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Resonance in chemistry6/20/2023 ![]() ![]() ![]() The true acetate ion’s nature is closer to structures A and B. The average of all the canonical structures is called the resonance hybrid, representing the averages of all the properties.įor example, the acetate ion can be depicted using three Lewis structures. For example, resonance structures 2 and 3 (shown below) explain the triple and single bond nature of CO 2, whereas structure 1 describes the double bond nature. In Benzene and aromatic systems, the delocalization of the electrons is shown as a circle-Įach canonical structure shows a localized bond explaining some of the molecules’ properties. The electron’s path indicating the partial bonds is shown in dashed lines (-) or curves. The path along which the electron movement takes place develop a partial bond character affecting the bond length, as seen in the bond lengths of Benzene and CO 2. Partial Double bond Character due to Resonance There are rules in writing resonance structures called the rules of resonance. Some structures are more favorable than others and are assigned the highest priority.Īs seen from the resonance structures, the relative positions of the atoms remain unchanged the only thing that changed was the electron position without affecting an atom’s valency.įor example, post electron delocalization, a Carbon atom cannot have more than four bonds, or Nitrogen cannot have five bonds. There are rules known as ‘Rules of resonance’ that assist in the predicting the various Lewis structures. The various structures cannot be isolated as individual molecules. The resonance structures do not represent another molecule but are hypothetical predictive structures. A double head arrow separates these canonical structures. The various Lewis structures are called the canonical/contributing/resonance structures. The probable electron movement is predicted using the curly arrows. The bond length discrepancy of both Benzene and carboxylic acid is thought to be due to electron delocalization. The electrons are assumed to be not localized between the two atoms of the bonds but are shared with more than two atoms’ nuclei.Īnd since such multiple Lewis structures were predominantly seen in molecules with multiple bonds and heteroatoms with lone pairs of electrons, it was assumed these labile electrons (pie bond and lone pairs) participate in the electron movement.įor example, in the above examples, Benzene and Carboxylic acid are represented with two Lewis structures. So, while Lewis structures depicted the concentration of two electrons between the atoms, its shortcoming in explaining the irregularities in the bond length was overcome by the proposition of electron delocalization. These observations led to the postulation that electrons in such structures aren’t localized but are delocalized. The bonds in CO 2 showed some triple bond character with the bond length of 1.15 A o intermediate to Carbon-Carbon double (bond length 1.22 A o ) and a triple bond (1.10 A o ). Similar observations were seen in double-bonded CO 2 and other molecules as well. Still, it could not explain why all the bonds behaved equally also, why Benzene showed a substitution reaction and not an addition reaction typical to the double bonds. Still, not one structure could explain all the properties.įor example, the Lewis structure of Benzene shows single and double bonds. In such instances, when a molecular structure could be represented by more than one electronic structure, each structure was found to explain most of the properties. Similarly, multiple Lewis structures can be drawn for molecules like Ozone (O 3), Acids (R-COOH), Carbon dioxide (CO 2), Carbonate ion (CO 3 2-), Nitrate ion (NO 3 -), 1,3-butadiene (CH 2=CH-CH=CH 2), etc. ![]()
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