Question

Draw the resonance structures for the formate ion, \(\mathrm{HCO}_{2}^{-},\) and find the formal charge on each atom. If an \(\mathrm{H}^{+}\) ion is attached to \(\mathrm{HCO}_{2}^{-}\) (to form formic acid), does it attach to \(\mathrm{C}\) or \(\mathrm{O} ?\)

Short answer

Draw resonance forms with double bonds on different oxygens; \( \mathrm{H}^{+} \) attaches to oxygen.

Step-by-step solution

Draw the Lewis Structure for Formate Ion

Start by drawing the Lewis structure of the formate ion, \( \mathrm{HCO}_{2}^{-} \). Place carbon (C) at the center due to its ability to form more bonds. Oxygen (O) atoms and the hydrogen (H) atom are placed around it. Carbon shares a single bond with hydrogen and is initially connected to both oxygens through single bonds. Count and adjust electrons to conform to the ion's negative charge, ensuring all atoms meet the octet rule where applicable.

Distribute Electrons and Assign Charges

The formate ion has a total of 18 valence electrons. After placing H-C-O bonds, distribute the remaining electrons among the oxygens as lone pairs to fulfill the octet rule. Assign formal charges: \( \text{Formal charge} = \text{valence electrons} - \text{non-bonding electrons} - \frac{1}{2}(\text{bonding electrons}) \). Initially, both oxygens will have non-zero formal charges.

Create First Resonance Structure

Move one lone pair from one of the oxygen atoms to form a double bond with carbon, converting that oxygen to an \( \text{O}^{0} \) (neutral charge) with no lone pairs, and the other oxygen becomes \( \text{O}^{-} \) due to having three lone pairs.

Create Second Resonance Structure

Reverse the double bond position: Move another lone pair from the singly bonded oxygen (from the other resonance structure) to form a double bond, converting this oxygen to neutral \( \text{O}^{0} \). The other oxygen with three lone pairs becomes \( \text{O}^{-} \).

Determine Attachment Site for Proton

Now consider where the \( \mathrm{H}^{+} \) would most likely attach to form formic acid. Analyze electron density regions and stability considerations. The proton attaches to the oxygen atom that carries the negative charge in the formate ion, due to higher electron density being more attractive to the proton.

Key concepts

Formate Ion

The formate ion, represented as \( \mathrm{HCO}_{2}^{-} \), is an anion that plays a crucial role in many chemical reactions, particularly as an intermediate in processes like oxidation and gas production. The negative charge on the ion signifies the presence of an extra electron, which influences how atoms within the ion are connected and how they interact with other chemical species.
Formate is the simplest carboxylate anion, derived from formic acid by deprotonation. Its resonance structures help illustrate the distribution of electrons within the molecule.
This highlights the different ways electrons can be arranged, making the ion more stable by delocalizing its charge across the molecule through resonance. Resonance doesn't mean the ion flips between different structures, rather its true form is a hybrid of all possible structures.

Formal Charge

Formal charge is a concept used to determine the distribution of electrons in a molecule, critical when drawing Lewis structures. It is calculated using the formula:
\[ \text{Formal charge} = \text{valence electrons} - \text{non-bonding electrons} - \frac{1}{2}(\text{bonding electrons}) \]
This calculation helps predict the stability and reactivity of various atoms within a molecule.

• Atoms prefer to have a formal charge as close to zero as possible, making the molecule more stable.
• In the formate ion, the formal charge assignment reveals that oxygen atoms often carry a partial negative charge, while carbon tends to have a neutral or positive charge.

Understanding formal charges aids in determining the most plausible resonance structures and the potential attachment sites for added protons or other atoms.

Lewis Structure

The Lewis structure is a simplified way to represent the valence electrons in an atom. In chemical diagrams, Lewis structures illustrate the bonding between atoms and the lone pairs of electrons that may exist in the molecule.
For the formate ion \( \mathrm{HCO}_{2}^{-} \), the central atom is carbon due to its ability to form multiple bonds, flanked by two oxygens and one hydrogen. While drawing:

• Start by connecting hydrogen directly to carbon, and both oxygens to carbon using single bonds.
• Next, distribute remaining valence electrons as lone pairs on oxygens to satisfy the octet rule.
• Consider resonance by moving lone pairs to form double bonds, observing changes in formal charges.

Lewis structures are fundamental in predicting molecular geometry, possible reactions, and understanding the concept of resonance.

Formic Acid

Formic acid, with the formula \( \mathrm{HCOOH} \), is the simplest carboxylic acid. It is formed when a hydrogen ion (\( \mathrm{H}^{+} \)) attaches itself to the negatively charged oxygen atom in the formate ion. This attachment is guided by electron density, as the proton naturally binds to the more electron-rich area to create a stable molecule.
Formic acid has multiple uses, ranging from its role in nature as a component in ant stings to its application in industrial processes. The addition of the proton to formic acid not only stabilizes the formate ion but also results in a new compound with unique chemical properties.

• Formic acid is known for its reducing properties and can partake in various chemical reactions.
• It serves as a starting material for many chemical syntheses.

Studying formic acid helps elucidate practical applications of resonance and formal charges in modifying molecular behavior.