Question

Write a Lewis formula for each of the following: (a) \(\mathrm{NO}_{3}{ }^{-}\) (c) \(\mathrm{SO}_{3}{ }^{2-}\) (e) \(\mathrm{NO}^{+}\) (b) \(\mathrm{SO}_{4}{ }^{2-}\) (d) \(\mathrm{NO}_{2}^{-}\)

Short answer

The Lewis structure for \(\mathrm{NO}_{3}{ }^{-}\) shows a nitrogen atom single-bonded to three oxygen atoms, with one of the oxygen atoms also bonded to the nitrogen atom via a double bond. There are resonance structures. Repeat the process to get Lewis structures for each of \(\mathrm{SO}_{3}{ }^{2-}\), \(\mathrm{NO}^{+}\), \(\mathrm{SO}_{4}{ }^{2-}\), and \(\mathrm{NO}_{2}^{-}\).

Step-by-step solution

Electron Counting

Count the total number of valence electrons that each atom in the molecule or polyatomic ion contributes. Remember that any change in charge should be accounted for. For instance, the \(\mathrm{NO}_{3}{ }^{-}\) ion will have 24 valence electrons, as nitrogen contributes 5, each oxygen contributes 6 and the extra electron contributes 1.

Initial Structure Drawing

Draw a structure by connecting the atoms with single bonds. The least electronegative atom is usually the central atom. In \(\mathrm{NO}_{3}{ }^{-}\), N is less electronegative than O, so it is central. This uses up 2 electrons per bond, so a total of 6 electrons are used.

Electron Distribution

Distribute the remaining electrons as lone pairs. The goal is to fill the octets of the atoms. For \(\mathrm{NO}_{3}{ }^{-}\), distribute 18 electrons as lone pairs (9 pairs), starting with the outer atoms (O), which will now each have 6 electrons (3 lone pairs).

Resonance Structures

If the central atom does not have a complete octet, create multiple resonance structures. For \(\mathrm{NO}_{3}{ }^{-}\), nitrogen only has 4 electrons, so one of the lone pairs on an oxygen atom can be used to create a double bond. This creates 3 equivalent resonance structures.

Steps 5-8: Repeat for Other Ions

Repeat steps 1 through 4 for each of the other ions, \(\mathrm{SO}_{3}{ }^{2-}\), \(\mathrm{NO}^{+}\), \(\mathrm{SO}_{4}{ }^{2-}\), and \(\mathrm{NO}_{2}^{-}\). Use the same rules to count electrons, draw structures, distribute electrons, and identify resonance structures.

Key concepts

Electron Counting

Counting electrons is the first and perhaps the most crucial step in creating a Lewis structure. Each element is characterized by a certain number of valence electrons, which are the outermost electrons involved in bonding. For instance, in constructing the Lewis structure for the nitrate ion, \(\mathrm{NO}_{3}^{-}\), nitrogen contributes its 5 valence electrons. Each of the three oxygen atoms adds 6 valence electrons. Because the ion carries a negative charge, an additional electron is added. This results in a total of 24 valence electrons for the molecule.

**Important Considerations for Electron Counting:**

• Always include electrons reflective of charges on ions. Add electrons for negative charges and subtract for positive ones.
• Sum the valence electrons of all atoms involved in the compound.
• Consider that the central atom often has fewer valence electrons and tends to be less electronegative. This helps in structuring the initial bonding framework.

Resonance Structures

Resonance structures are different ways of representing the same molecule. In certain molecules, like \(\mathrm{NO}_{3}^{-}\), electrons can be distributed in various configurations, leading to multiple valid structures. These structures don't actually exist separately; rather, the true structure is a hybrid or average of all possible resonance forms.

**Understanding Resonance Better:**

• Resonance primarily addresses the distribution of electrons, not the positions of the atoms.
• Structures are often equivalent and reduce the formal charge of atoms involved, creating more stability.
• The resonance hybrid has properties that reflect aspects of each contributing structure, especially concerning bond lengths and energy levels.

Remember, resonance is a fundamental concept when dealing with molecules that have delocalized electrons, enhancing the molecule’s stability beyond a single structural representation.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. They are crucial in the formation of chemical bonds between atoms. Valence electrons determine how atoms interact and bond with other atoms to form different compounds. In the context of drawing Lewis structures, valence electrons are represented as either pairs or as lone dots around elemental symbols.

**Key Points about Valence Electrons:**

• The number of valence electrons corresponds to an atom’s position in the periodic table. For instance, oxygen, in group 16, always has 6 valence electrons.
• During bonding, atoms generally aim to fulfill the octet rule, seeking 8 electrons in their valence shell, like noble gases.
• Knowing the number of valence electrons helps predict molecular shape, bond numbers, and whether atoms will participate in single, double, or triple bonds.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It is crucial for understanding the physical and chemical properties of molecules. While Lewis structures give us a good insight into bonding and molecular layout, they are primarily two-dimensional. Geometries like linear, trigonal planar, tetrahedral, among others, provide a more comprehensive understanding.

**Diving Deeper into Molecular Geometry:**

• Geometry is determined by the number of bonding groups and lone pairs around a central atom. For example, \(\mathrm{NO}_{3}^{-}\) is trigonal planar because it has three bonding regions around nitrogen and no lone pair on the nitrogen atom.
• The concept of electron pair repulsion helps predict shapes. This theory states that electron pairs, both bonding and lone pairs, repel each other to achieve the maximum distance apart.
• Understanding molecular geometry is vital for anticipating molecular reactivity and interaction, crucial in fields like pharmacology and material science.