In terms of the bonds present, explain why acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\), contains two distinct types of carbon-oxygen bonds, whereas the acetate ion, formed by loss of a hydrogen ion from acetic acid, only contains one type of carbonoxygen bond. The skeleton structures of these species are shown:

Understanding the chemical bonds present in acetic acid and acetate ion is crucial to explaining their distinct structural features. In acetic acid, \(\mathrm{CH}_{3}\mathrm{CO}_{2}\mathrm{H}\), there are two types of carbon-oxygen bonds, a single bond and a double bond. The double bond, represented as C=O, involves the sharing of four electrons between the carbon and oxygen atoms, leading to a shorter and notably stronger connection. On the other hand, the single bond, denoted as C-O, involves only two electrons being shared, resulting in a bond that is comparatively longer and weaker.

The transition from acetic acid to its conjugate base, the acetate ion, is accompanied by the loss of a hydrogen ion \(\mathrm{H^{+}}\), which remarkably changes the bonding scenario. In acetate ion \(\mathrm{CH}_{3}\mathrm{CO}_{2}^{-}\), the negative charge is distributed between the two oxygen atoms through a process known as resonance, which effectively equalizes the bond strength, making them appear the same. This uniformity replaces the distinct single and double bonds seen in the neutral acetic acid molecule.

The electron distribution within a molecule dictates the nature and strength of its chemical bonds. In acetic acid, electrons are distributed such that the carbonyl group (containing the C=O bond) has a higher electron density compared to the C-O bond. This uneven distribution of electron density accounts for the different bond strengths and lengths in the molecule.

When acetic acid becomes an acetate ion, a fascinating shift in electron distribution occurs. Losing a proton (\(\mathrm{H^{+}}\)) results in a negative charge that must be borne by the remaining atoms. This charge is not localized on a single oxygen atom; instead, it is delocalized across the two oxygens, creating a conjugated system. The electrons in the acetate ion are thus evenly spread over the molecule, giving it the characteristic symmetry and equivalent Carbon-Oxygen bonds.

Electron delocalization plays a crucial role in the stability of ions and molecules. In the case of the acetate ion, it is this delocalization that allows the molecule to distribute its negative charge more efficiently, lowering the overall energy and giving rise to equivalent bonding characteristics throughout.

Resonance structures are a means to portray molecules with delocalized electrons that cannot be represented by a single Lewis structure. The acetate ion exemplifies this by having non-localizable electrons that are spread over the two oxygen atoms.

These resonance structures don't represent different individual entities but rather, they show the possible distribution patterns of electrons within the entity. For the acetate ion, it translates to two resonance forms in which the negative charge is alternately associated with each oxygen. Hence, each individual Carbon-Oxygen bond is essentially a hybrid of both single and double bond character. This superposition of structures leads to the bond equivalence observed in the acetate ion.

Understanding resonance is essential as it affects molecular properties such as stability, reactivity, and bond lengths. A molecule like the acetate ion, which exhibits resonance, will typically have a lower energy state and greater stability compared to a similar structure without resonance. This is because the energy of the electrons is more spread out, rather than being concentrated in one area.