A Brief Note on Resonance

Resonance is a mental exercise under the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. It includes generating numerous Lewis structures that, when assembled, represent the complete electrical structure of the molecule. Resonance structures are employed when a single Lewis structure cannot properly characterize the bonding; the combination of possible resonance structures is defined as a resonance hybrid, which depicts the total delocalization of electrons inside the molecule. In general, molecules with several resonance structures will be more stable than one with fewer and some resonance structures contribute more to the stability of the molecule than others – formal charges aid in establishing this.

Resonance is a term used to describe delocalized electrons within specific compounds or polyatomic ions whose bonding cannot be represented using a single Lewis formula. Numerous resonance structures are used to depict a molecule or ion with such delocalized electrons. The Lewis nuclear skeleton.The structure of these resonance structures is unchanged; only the electron positions are altered. This is the case for ozone (O3), an oxygen allotrope with a V-shaped structure and a 117.5° O–O–O angle.

DIAGRAMMATIC REPRESENTATION

Contributing structures are often denoted in diagrams by double-headed arrows (↔). The arrow should not be confused with the equilibrium arrows pointing right and left (⇌). All structures may be surrounded by huge square brackets to indicate that they represent a single molecule or ion, rather than many species in chemical equilibrium.

In addition to contributing structures, a hybrid structure can be employed in diagrams. Pi bonds involved in resonance are typically depicted as curves or dotted lines in a hybrid structure, indicating that they are partial rather than complete pi bonds. The delocalized pi-electrons of benzene and other aromatic rings are sometimes shown as a solid circle.

BOTH MAJOR AND MINOR CONTRIBUTORS 

One of the contributing structures may bear a greater resemblance to the actual molecule than another (in the sense of energy and stability). Potential energy structures with a low value are more stable than those with a high value and more closely match the actual structure. Major contributors are the most stable contributing structures. Energetically unfavorable and thus less advantageous structures play a limited role. With rules listed roughly in decreasing order of relevance, important contributors are often structures. 

• adhere to the octet rule to the greatest extent possible (8 valence electrons around each atom rather than shortfalls or surpluses, or 2 electrons for Period 1 elements)

• have the greatest possible amount of covalent bonds

• carry the fewest number of formally charged atoms possible, with the spacing of opposite and like charges minimized and maximized, respectively

• Negative charge, if any, should be applied to the most electronegative atoms and positive charge, if any, should be applied to the most electropositive atoms

• do not significantly differ from idealized bond lengths and angles (for example, the relative insignificance of Dewar-type resonance contributors to benzene)

• Locally, preserve aromatic substructures while avoiding anti-aromatic ones

SETS OF LEWIS STRUCTURE

A Lewis Structure is a highly simplistic representation of the valence shell electrons of a molecule. It is used to depict how the electrons are organized around specific atoms in a molecule. Electrons are portrayed as “dots” or for bonding electrons as a line connecting the two atoms. The purpose is to obtain the “best” electron configuration, i.e. the octet rule and formal charges need to be satisfied.

Additionally, we use Lewis symbols to denote the formation of covalent bonds, which are depicted in Lewis structures, which are diagrams illustrating the bonding of molecules and polyatomic ions. For instance, two chlorine atoms share one pair of electrons when they create a chlorine molecule:

According to the Lewis structure, each Cl atom contains three pairs of electrons that are not used in bonding (referred to as lone pairs) and one pair of electrons that is shared (written between the atoms). A dash (or line) is occasionally used to denote a pair of electrons that are shared:

A single bond is a shared pair of electrons. Each Cl atom has eight valence electrons: six in lone pairs and two in single bonds.

CONCLUSION- 

Knowing an atom’s Lewis structure enables you to predict how and how many bonds it will form. Eventually, this understanding will enable us to comprehend the shapes of molecules and their chemical properties. A resonance form is another technique to represent the Lewis dot structure of a molecule. Lewis structures that are equivalent are referred to as resonance forms. They are utilized when there are multiple possible configurations for double bonds and lone pairs on atoms.