Question

Draw four reasonable resonance structures for the \(\mathrm{PO}_{3} \mathrm{~F}^{2-}\) ion. The central \(\mathrm{P}\) atom is bonded to the three \(\mathrm{O}\) atoms and to the \(\mathrm{F}\) atom. Show formal charges.

Short answer

Draw four resonance forms by rearranging electrons, ensuring they all satisfy octet and formal charge rules.

Step-by-step solution

Understand the Resonance Concept

Resonance structures are different Lewis structures for the same molecule that only differ in the position of electrons, not atoms. Each structure must follow the rules of valence and the octet.

Determine the Valence Electrons

Count the total number of valence electrons in \( \mathrm{PO}_{3}\mathrm{~F}^{2-} \). Phosphorus (P) has 5, each oxygen (O) has 6, fluorine (F) has 7, and there are 2 additional electrons from the \( 2- \) charge. Total: \(5+3(6)+7+2=32\) electrons.

Draw the Basic Structure

Place P in the center with single bonds to each O and F atom, distributing single electrons as needed. Assume equal distribution of electrons initially, respecting the octet rule.

Create Resonance Structures by Redistributing Electrons

Adjust bonds to create different resonance structures. Move pairs of electrons to form double bonds between P and O atoms, ensuring that each structure also maintains 32 electrons. Keep rearranging electrons to present each structure while ensuring the octet rule is satisfied.

Assign Formal Charges

For each structure, calculate the formal charge for each atom using the formula: \( \text{Formal charge} = \text{Valence electrons} - \text{Non-bonding electrons} - \frac{1}{2}\text{Bonding electrons} \). Ensure that the sum of formal charges matches the overall ion charge (\(2-\)).

Verify and Adjust

Check that all resonance structures are valid by ensuring they adhere to formal charge rules and have an accurate total charge. Adjust structures if necessary for clarity and accuracy.

Key concepts

Valence Electrons

Valence electrons are the key players in understanding chemical bonding and structure. These are the electrons found in the outermost shell of an atom and are crucial because they participate in bonding. To determine how a molecule like \( \mathrm{PO}_3\mathrm{~F}^{2-} \) will interact, you start by counting all the valence electrons. Each element contributes its own number of valence electrons based on its position in the periodic table.
For example, phosphorus (P) contributes 5 electrons, each oxygen (O) contributes 6, and fluorine (F) contributes 7. Since \( \mathrm{PO}_3\mathrm{~F}^{2-} \) carries a \( 2- \) charge, you add 2 additional electrons to your total. This gives us \( 32 \) valence electrons to work with.
Recognizing valence electrons informs how atoms will connect and what types of bonds, like single or double, might form in different structures.

Formal Charges

Formal charges are useful for understanding the stability and reactivity of molecules. They help identify how electrons are distributed in a molecule. Imagine formal charges as a bookkeeping system for electrons, ensuring each atom follows the typical guidelines.
The formula to calculate formal charge is:

• \( \text{Formal charge} = \text{Valence electrons} - \text{Non-bonding electrons} - \frac{1}{2}\text{Bonding electrons} \)

By assessing the formal charges in \( \mathrm{PO}_3\mathrm{~F}^{2-} \), you verify that the total charge of the molecule aligns with its overall charge, which is \( 2- \). Each resonance structure will display a different distribution of electrons, and thus, yield different formal charges. However, the sum of these charges must remain consistent across all resonance structures, ensuring accurate portrayal of the ion's actual charge.

Octet Rule

The octet rule is a guiding principle in chemistry, particularly in drawing Lewis structures. It states that atoms tend to form bonds until they are surrounded by eight electrons, forming a stable electron configuration. This rule is significant when considering resonance structures.
For the \( \mathrm{PO}_3\mathrm{~F}^{2-} \) ion, each oxygen and the fluorine will aim to have an octet. Phosphorus, although typically surrounded by more due to its ability to expand its octet, generally will follow this rule as well.
When drawing the resonance structures, the redistribution of electrons must consider fulfilling the octet requirement. Generally, moving pairs of electrons to form double bonds between phosphorus and oxygen atoms can help maintain the octet across different resonance forms.

Lewis Structures

Lewis structures are diagrams that represent molecules' electronic structure. They show how atoms are bonded and where the lone pairs of electrons are positioned. Creating a Lewis structure involves connecting atoms with lines for bonds, filling valence shells to satisfy the octet rule, and minimizing formal charges.
For \( \mathrm{PO}_3\mathrm{~F}^{2-} \), you start by placing phosphorus in the center, given its lower electronegativity, and link it with oxygen and fluorine using initial single bonds. The total count of electrons used should sum to \( 32 \), as calculated from the valence electrons.
Resonance structures are variants of Lewis structures which showcase possibilities of electron placement. They provide a more comprehensive understanding of electron distribution, which cannot always be captured by a single static diagram. Through resonance and Lewis structures, we can appreciate both the molecule's stable state and the dynamic nature of electron positions.