Problem 56 - Chemistry The Molecular Science

Question

Draw resonance structures for each of these ions: \(\mathrm{NSO}^{-}\) and \(\mathrm{SNO}^{-}\). (The atoms are bonded in the order given in each case, that is, \(\mathrm{S}\) is the central atom in \(\mathrm{NSO}^{-} .\).) (a) Use formal charges to determine which ion is likely to be more stable. (b) Explain why the two ions cannot be considered resonance structures of each other.

Step-by-Step Solution

Answer

NSO$^{-}$ is more stable with a central sulfur based on formal charges. NSO$^{-}$ and SNO$^{-}$ cannot be resonance structures of each other because they have different atom connectivity.

Step 1: Understand Resonance Structures

Resonance structures are different ways of arranging electrons in a molecule that can show the same overall connectivity of atoms. They distribute electrons differently around a central skeleton of atoms, helping to explain electron behavior more accurately than any single depiction.

Step 2: Draw Main Structure for NSO$^{-}$

For \(\mathrm{NSO}^{-}\), draw the central sulfur atom bonded to nitrogen and oxygen. Start with a single bond between each atom, and keep track of valence electrons. Sulfur (S) can hold more than 8 valence electrons due to its position in the periodic table.

Step 3: Create Resonance Structure for NSO$^{-}$

Adjust the positions of double bonds and lone pairs around the structure to form a plausible resonance structure. Ensure the structure is within octet (or expanded octet) rules: a double bond can form between sulfur and nitrogen, and nitrogen and oxygen can form a triple bond to create one resonance structure.

Step 4: Draw Main Structure for SNO$^{-}$

For \(\mathrm{SNO}^{-}\), set sulfur as the central atom between nitrogen to the left and oxygen to the right. Begin with single bonds, similar to the NSO arrangement.

Step 5: Create Resonance Structure for SNO$^{-}$

Position electrons differently by moving double or triple bonds. Here, you might form a triple bond between the sulfur and nitrogen, while oxygen retains a single/double bond with sulfur to ensure charge distribution and atom completion of valence shells.

Step 6: Determine Formal Charges

Calculate the formal charge for each atom in both ions by using the formula \( \text{Formal charge} = \text{Valence electrons} - \text{Non-bonding electrons} - \text{Bonding electrons}/2 \). Do this for every resonance structure to determine which is the more stable form of each ion.

Step 7: Compare Stability via Formal Charges

Compare the resonance structures for both \(NSO^{-}\) and \(SNO^{-}\). The structure where atoms possess formal charges closer to zero (or the distribution is more even) represents a more stable ion.

Step 8: Explain Structural Differences

Even though the atoms involved are the same, their arrangements (NSO vs. SNO) create different structures with differing bond orders and electron distributions. As such, these ions cannot be considered resonance structures of each other because their atomic connectivity isn't the same.

Key Concepts

Formal Charges

Formal charges are a vital concept in understanding the stability of different resonance structures. They are calculated to give us an idea of the electron "ownership" by each atom in a molecule. To find the formal charge, use the formula \[\text{Formal charge} = \text{Valence electrons} - \text{Non-bonding electrons} - \frac{\text{Bonding electrons}}{2}\]This formula helps in balancing the charge across the molecule. When drawing resonance structures for ions like \( \text{NSO}^{-} \) and \( \text{SNO}^{-} \), each atom's formal charge tells us how close the structure is to being energetically favorable. Ideally, the atoms in the molecule should have a formal charge of zero, or as close to zero as possible, to reflect stability.

• Calculate formal charges carefully for all resonance structures.
• Smaller formal charges often indicate a more stable structure.
• Formal charges help in comparing different resonance structures for stability.

Electron Distribution

In molecules and ions, electron distribution illustrates how electrons are spread among the atoms within a compound. For ions like \( \text{NSO}^{-} \) and \( \text{SNO}^{-} \), electron distribution plays a key role in predicting how these compounds will behave.
These ions have unique electron arrangements because their central atom (sulfur) can expand its octet, accommodating more than eight electrons. This flexibility means we can rearrange electrons to minimize formal charges and assess possible resonance structures.

• Electron distribution can vary with different structures, affecting stability and reactivity.
• Ions typically try to distribute electrons to complete an octet around each atom.
• Expanded octets, as seen with sulfur, help explain unconventional compounds.

Understanding how electrons are distributed helps in predicting the bonding and stability scenarios within the compound.

Stability Comparison

Comparing the stability of two ions involves looking at their resonance structures and their formal charges. When examining \( \text{NSO}^{-} \) and \( \text{SNO}^{-} \), it is essential to note that these are not resonance structures of each other, despite having similar atoms.
Each ion has a different arrangement and therefore different stabilities. The resonance structure where the atoms have formal charges closest to zero is often the most stable. This is because it represents the least energetic and most balanced state of the molecule.

• Stable ions have resonance structures with minimal formal charges.
• Resonance structures cannot be the same if the connectivity of atoms is different.
• Structures with better electron distribution usually depict more stable molecules.

By examining formal charges and electron distribution, we can identify the more stable ion and understand its chemical behavior better.