Write resonance forms that describe the distribution of electrons in each of these molecules or ions. (a) sulfur dioxide, \(\mathrm{SO}_{2}\) (b) carbonate ion, \(\mathrm{CO}_{3}^{2-}\) (c) hydrogen carbonate ion, \(\mathrm{HCO}_{3}^{-}\) (C is bonded to an OH group and two O atoms) (d) pyridine: (e) the allyl ion:

When it comes to sulfur dioxide (SO₂) resonance, envision two distinct structures where the position of the double bond shifts between the sulfur atom and each of the oxygen atoms. In essence, resonance structures help us represent all the possible ways electrons can be arranged based on the actual distribution of electrons across a molecule.

For SO₂, think of sulfur as the central atom that likes to share electrons in different configurations. The alternating single and double bonds to the oxygen atoms illustrate a scenario where electrons are 'delocalized', or shared among more than two atoms, giving us a more accurate picture of the electronic distribution within the molecule. This delocalization contributes to the stability of the molecule and is a fundamental concept in understanding molecular structure and reactivity.

• The first structure has sulfur with a double bond to one oxygen and a single bond to another.
• The second structure is where the double bond 'switches' to the other oxygen.

The carbonate ion (CO₃^2-) presents a case where one carbon atom is surrounded by three oxygen atoms, leading to multiple resonance forms. Resonance in the carbonate ion exemplifies the concept of equivalent structures - where multiple resonance forms have the same arrangement of atoms, but the electrons are distributed differently.

In the carbonate ion's resonance structures, the double bond rotates among the three oxygen atoms, and the negative charges are also shared among the oxygen atoms. This redistribution of charge and bonding electrons demonstrates the fluid nature of electron sharing in such ions, revealing the importance of considering multiple structures for a complete understanding of the electron distribution:

• Each resonance form shows a different oxygen atom double-bonded to carbon.
• The negative charges on the oxygen atoms are delocalized over the entire anion.

Delving into the hydrogen carbonate ion (HCO₃^-) resonance reveals interesting aspects of electron distribution when an additional atom is involved. In this ion's case, a hydrogen is bonded to one of the oxygen atoms, forming a hydroxyl group.

The Lewis structures suggest that there are two primary ways to distribute the electrons in the presence of this hydroxyl group:

• One with a double bond between carbon and one oxygen, and the other oxygen bearing a negative charge.
• Another structure with roles reversed, displaying the flexibility of electron pairs to form multiple valid structures.

These resonance structures illustrate the concept of electron delocalization in polyatomic ions and how it affects chemical reactivity and bonding.

The chemical structure of pyridine, an aromatic compound, showcases a nitrogen atom integrated within a six-membered ring. The resonance structures of pyridine are unique because the nitrogen atom carries a lone pair of electrons that do not partake in resonance - it does not contribute to the delocalized pi electron system of the ring.

The key to understanding pyridine's resonance is recognizing that the aromaticity, or the continuous overlap of p-orbitals across the ring, is preserved in all resonance forms. Despite the presence of this lone pair on the nitrogen, the pi electrons can still shift around the carbon atoms of the ring. This characteristic distinguishes pyridine from other similar heterocyclic structures:

• Pyridine maintains its aromaticity despite the nitrogen's non-participating lone pair.
• The resonance forms indicate different placements of pi electrons around the ring, but not involving the nitrogen's lone pair.

Resonance in the allyl ion, a simple carbon-based ion, is an excellent demonstration of the mobility of pi electrons in unsaturated hydrocarbon chains. The allyl ion possesses a positive charge due to the absence of an electron, which is significantly involved in its resonance forms.

Resonance structures for the allyl ion show the positive charge alternating between the terminal carbon atoms. Electron delocalization is key here, exemplifying how resonance can stabilize ions by spreading out the electron deficiency (the positive charge) over multiple atoms:

• The first resonance form has a positive charge on one end of the chain.
• Then, by shifting electrons through the pi-bond, the positive charge appears at the opposite end in the second form.

This underscores the fundamental principle of resonance - providing a dynamic view on how electrons can be present in multiple configurations, thus stabilizing the overall structure.