Question

Write Lewis structures for the following. Show all resonance structures where applicable. a. \(\mathrm{NO}_{2}^{-}, \mathrm{NO}_{3}^{-}, \mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{N}_{2} \mathrm{O}_{4} \text { exists as } \mathrm{O}_{2} \mathrm{N}-\mathrm{NO}_{2} .\right)\) b. \(\mathrm{OCN}^{-}, \mathrm{SCN}^{-}, \mathrm{N}_{3}^{-}\) (Carbon is the central atom in \(\mathrm{OCN}^{-}\) and \(\mathrm{SCN}^{-} . )\)

Short answer

In the short answer, we have drawn the Lewis structures for the given compounds: a. 1. \(\mathrm{NO}_{2}^{-}\): N has a single bond to one O and a double bond to the other O. 2. \(\mathrm{NO}_{3}^{-}\): N has a double bond to one O and single bonds to the other two O atoms; there are two resonance structures. 3. \(\mathrm{N}_{2}\mathrm{O}_{4}\): Exists as \(\mathrm{O}_{2}\mathrm{N-NO}_{2}\), where one N atom has single bonds to two O atoms and the other N atom has a double bond to an O atom. b. 1. \(\mathrm{OCN}^{-}\): C has a double bond to O and a triple bond to N; resonance structure with triple bond to O and double bond to N. 2. \(\mathrm{SCN}^{-}\): C has a double bond to S and a triple bond to N; three resonance structures possible with different bond orders between C, S, and N. 3. \(\mathrm{N}_{3}^{-}\): Central N atom has a single bond to one terminal N atom and a triple bond to the other terminal N atom; two resonance structures for the different triple-bonded terminal N atoms.

Step-by-step solution

1. Drawing the Lewis Structure for \(\mathrm{NO}_{2}^{-}\)

1. Total valence electrons: N (5) + 2 × O (6 each) + 1 (for the negative charge) = 18. 2. Nitrogen is the central atom. 3. Form single bonds between N and O (2 × 2 = 4 electrons used). 4. Complete the octet for the outer O atoms (2 × 6 = 12 electrons used). 5. Remaining electrons = 18 - 16 = 2, place these on the central N atom. 6. At this point, the N atom has 6 electrons, so it forms a double bond with one of the O atoms to satisfy the octet rule. 7. No resonance structures are required for this compound.

2. Drawing the Lewis Structure for \(\mathrm{NO}_{3}^{-}\)

1. Total valence electrons: N (5) + 3 × O (6 each) + 1 (for negative charge) = 24. 2. Nitrogen is the central atom. 3. Form single bonds between N and each O atom (3 × 2 = 6 electrons used). 4. Complete the octet for the outer O atoms (3 × 6 = 18 electrons used). 5. Remaining electrons = 24 - 24 = 0, no need to place any extra electrons. 6. Form double bonds with one of the O atoms to satisfy the octet rule for nitrogen. 7. There are two resonance structures for this compound, corresponding to the different O atoms that can form the double bond with nitrogen.

3. Drawing the Lewis Structure for \(\mathrm{N}_{2}\mathrm{O}_{4}\)

1. Total valence electrons: 2 × N (5 each) + 4 × O (6 each) = 34. 2. The central atoms are the two N atoms. 3. \(\mathrm{N}_{2}\mathrm{O}_{4}\) exists as \(\mathrm{O}_{2}\mathrm{N-NO}_{2}\), so we will form single bonds between the N atoms and each O atom (4 × 2 = 8 electrons used). 4. Complete the octets for the outer O atoms (4 × 6 = 24 electrons used). 5. Remaining electrons = 34 - 32 = 2, place these on the N atom not bonded to the other N atom. 6. The N atom bonded to the other N atom has 6 electrons, so it forms a double bond with one of its O atoms to satisfy the octet rule. 7. There are no resonance structures required for this compound. #exercise b# We will draw the Lewis structures for the following compounds: 1. \(\mathrm{OCN}^{-}\) 2. \(\mathrm{SCN}^{-}\) 3. \(\mathrm{N}_{3}^{-}\)

4. Drawing the Lewis Structure for \(\mathrm{OCN}^{-}\)

1. Total valence electrons: O (6) + C (4) + N (5) + 1 (for negative charge) = 16. 2. Carbon is the central atom. 3. Form single bonds between C and the O and N atoms (2 × 2 = 4 electrons used). 4. Complete the octets for the outer O and N atoms (6+6 = 12 electrons used). 5. Remaining electrons = 16 - 16 = 0, no need to place any extra electrons. 6. Form double/triple bonds with remaining electron pairs: There are two possible combinations: triple bond between C and N, double bond between C and O. This leads to a resonance structure.

