Resonance Structure of O3

Resonance structures are the Lewis structure sets that characterize the electron’s delocalization in a molecule or a polyatomic ion. Resonance structures are a more accurate representation of a Lewis dot structure than Lewis dot structures because they clearly illustrate the bonding between molecules. Not all resonance structures are created equal; some are superior to others. The better ones have the fewest formal charges, the most electronegative atoms have the most formal charges, and the structure maximizes bonding. The more resonance forms a molecule has, the more stable the molecule is. They are connected by a double-headed arrow, indicating that the true structure is between the resonance structures. Curved arrow notation was employed to depict the flow of electrons from one resonance type to the next.

Resonating structures of Ozone

Ozone (O3) is an oxygen allotrope composed of three oxygen atoms. Ozone has one double bond and one single bond in its Lewis structure. Additionally, two oxygen atoms in the O3 Lewis structure have charges. The Lewis structure of O3 can be deduced in various phases starting with the valence electrons of oxygen atoms. This lesson walks you through each step of drawing the Lewis structure of O3. After drawing the Lewis structure of NH3, the shape of the O3 molecule can be determined.

The ozone (O3) molecule is composed of a core oxygen atom that is linked singly to one oxygen atom and doubly to another. Although this molecule has no net charge, the Lewis structures reveal a +1 charge on the central oxygen and a -1 charge on the singly bonded oxygen. The ozone molecule’s two resonance configurations are seen below.

Ozone has an angular structure with two oxygen-oxygen bonds measuring approximately 1.278 Angstroms in length. The molecule’s 1/2 bond order, which indicates the number of bonds between two atoms, also supports the notion that the hybrid structure incorporates the ozone molecule’s two principal resonance configurations. As a result, the triple bond formed by the three oxygen atoms will exhibit strong single bond properties.

These structures will contribute relatively little because, among other things, both lack a complete octet of oxygen and have fewer covalent bonds than the other two structures, another characteristic that severely reduces structure stability.

Contributors, both major and minor

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 outlined in rough order of decreasing importance, substantial contributors are often structures that adhere to 

• the octet rule to the greatest extent feasible (8 valence electrons around each atom rather than deficits or surplus, 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.

Concluson

When various approaches of constructing a Lewis dot diagram that satisfies the octet rule exist, resonance structures arise. Bear in mind that the octet rule describes the process by which an atom acquires, loses, or shares electrons in order to have an outer electron shell with eight electrons. When a single molecule is insufficient to completely represent the bonding between surrounding atoms in comparison to empirical data on the actual bond lengths between those atoms, resonance structures are used.