Resonance structures are a conceptual tool used in chemistry to represent bonding when a single Lewis structure fails to accurately describe the distribution of electrons. Lewis structures depict valence electrons as localized in specific bonds or on particular atoms. However, many molecules and ions exhibit properties that contradict this localized model, suggesting that electrons are actually spread out, or delocalized, over multiple atoms. Resonance structures, also called canonical forms, are a set of two or more Lewis structures that collectively represent this electron delocalization.
The Need for Resonance Structures
The concept of resonance is required because a single Lewis structure often cannot account for the experimentally observed properties of a molecule. For instance, in the carbonate ion (\(\text{CO}_3^{2-}\)), a single Lewis structure predicts one double bond and two single bonds, implying two different bond lengths. Since double bonds are shorter and stronger than single bonds, this structure predicts two different bond lengths. Experimental measurements, however, show that all three carbon-oxygen bonds in the carbonate ion are identical in length and strength, intermediate between a typical single and double bond. This discrepancy demonstrates that the pi electrons are delocalized over all three oxygen atoms. Resonance structures provide a way to show that the actual structure is an average of all possible localized arrangements, resolving the conflict between theory and observation.
Rules for Drawing Resonance Structures
Drawing valid resonance structures follows a strict set of rules, as they are theoretical representations of the same single molecule. The most fundamental rule is that only electrons can move; the position of all atomic nuclei (the “sigma framework”) must remain unchanged. Specifically, only pi electrons and non-bonding lone pairs are permitted to be shifted. The movement of these electrons is indicated using curved arrows, starting at an electron pair and pointing to the new position. This movement must not violate the octet rule for second-row elements like carbon, nitrogen, and oxygen. Importantly, the overall net charge of the molecule or ion must be conserved across all resonance structures. A double-headed arrow (\(\leftrightarrow\)) is always used to link the individual resonance structures, signifying that they are contributors to a single, real structure, not molecules in equilibrium.
Understanding the Resonance Hybrid
The individual resonance structures are hypothetical, but the true electronic structure of the molecule is the resonance hybrid. The hybrid is a single, accurate representation that exists as a weighted average of all valid contributing structures. It is the only structure that truly exists in nature and exhibits all experimentally determined properties, such as equal bond lengths and charge distribution. This averaged nature means that the bonds involved in resonance possess fractional bond orders, which are non-integer values. For example, the carbon-oxygen bond in the carbonate ion has a bond order of 1.33, reflecting the delocalization over three positions. Furthermore, the charge is also delocalized, resulting in partial charges (\(\delta+\) or \(\delta-\)) on atoms in the hybrid structure. These partial charges represent the averaged formal charges from all the contributing structures.
How Resonance Affects Stability and Reactivity
The delocalization of electrons across multiple atoms in the resonance hybrid leads to a significant increase in molecular stability, a phenomenon known as resonance stabilization. By spreading electron density over a larger area, the system’s potential energy is lowered, and electron-electron repulsion is minimized. The difference in energy between the actual molecule and the most stable, hypothetical contributing structure is called the resonance energy, which serves as a measure of this added stability. Not all resonance structures contribute equally to the hybrid; the most stable contributors are those that follow the octet rule for all atoms and have the fewest number of formal charges. Structures that place a negative charge on the more electronegative atom are also more stable and therefore contribute more significantly. This stability profoundly impacts a molecule’s reactivity, often making it less reactive than similar molecules with localized electrons. The delocalization of charge can also direct the site of a chemical reaction by creating specific areas of higher or lower electron density.