5. Drawing the Lewis Structure for \(\mathrm{SCN}^{-}\)

1. Total valence electrons: S (6) + C (4) + N (5) + 1 (for negative charge) = 16. 2. Carbon is the central atom. 3. Form single bonds between C and the S and N atoms (2 × 2 = 4 electrons used). 4. Complete the octets for the outer S and N atoms (6+6 = 12 electrons used). 5. Remaining electrons = 16 - 16 = 0, no need to place any extra electrons. 6. Form double/triple bonds with remaining electron pairs. There are three different resonance structures: a triple bond between C and N, a double bond between C and S, a triple bond between C and S, and a double bond between C and N.

6. Drawing the Lewis Structure for \(\mathrm{N}_{3}^{-}\)

1. Total valence electrons: 3 × N (5 each) + 1 (for negative charge) = 16. 2. Two terminal N atoms and one central N atom. 3. Form single bonds between the central N atom and the two terminal N atoms (2 × 2 = 4 electrons used). 4. Complete the octets for the terminal N atoms (2 × 6 = 12 electrons used). 5. Remaining electrons = 16 - 16 = 0, no need to place any extra electrons. 6. Form a triple bond between the central N atom and one of the terminal N atoms (there are two resonance structures for this compound).

Key concepts

Resonance Structures

Resonance structures are a fascinating aspect of chemistry. They show different possible ways that electrons can be arranged in a molecule, which can affect the molecule's properties. Imagine you’re trying to capture a photo of a dancing figure. Each photo looks slightly different depending on the moment captured. Resonance structures are like these photos. They represent various "snapshots" of electron distribution in a molecule.For compounds like \(NO_3^{-}\) and \(SCN^{-}\), electrons can form different bonding patterns. Multiple resonance structures provide a more accurate description of the molecule's true electronic structure. In practice, this means that the actual structure is a blend, or "resonance hybrid," of all possible forms. Understanding resonance structures can help predict the molecule's stability and reactivity. It provides valuable insight into how electrons move and bond within the molecule.

Valence Electrons

Valence electrons are the key players in forming chemical bonds. They are the electrons present in the outermost shell of an atom. These electrons are crucial because they determine an atom's ability to bond with others.Think of valence electrons as social butterflies at a party. They are always looking to "mingle" with electrons from other atoms. By sharing or transferring these electrons, atoms can achieve a stable electron configuration, similar to having a full dance card at a lively party. When drawing Lewis structures, counting the total valence electrons is the first step. This helps in distributing the electrons around atoms in a molecule. For instance, in \(NO_2^{-}\), nitrogen contributes five valence electrons, while oxygen provides six. The additional electron is due to the negative charge.Recognizing the number of valence electrons involved allows us to predict potential bonding scenarios and is fundamental to understanding molecular structures.

Octet Rule

The octet rule is a simple yet powerful concept in chemistry. It states that atoms tend to form bonds until they are surrounded by eight valence electrons. This rule is based on the idea that having a full outer shell leads to greater stability.You can think of the octet rule as a quest for satisfaction. Atoms bond, share, or exchange electrons to complete their valence shell, much like gathering enough companions to fill a dinner table.In Lewis structures, ensuring each atom meets the octet rule is crucial. For example, in \(NO_3^{-}\), nitrogen forms a double bond with one oxygen atom so that each atom fulfills the rule. While there are exceptions, particularly with elements in the third period and beyond, the octet rule provides a helpful guideline for predicting molecular shapes and stability.Applying the octet rule effectively can provide a clearer visual representation of how atoms "connect" and stabilize one another in a compound.

Central Atom

Choosing the central atom in a molecule is a smart strategy when drawing Lewis structures. Usually, the atom with the lowest electronegativity or the one that can form the most bonds is selected.Visualize the central atom as the "hub" of a network. It's like the main character at a party who holds all social links, or bonds, together. It forms the core around which other atoms arrange themselves.For instance, in \(OCN^{-}\) and \(SCN^{-}\), carbon serves as the central atom. Despite having fewer valence electrons than oxygen or sulfur, carbon can easily form multiple bonds with other atoms, stabilizing the structure.Understanding the concept of a central atom is vital in predicting how molecules are structured and function, offering insights on how atoms prefer to associate in complex chemical systems.

Double Bond

Double bonds are a common feature in many molecular structures. They are formed when two atoms share two pairs of electrons.Imagine a double bond as a stronger handshake at a business meeting, signaling a deeper connection between two parties. They not only hold atoms together but also influence the molecule's shape and properties.Double bonds are crucial in satisfying the octet rule. For instance, nitrogen in \(NO_2^{-}\) forms a double bond with one oxygen to satisfy its requirement for eight electrons around it. This connection is tighter and often alters the energy, reactivity, and overall geometry of the molecule.Grasping the concept of double bonds is essential for understanding how more complex molecules form and behave. They provide valuable insight into the dynamics of molecular interactions.