Question

Draw three resonance structures for the molecule \(\mathrm{N}_{2} \mathrm{O}\) in which the atoms are arranged in the order NNO. Indicate formal charges.

Short answer

The three resonance structures for N2O (NNO arrangement) show different placements of electron pairs, leading to different distributions of formal charge in each structure. The first structure has formal charges of +1 on N (leftmost), 0 on N (middle) and -1 on O. The second structure presents formal charges of +1 on N (leftmost), 0 on both N (middle) and O. The third structure has formal charges of +1 on both N atoms and 0 on O.

Step-by-step solution

- Determine the central atom and draw the structure

The first atom in the given order, Nitrogen (N), is the leftmost atom, then another Nitrogen (N) in the middle, and Oxygen (O) is the rightmost. Connect all three atoms with single bonds. Subtract the total valence electrons used in bonding (6 electrons) from the total valence electrons available (16 electrons: 5 from each Nitrogen and 6 from Oxygen). There are 10 electrons left, which are placed as lone pairs on the atoms, starting from the outer atoms.

- Draw the first resonance structure

The first resonance structure places all the remaining 10 electrons on the Oxygen atom. Oxygen now has a complete octet (4 pairs of electrons; 2 pairs from bonds and 2 lone pairs). However, the left Nitrogen only has 3 electron pairs (1 from the bond and 1 lone pair), leading to a positive formal charge. The second Nitrogen has 4 electron pairs (1 from two bonds and 1 lone pair), and so does not carry any charge. Oxygen negatively charged due to having an extra pair of electrons.

- Draw the second resonance structure

The second resonance structure has one of the lone pairs on Oxygen forming a double bond with the second Nitrogen. This leaves Oxygen with 3 electron pairs (2 pairs from bonds and 1 lone pair). The second Nitrogen has a complete octet (4 electron pairs; 3 from bonds and 1 lone pair), while the first Nitrogen remains positively charged. Oxygen does not have any formal charge in this structure.

- Draw the third resonance structure

The third resonance structure has both lone pairs on Oxygen forming double bonds with the Nitrogen atoms. Oxygen now has a complete octet (4 electron pairs in bonds). The Nitrogen atoms have 3 electron pairs each (2 pairs from bonds and 1 lone pair), giving them both a positive formal charge.

Key concepts

Formal Charge

In chemistry, understanding formal charge is critical for figuring out the most plausible resonance structures of a molecule. Simply put, a formal charge is an accounting system for electrons. It helps chemists determine the distribution of electrons in a molecule by assigning a hypothetical charge to atoms. This charge is calculated based on the electrons an atom owns in the molecule as compared to the electrons it would have on its own in a free state.
To calculate the formal charge, you can use this simple formula:

• Formal Charge = (Valence Electrons) - (Non-bonding Electrons + 0.5 * Bonding Electrons)

The main idea is to help predict and represent the most accurate charge distribution within a molecule. It assists in identifying which resonance structure is more stable and likely to occur in nature.
In the context of the N₂O molecule, calculating the formal charge on each atom helps determine which arrangement of electrons across the atoms minimizes the formal charges. This gives us insight into which resonance structures contribute most to the molecule’s overall stability.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and play a vital role in chemical bonding. They are those electrons involved directly when atoms form bonds to create molecules. The number of valence electrons dictates how an atom will interact with others, making them key in determining how atoms form molecules.
For N₂O, calculating total available valence electrons is a critical initial step. Nitrogen has 5 valence electrons, and oxygen has 6. For two nitrogens and one oxygen, we have:

• 2 Nitrogen atoms = 2 x 5 = 10 valence electrons
• 1 Oxygen atom = 6 valence electrons
• Total = 16 valence electrons

Once the total valence electrons are counted, chemists distribute them according to bonding and lone pair needs, respecting the octet rule and minimizing formal charge ready to explore resonance structures. By doing so, one can demonstrate the different possible structures while staying within the chemical and physical constraints of the electron availability.

Octet Rule

The octet rule is a core concept in chemistry that states atoms aim to have eight electrons in their valence shell. This achieves a stable electronic configuration, similar to the noble gases, with filled s and p orbitals. The rule is a guiding principle when drafting chemical bonds and structures for molecules.
In the case of N₂O, knowing the octet rule helps distribute the electrons so each atom achieves the desired electronic configuration. Here, each atom will ideally share, gain, or lose electrons in bonding until it reaches eight valence electrons, creating a "full" outer shell.
All resonance structures must respect this trend, providing plausible scenarios of electron delocalization which satisfy the octet fulfillment. For resonance, one key step is ensuring that each structure respects this rule especially when adjusting electron distribution between bonds and lone pairs.
This helps chemists distinguish between possible resonance structures, evaluating stability based on how well they satisfy the octet. Particularly for central nitrogen or peripheral oxygen, such structures are deemed better if they preserve this electron total effectively and distribute formal charges efficiently across the molecule.