Some resonance structures are more favorable than others. A molecule that has several resonance structures is more stable than one with fewer. The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule. Resonance structures are used when one Lewis structure for a single molecule cannot fully describe the bonding that takes place between neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. Negative charge is stabilized as we go from sp 3 to sp 2 to sp hybridization (alkyl to alkene to alkyne), since the negative charge is held in orbitals with increasing s-character closer to the positively charged nucleus. These structures are written with a double-headed arrow between them, indicating that none of the Lewis structures accurately describes the bonding but that the actual structure is an average of the individual resonance structures. Factor 6: Electronegativity Of An Atom Increases As The s-Character Of Its Hybridization Increases This Destabilizes Positive Charge. Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. Some molecules have two or more chemically equivalent Lewis electron structures, called resonance structures. Single bonds, double bonds, triple bonds, +1 charges, -1 charges, these are our limitations in explaining the structures, and the true forms can be in between - a carbon-carbon bond could be mostly single bond with a little bit of double bond character and a partial negative charge, for example. Resonance structures is a mechanism that allows us to use all of the possible resonance structures to try to predict what the actual form of the molecule would be. The best measurements that we can make of benzene do not show two bond lengths - instead, they show that the bond length is intermediate between the two resonance structures. One would expect the double bonds to be shorter than the single bonds, but if one overlays the two structures, you see that one structure has a single bond where the other structure has a double bond.
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