Problem 5 - Chemistry and Chemical Reactivity

Question

Draw a Lewis structure for each of the following molecules or ions. (a) \(\mathrm{NF}_{3}\) (b) \(\mathrm{ClO}_{3}\) (c) HOBr (d) \(\mathrm{SO}_{3}^{2-}\)

Step-by-Step Solution

Answer

NF3: Octets complete, lone pairs on N and F. ClO3: Add single/double bonds for octets. HOBr: O in center, octets full. SO3^2-: Check charge and form double bonds.

Step 1: Count the Valence Electrons for NF3

Nitrogen has 5 valence electrons, and each fluorine has 7 valence electrons. Therefore, the total number of valence electrons is \(5 + 3 \times 7 = 26\).

Step 2: Draw the Skeleton Structure for NF3

Place nitrogen in the center and connect it to three fluorine atoms with single bonds.

Step 3: Distribute Remaining Electrons for NF3

After forming three \(N-F\) bonds, 6 electrons are used, leaving 20. Distribute the remaining 20 electrons as lone pairs on fluorine atoms (6 each) and the nitrogen (2). Ensure each atom satisfies the octet rule.

Step 4: Count the Valence Electrons for ClO3

Chlorine has 7 valence electrons, and oxygen has 6. Thus, the total number of valence electrons is \(7 + 3 \times 6 = 25\).

Step 5: Adjust for Charge in ClO3

As ClO3 is a polyatomic ion, add an electron to account for the negative charge, resulting in 26 valence electrons.

Step 6: Draw the Skeleton Structure for ClO3

Place chlorine in the center and connect it to three oxygen atoms with single bonds.

Step 7: Distribute Remaining Electrons for ClO3

After bonding, 6 electrons are used, leaving 20. Distribute remaining electrons to fulfill oxygen's octet. Put extra pairs on chlorine if required, forming double bonds if necessary.

Step 8: Count the Valence Electrons for HOBr

Hydrogen has 1 valence electron, oxygen has 6, and bromine has 7, totaling \(1 + 6 + 7 = 14\).

Step 9: Draw the Skeleton Structure for HOBr

Place oxygen in the center since hydrogen can only form one bond, and bond it to hydrogen and bromine.

Step 10: Distribute Remaining Electrons for HOBr

Assign the remaining 10 electrons to complete the octet for oxygen and bromine (8 for O, 2 for Br completing octet) with no formal charges.

Step 11: Count the Valence Electrons for SO3^2-

Sulfur has 6 valence electrons, and each oxygen has 6. Thus, the total is \(6 + 3\times6 = 24\).

Step 12: Adjust for Charge in SO3^2-

Add two extra electrons for the \(2-\) charge, totaling 26 valence electrons.

Step 13: Draw the Skeleton Structure for SO3^2-

Place sulfur in the center and connect it to three oxygen atoms with single bonds.

Step 14: Distribute Remaining Electrons for SO3^2-

Use remaining electrons to satisfy oxygen's octet, forming double bonds if necessary. Place any leftover electrons on sulfur.

Key Concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom that participate in chemical bonding. For a given element, you can find the number of valence electrons by looking at its group number in the periodic table. For example:
- Nitrogen ( f{N} ) has 5 valence electrons.
- Fluorine ( f{F} ) has 7 valence electrons.
- Chlorine ( f{Cl} ) also has 7 valence electrons.

Understanding valence electrons is essential to drawing Lewis structures since these electrons form bonds and lone pairs. More specifically, knowing how many valence electrons each atom has helps you determine how atoms bond together. To count their total in a molecule, sum up all valence electrons from each atom. For example, in f{NF_{3}} , nitrogen contributes 5 valence electrons, and each fluorine adds 7. Therefore, the total is 26 valence electrons. Counting valence electrons is the first step in drawing Lewis structures and sets the stage for understanding how the octet rule is applied. It helps visualize the interactions among atoms, guiding you in constructing both skeleton structures and any subsequent adjustments needed to satisfy all atoms.

Octet Rule

The "Octet Rule" is a crucial guideline in chemistry stating that atoms tend to form bonds to have eight electrons in their valence shell, resembling the electron configuration of noble gases. This rule is particularly relevant for main-group elements. For instance:
- In f{NF_{3}} , after forming three f{N-F} bonds, nitrogen ends up with a complete octet due to shared electrons.
- Similarly, each fluorine atom attains an octet through bonding and its lone pairs.

The octet rule helps predict the chemical reactivity and bonding of elements. However, there are exceptions. Some atoms can exceed the octet due to available d-orbitals, as seen in elements starting from period 3. Additionally, molecules or ions with an odd number of electrons or those where achieving an octet is impossible may violate this rule.

While drawing Lewis structures, ensure each atom fulfills the octet rule where applicable by distributing electrons around atoms in such a way that each achieves a stable configuration by either sharing, transferring, or pooling electrons. This helps to determine Lewis structures that best describe the stable form of molecules or ions.

Polyatomic Ions

Polyatomic ions are charged species composed of two or more atoms covalently bonded, carrying a net positive or negative charge. Examples include f{ClO_{3}^{-}} and f{SO_{3}^{2-}} . Understanding polyatomic ions is crucial for drawing accurate Lewis structures, particularly for adjusting electron counts to account for charge.

When calculating total valence electrons for polyatomic ions:

• Add electrons if the ion is negatively charged.
• Subtract electrons if it's positively charged.

For f{ClO_{3}^{-}} , chlorine and oxygen contribute 25 electrons together. The negative charge adds one more, resulting in 26 electrons. Likewise, in f{SO_{3}^{2-}} , the sulfur and oxygen atoms initially contribute 24 electrons. To account for the f{2-} charge, add two electrons, totaling 26. Once adjusted, these electrons are used to form bonds and distribute lone pairs, maintaining the integrity of the octet rule whenever possible. These steps ensure a complete and balanced Lewis structure reflecting the true nature of these charged entities.

Formal Charges

Formal charge is a theoretical charge assigned to atoms within a molecule or ion, indicating if electrons are evenly shared in bonds. Calculating formal charges helps determine the most stable Lewis structure configuration.

The formula for formal charge is:\[\text{Formal charge} = \left( \text{Valence electrons} - \text{(Non-bonding electrons + } \frac{1}{2} \text{Bonding electrons)} \right)\]Consider the example of f{HOBr}: The formal charge calculation can guide adjustments in the structure to minimize charges across atoms. Ideally, the most stable Lewis structure has formal charges as close to zero as possible for each atom.

When drawing structures for f{SO_{3}^{2-}}, use formal charge calculations to justify double bonds and electron placements, ensuring the negative charge is delocalized if applicable. Good practice is to strive for a formal charge distribution where the most electronegative atoms can bear additional electrons. Looking at formal charges ensures the drawing accurately reflects the most stable arrangement of electrons, optimizing molecular stability.