Ozone formal charge

Seif-Test 2.1 IB Calculate the formal charges for the three oxygen atoms in one of the Lewis structures of the ozone resonance structure (Example 2.5). 

Problem 2.24 Write contributing structures, showing formal charges when necessary, for (a) ozone, O, (b) CO (c) hydrazoic acid, HN, (d) isocyanic acid, HNCO. Indicate the most and least stable structures and give reasons for your choices. Give the structure of the hybrid. ... 

Give the formal charge for each atom in the nitrate ion, N03, and ozone, O3 (yes, even neutral molecules can have elements with formal charges). 

Let us illustrate the concept of formal charge using the ozone molecule (O3). Proceeding by steps, as we did in Examples 9.3 and 9.4, we draw the skeletal structure of O3 and then add bonds and electrons to satisfy the octet rule for the two end atoms ... 

We can now write the Lewis structure for ozone and show the formal charges ... 

Consider nitrogen dioxide (NO2) as an example. A major contributor to urban smog is formed when the NO in auto exhaust is oxidized. NO2 has several resonance forms. Two involve the O atom that is doubly bonded, as in the case of ozone. Several others involve the location of the lone electron. Two of these resonance forms are shown below. The form with the lone electron on the singly bonded O has zero formal charges (right) ... 

A What is the formal charge on the indicated atom in each of the following species (a) sulfur in SO2 (b) nitrogen in N2H4 (c) each oxygen atom in ozone, O3. 

A molecule of ozone (O3) consists of three oxygen atoms covalently bonded together in a V-shaped molecule. As such it provides an appropriate example for discussing different aspects of the various approaches to covalent bonding resonance forms and bond order, formal charge and bond polarization, and the hybridization of molecular orbitals. Ozone is also of note in terms of its environmental significance both in the lower levels of the stratosphere and at ground level . 

The bonds in ozone are between oxygen atoms and so are in themselves non-polar. However, the molecule as a whole has a net dipole moment which can be explained by checking the formal charges on each atom (Figure 14.60). This shows the uneven distribution of electrons through the structure. 

Figure 14.60 The formal charges on the atoms of the ozone molecule, showing the net dipole of the structure... 

An atom s formal charge is calculated as follows We can illustrate the concept of formal charge using the ozone molecule (O3). Use the step-by-step method for drawing Lewis structures to draw the Lewis structure for ozone, and then determine the formal charge on each O atom by subtracting the number of associated electrons from the number of valence electrons. 

Confirm that there is no legitimate Lewis structure of ozone with all zero formal charges. 

Little is known of the actual mechanism. A mode of reaction is possible, in which the oxygen atom at the top of the ozone molecule with a formal positive charge (p. 230) reacts with an electron pair, not localized in a bond but on one carbon atom, and in which the ozone therefore reacts by an electrophilic mechanism (Wibaut, Sixma and Kampschmidt). However, in order to explain the differences between the reaction course for ozonization and for other electrophilic reactions, e.g., bromination and nitration with pyrene, these authors assume also an interaction of one of the other oxygen atoms with the adjacent carbon atom. The net result is, however, about the same as that predicted by the bond localization hypothesis. 

The atoms need not be carbon it is the orbitals that matter. Charged species can be involved, for example the allyl cation CH2=CH-CH2+ has two n electrons (formally two from the double bond and none from the carbocation centre), the allyl anion CH2=CH-CH2 has four n electrons (two from the double bond and two from the lone pair on the negatively charged carbon). Ozone, 0=0+-0, is isoelectronic with the allyl anion, and will also bring four n electrons to the reaction (two from the double bond and two from the lone pair on the negatively charged oxygen). 

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