How many electrons are delocalized in the system of ozone

Resonance hybrids necessarily contain some "abnormal" electrons. A lone pair may seem to have some bonding characteristics; instead of sticking near one atom, it visits two atoms. A bond pair may appear to move between two different pairs of atoms. These electrons step outside the boundaries that Lewis' theory has set for them, and we consider them to be delocalized.

The easiest way to spot delocalized electrons is to compare electron locations in two resonance forms. If a pair appears in one place in one form, and in a different place in another form, the pair is delocalized. You can see delocalized behavior in resonance forms I and II below.

Both O 2 and O 3 accept electrons and hydrogen ions to produce water see reduction half-reactions in Equations 1 and 2. Note: the reduction of O 2 is the basis of the electron transport chain, which is discussed in the tutorial Energy for the Body: Oxidative Phosphorylation. The more positive the standard reduction potential, the more readily the substance accepts electrons. DG is a measure of spontaneity, and negative values for DG indicate a spontaneous reaction.

Negative values of DG correspond to positive values for the reduction potential, e. Hence, we see that O 3 is a stronger oxidizing agent than O 2. Our knowledge of chemistry predicts that high concentrations of ozone might be unhealthy because it is a stronger oxidizer than oxygen and might therefore damage lung tissue. A number of tissue culture and animal studies confirm this hypothesis.

In vitro exposure to ozone causes airway epithelium damage and lipid oxidation. Human bronchial epithelial cells exposed to ozone in tissue culture release molecules that cause inflammation. We have seen them in compounds like nitrogen.

A second bond is generally made through a pi bonding interaction. Therefore, we are only going to worry about the orbitals that will form pi bonds. Populating these orbitals, and getting an exact energy, is not possible given the huge approximations we have made.

However, in focusing on the pi bonding, we see something that we can't see in Lewis terms. There really is a pi bond that stretches the entire length of the ozone molecule. This is the lowest energy combination, with a wavelength steretching over twice the length of the molecule. Because it is low in energy, the extended pi bond is pretty certain to be populated by electrons, and it will make some contribution to the structure of ozone.

How much impact it has would depend on the population of the other combinations, which we can't predict without a more careful approach. In each of the following cases, there may or may not be conjugation involving lone pairs and pi bonds.

Show why or why not, using drawings of the orbitals involved. The molecule acetamide is shown in problem MO Explain the following structural features. In acetamide, the C-N and C-O bond lengths are 1. For comparison, some typical bond lengths are C-N 1. The two C atoms, plus the O, the N and the two hydrogens on the N lie in a plane. The two hydrogens on the N are not in identical chemical environments.

All of the answers depend on an understanding of the contributions of two resonance structures to the overall picture of acetaminde, or alternatively, that actetamide forms a conjugated pi system with four electrons delocalized over the O, C and N. Contribution of the second resonance structure introduces some double bond character to the C-N bond and some single bond character to the C-O bond.