Question

Use Lewis structures and formal charges to determine the bond type (single, double, or triple) for each bond in (a) \(\mathrm{POF}_{3}\) (b) \(\mathrm{SOF}_{4}\)

Short answer

In \( \mathrm{POF}_{3} \), \( \mathrm{P} \text{-} \mathrm{O} \) is a double bond, \( \mathrm{P} \text{-} \mathrm{F} \) are single bonds. In \( \mathrm{SOF}_{4} \), all bonds are single.

Step-by-step solution

Determine the Valence Electrons

First, sum up the valence electrons for each atom in the molecule.(a) For \( \mathrm{POF}_{3} \):- \( \mathrm{P} \) (Phosphorus): 5 valence electrons- \( \mathrm{O} \) (Oxygen): 6 valence electrons- \( \mathrm{F} \) (Fluorine): 7 valence electrons each \( \times 3 = 21 \) - Total for \( \mathrm{POF}_{3} \): \( 5 + 6 + 21 = 32 \) valence electrons(b) For \( \mathrm{SOF}_{4} \):- \( \mathrm{S} \) (Sulfur): 6 valence electrons- \( \mathrm{O} \) (Oxygen): 6 valence electrons- \( \mathrm{F} \) (Fluorine): 7 valence electrons each \( \times 4 = 28 \)- Total for \( \mathrm{SOF}_{4} \): \( 6 + 6 + 28 = 40 \) valence electrons.

Draw the Skeletal Structure

Create a skeletal structure for each compound with the least electronegative atom in the center.(a) For \( \mathrm{POF}_{3} \):- Place \( \mathrm{P} \) in the center, connected to \( \mathrm{O} \) and three \( \mathrm{F} \) atoms.(b) For \( \mathrm{SOF}_{4} \):- Place \( \mathrm{S} \) in the center, connected to \( \mathrm{O} \) and four \( \mathrm{F} \) atoms.

Distribute the Electrons

Distribute the total valence electrons around the structure to fulfill the octet rule as much as possible, starting with the outer atoms.(a) \( \mathrm{POF}_{3} \):- First, assign electrons to fluorine atoms, as they prefer to have a full octet.- Remaining electrons are placed as lone pairs on the central phosphorus atom. Balance the distribution to satisfy the octet, creating multiple bonds where necessary.(b) \( \mathrm{SOF}_{4} \):- Start with fluorine and oxygen by direct single bonds with \( \mathrm{S} \), fulfilling their octets. - Extra electrons are adjusted around the central \( \mathrm{S} \), potentially placing lone pairs where applicable.

Calculate Formal Charges

Calculate the formal charge for each atom to ensure the most stable Lewis structure (ideally close to zero).(a) \( \mathrm{POF}_{3} \):- \( \mathrm{P} = 5 - (N_{bonding}/2 + N_{lone}) = 5 - (8/2 + 0) = 0 \)- \( \mathrm{O} = 6 - (N_{bonding}/2 + N_{lone}) = 6 - (4/2 + 4) = 0 \)- \( \mathrm{F} = 7 - (N_{bonding}/2 + N_{lone}) = 7 - (2/2 + 6) = 0 \)(b) \( \mathrm{SOF}_{4} \):- \( \mathrm{S} = 6 - (N_{bonding}/2 + N_{lone}) = 6 - (10/2 + 0) = +1 \)- \( \mathrm{O} = 6 - (N_{bonding}/2 + N_{lone}) = 6 - (2/2 + 6) = 0 \)- \( \mathrm{F} = 7 - (N_{bonding}/2 + N_{lone}) = 7 - (2/2 + 6) = 0 \) - The structure with the lowest sum of formal charges is usually the most stable.

Identify the Bond Types

Determine the bond types by examining the skeletal structure and electron distribution.(a) \( \mathrm{POF}_{3} \):- \( \mathrm{P} \text{-} \mathrm{O} \) forms a double bond (due to oxygen's double bond tendency and formal charge adjustment).- Each \( \mathrm{P} \text{-} \mathrm{F} \) bond is a single bond.(b) \( \mathrm{SOF}_{4} \):- \( \mathrm{S} \text{-} \mathrm{O} \) is a single bond.- \( \mathrm{S} \text{-} \mathrm{F} \) bonds are all single bonds.

Key concepts

Understanding Formal Charges

In the world of chemistry, understanding formal charges is critical for drawing accurate Lewis structures. The formal charge of an atom is a tool that helps us keep track of electrons in molecules. It is calculated using the formula: \[\text{Formal Charge} = \text{Valence Electrons} - \left(\frac{\text{Bonding Electrons}}{2} + \text{Non-bonding Electrons}\right)\]This formula essentially balances the number of valence electrons an atom should have with the electrons it actually "owns" in a structure. The goal is to arrive at a structure where formal charges are as close to zero as possible, indicating stability. Do note that exceptions apply, especially in larger and more complex molecules. Always remember:- **Atoms prefer to have formal charges close to zero.**- **Negative charges are more stable on more electronegative elements.**- **Don't dismiss structures with small formal charges; they may be significant for understanding reactivity.**Balancing formal charges helps us predict the most stable configuration of a molecule, which aligns neatly with how molecules would naturally stabilize.

Types of Chemical Bonds

Chemical bonds are the glue that hold atoms together. In the context of Lewis structures, bonds are represented by lines connecting atoms. Understanding bond types is pivotal in explaining how atoms are held together. There are three primary types of bonds:

• **Single Bond:** This involves one pair of shared electrons (2 electrons). Represented by a single line, such as the bonds in most of the fluorine atoms in both \( \mathrm{POF}_3 \) and \( \mathrm{SOF}_4 \).
• **Double Bond:** This involves two pairs of shared electrons (4 electrons). Double bonds often occur with oxygen as in the \( \mathrm{P} - \mathrm{O} \) bond in \( \mathrm{POF}_3 \).
• **Triple Bond:** This involves three pairs of shared electrons (6 electrons). Triple bonds are less common in the molecules discussed, but a classic example is the nitrogen-nitrogen bond in nitrogen gas \( \mathrm{N}_2 \).

Each bond type, signified by the number of electron pairs shared, affects the molecule’s characteristics, including its length, strength, and reactivity. Double and triple bonds are generally stronger and shorter than single bonds, and knowing these differences aids in predicting molecular behavior.

The Role of Valence Electrons

Valence electrons play a crucial role in determining the bonding behavior of atoms. These are the electrons found in the outermost shell of an atom and are the ones involved in forming chemical bonds. For drawing Lewis structures, beginning with the number of valence electrons in each atom is vital.Here’s how you can determine the number of valence electrons:- **Phosphorus (\( \mathrm{P} \)):** 5 valence electrons.- **Oxygen (\( \mathrm{O} \)):** 6 valence electrons.- **Sulfur (\( \mathrm{S} \)):** 6 valence electrons.- **Fluorine (\( \mathrm{F} \)):** 7 valence electrons.When you add the valence electrons for each atom in a molecule, you can distribute them to satisfy the octet rule, where each atom (except hydrogen) seeks to have eight electrons completing its valence shell. Remember the importance of the octet rule:

• Atoms "prefer" configurations where the valence shell resembles that of a noble gas.
• Exceptions include molecules with odd numbers of valence electrons, central atoms that can hold more than eight electrons, and molecules with too few valence electrons to complete an octet.

In constructing Lewis structures like those for \( \mathrm{POF}_3 \) and \( \mathrm{SOF}_4 \), accurately accounting for these electrons is key to achieving stable and realistic representations of molecular structure.