Question

Draw Lewis structures for the following species, indicating formal charges and resonance where applicable: (a) \(\mathrm{HCO}_{2}=\) (b) \(\mathrm{HCO}_{3}^{-}\) (c) \(\mathrm{FSO}_{3}^{-}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{3}^{2-}\) (the nitrogen atoms are joined centrally with one oxygen atom on one \(\mathrm{N}\) and two on the other)

Short answer

The Lewis structures can be drawn by following these steps: counting the total number of valence electrons, arranging the atoms with the less electronegative atom usually in the center, placing a single bond between the central atom and the surrounding atoms, distributing the remaining electrons in the form of lone pairs, and creating double or triple bonds if necessary to fulfill the octet rule. Resonance structures may be present when there are several valid ways to place the remaining electrons.

Step-by-step solution

Calculate Total Valence Electrons

This is determined by adding up the number of valence electrons each atom in the molecule/ion has. Hydrogen has 1, oxygen has 6, carbon has 4, nitrogen has 5, sulfur has 6, and fluorine has 7 valence electrons. For ions, remember to add electrons if the molecule/ion carries a negative charge and subtract electrons if it carries a positive charge.

Arrange Atoms

The less electronegative atom is typically the central atom (hydrogen never goes in the center). The arrangement of atoms in \( \mathrm{N}_{2} \mathrm{O}_{3}^{2-} \) is mentioned in the question.

Place Four Two-Electron Bonds

Place a single bond (two electrons) between the central atom and surrounding atoms.

Spread Out Remaining Electrons

Starting from the outer atoms, remaining electrons are placed in the form of lone pairs. Each atom, except hydrogen, should be surrounded by 8 electrons (Octet Rule). If we've not used all the electrons after the outer atoms have obtained octets, place left electrons around the central atom.

Create Double or Triple Bonds

If any atoms don’t have a full octet, move electrons from lone pairs on the outer atoms to create double or triple bonds. You may see different resonance forms, if applicable, since those are alternate ways the same molecule could look.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom. They play a crucial role in bonding because they can be shared or transferred to achieve a stable electron configuration. To find the number of valence electrons, look at the group number in the periodic table for each element:

• Hydrogen has 1 valence electron.
• Oxygen has 6 valence electrons.
• Carbon has 4 valence electrons.
• Nitrogen has 5 valence electrons.
• Sulfur has 6 valence electrons.
• Fluorine has 7 valence electrons.

For ions, you need to adjust the total by adding or subtracting electrons based on the charge. For example, a negative charge means adding electrons, while a positive charge involves subtracting them. Understanding valence electrons is the first step in drawing Lewis structures, as they determine how atoms bond together.

Formal Charges

Formal charges help us determine the most representative Lewis structure for a molecule. It shows whether atoms within a molecule possess a hypothetical charge. To calculate formal charge:

• Count the number of valence electrons in the free atom.
• Subtract the non-bonding (lone pair) electrons on the atom in the structure.
• Subtract half the bonding (shared) electrons.

The formula looks like this: \[ \text{Formal charge} = \text{Valence electrons in free atom} - \left(\text{Non-bonding electrons} + \frac{1}{2} \times \text{Bonding electrons}\right)\]
The goal is to have formal charges close to zero for each atom in the molecule. This rule helps guide adjustments in Lewis structures, making them more stable.

Resonance Structures

Resonance structures represent different possible configurations for the same molecule, showcasing electron distribution variations. Certain molecules cannot be accurately depicted by a single Lewis structure. Instead, multiple valid structures may exist.

Resonance helps in illustrating multiple ways to arrange electrons. This is important when a molecule has atoms that can form multiple equivalent bonds. Each form does not exist individually; rather, the true structure is a hybrid of all resonance forms.

• For example, ozone \(\text{O}_3\) has two resonance structures involving alternating double and single bonds.
• In ionized structures, like nitrate \(\text{NO}_3^-\), multiple resonance forms become even more significant due to electron movement.

Recognizing resonance in molecules assists in predicting behavior and stability.

Octet Rule

The octet rule is a chemical guideline stating that atoms tend to form compounds in ways that allow each atom to have eight electrons in its valence shell, resembling the electron configuration of noble gases. This concept is central in drawing Lewis structures, as it guides how bonds are formed.

Exceptions exist, for example:

• Hydrogen is satisfied with 2 electrons.
• Elements in the third period and beyond, like sulfur and phosphorus, can have expanded octets.

Most atoms in compounds satisfy the octet rule through

• Sharing electrons in covalent bonds.
• Forming multiple bonds to achieve a stable structure.

Knowledge of the octet rule allows students to predict bonding patterns and understand molecular geometry. It’s a fundamental concept in chemistry that simplifies understanding how molecules come together.