resonance hybrid, as is so often required in the Lewis theory. Consider the ozone molecule, O3, that we used to introduce the concept of resonance in Section 10-5. In place of the resonance hybrid based on these contributing structures. 

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. 

In ozone, each of the three O atoms is bound to its neighbor through a a-bond (1.278 A), the angle being 116.5 to 7° the two terminal atoms are 2.18 A apart, a distance shorter than the van der Waals distance (2.8 A). Owing to the delocalization of the It electrons, the 0—0 bond in ozone has a multiple bond character, which is reflected in the bond distance. 

Resonance attempts to correct a fundamental defect in Lewis formulas. Lewis formulas show electrons as being localized they either are shared between two atoms in a covalent bond or are unshared electrons belonging to a single atom. In reality, electrons distribute themselves in the way that leads to their most stable arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei. What we try to show by the resonance description of ozone is the delocalization of the lone-pair electrons of one oxygen and the electrons in the double bond over the three atoms of the molecule. Organic chemists often use curved arrows to show this electron... 

The total bond energy of a substance for which resonance structures are written is greater than would be expected if there were only one formal Lewis structure. This additional stabilization is called resonance energy. It arises from the same principle that is responsible for covalent bond energy, the delocalization of electrons about the atoms forming the bond. As a result of resonance in ozone, for example, the electrons constituting the second pair of the double bond are delocalized around the 3... 

Ozone is a pale blue gas with a boiling point of -111 °C and a melting point of -193 °C. It consists of three oxygen atoms which are aligned in a binding angle of 117°. The distances between the inner and the outer atoms is 128 pm, which suggests a delocalized 7r-bond [331]. 

Use Lewis formulas to depict the resonance structures of the following species from the valence bond point of view, and then sketch MOs for the delocalized tt systems (a) SO2, sulfur dioxide (b) O3 ozone (c) HC02, formate ion (H is bonded to C). 

Our need for more than one Lewis structure to depict the ozone molecule is the result of electron-pair delocalization. In a single, double, or triple bond, each electron pair is attracted by the nuclei of the two bonded atoms, and the electron density is greatest in the region between the nuclei each electron pair is localized. In the resonance hybrid for O3, however, two of the electron pairs (one bonding and one lone pair) are delocalized their density is spread over the entire molecule. In O3, this results in two identical bonds, each consisting of a single bond (the localized electron pair) and a partial bond (the contribution from one of the delocalized electron pairs). We draw the resonance hybrid with a curved dashed line to show the delocalized pairs ... 

What is delocalized tt bonding and what does it explain Explain the delocalized TT bonding system in CgHg (benzene) and O3 (ozone). 

When electrons become delocalized, it s important to remember that all the resonant Lewis structures are just partial pictures of how the molecule may appear. Just as the ozone molecule can be drawn with the double bond left or right, it can more correctly be drawn with a partial double bond to both oxygen atoms as shown on the right. 

In ozone, O3, the two oxygen atoms on the ends of the molecule are equh ent to one another, (a) What is the best choice of hybridization scheme for the atoms of ozone (b) For one of the resonance forms of ozone, which of the orbitals are used to make bonds and which are used to hold nonbonding pairs of electrons (c) Which of the orbitals can be used to delocalize the TT electrons (d) How many electrons are delocalized in the TT system of ozone ... 

Bonding in many polyatomic molecules, such as ozone or benzene, can often be best described in toms of delocalized molecular orbitals. 

Consider the free radical nitrogen dioxide, NO2, a major contributor to urban smog that is formed when the NO in auto exhaust is oxidized. NO2 has several resonance forms. Two differ in terms of which O atom is doubly bonded, as in the case of ozone. Two others have the lone electron residing on the N or on an O, so the resonance hybrid has the lone electron delocalized over these two atoms ... 

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