Problem 60 - Chemistry Principles and Practice

Question

Write the Lewis structures showing formal charge for the following species. (a) \(\mathrm{ClO}_{3}^{-}\) (b) NCCN (c) \(\mathrm{SOCl}_{2}\)

Step-by-Step Solution

Answer

The Lewis structures with formal charges are: (a) \\([\mathrm{Cl}^{+1} = \mathrm{O}^{-1} = \mathrm{O}^{0} = \mathrm{O}^{-1}](\mathrm{ClO}_3^{-})\\) (b) \\([\mathrm{N} \equiv \mathrm{C} - \mathrm{C} \equiv \mathrm{N}](\mathrm{NCCN})(0 \text{ charge each})\\) (c) \\([\mathrm{Cl} - \mathrm{S}^{0} = \mathrm{O}^{0} - \mathrm{Cl}](\mathrm{SOCl}_2)\\)

Step 1: Count Total Valence Electrons for ClO3-

Calculate the total number of valence electrons. Chlorine (Cl) has 7 valence electrons, each oxygen (O) has 6, and the negative charge contributes an extra electron. Thus, the total is: \(7 + (3 \times 6) + 1 = 26\) valence electrons.

Step 2: Draw the Lewis Structure for ClO3-

Position Cl in the center with three O atoms surrounding it. Create single bonds between Cl and each O atom. Distribute remaining electrons to satisfy the octet rule, placing double bonds as needed to minimize formal charges. Adjust to form one double bond with one O atom, leaving two single bonded O atoms.

Step 3: Calculate Formal Charges for ClO3-

Determine formal charges using the formula: \(\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + 0.5 \times \text{Bonding electrons})\). Chlorine ends with a formal charge of +1, two oxygen atoms have a formal charge of -1, and one oxygen has a formal charge of 0.

Step 4: Count Total Valence Electrons for NCCN

Each nitrogen has 5 valence electrons and each carbon has 4. Total is: \(2 \times 5 + 2 \times 4 = 18\) valence electrons.

Step 5: Draw the Lewis Structure for NCCN

Arrange N-C-C-N linearly. Use a triple bond between each C and N to satisfy the octet rule. Place a single bond between the two C atoms. Assign any remaining electrons to complete the octets on terminal atoms.

Step 6: Calculate Formal Charges for NCCN

For each element, use the formula: \(\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + 0.5 \times \text{Bonding electrons})\). Both nitrogen atoms end with a formal charge of 0, as do the carbon atoms.

Step 7: Count Total Valence Electrons for SOCl2

Sulfur (S) has 6 valence electrons, oxygen has 6, and each chlorine has 7. Total is: \(6 (S) + 6 (O) + 2 \times 7 (Cl) = 26\) valence electrons.

Step 8: Draw the Lewis Structure for SOCl2

Place S in the center and connect it to O and both Cl atoms with single bonds. Add remaining electrons primarily to outer atoms to complete their octets and then adjust bonding on S-O to form a double bond if needed to satisfy the octet.

Step 9: Calculate Formal Charges for SOCl2

Apply \(\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + 0.5 \times \text{Bonding electrons})\). Sulfur generally shows a formal charge of 0 when correct bonding is applied with a double bond to oxygen; both oxygens and chlorines will also show neutral formal charges.

Key Concepts

formal charges

Formal charges help us determine the most stable structure by assigning an apparent charge to each atom within a molecule. To calculate the formal charge, you use the formula: \[\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + 0.5 \times \text{Bonding electrons})\] This allows us to check if everything is balanced. Ideally, the formal charges on the atoms should be as close to zero as possible or match the charge of the ion you are working with.

• Positive formal charges are assigned to atoms that effectively lose electrons in bonding.
• Negative formal charges are observed on atoms gaining electrons compared to their usual state.
• Zero charge indicates a neutral, balanced atom in the molecular structure.

Adjusting bonds, such as creating double bonds where needed, often helps balance these charges. This is an essential step to ensure the stability and accuracy of the Lewis structure.

valence electrons

Valence electrons are the electrons located in the outermost shell of an atom and play a key role in chemical bonding and reactions. They determine how an atom will interact or bond with other atoms.

• The number of valence electrons can typically be found in the element's group number on the periodic table.
• These electrons are the ones that get shared, lost, or gained in chemical reactions, leading to the formation of molecules.

When drawing Lewis structures, properly counting the total number of valence electrons is crucial to represent the molecule accurately. For example, when considering the molecule \(\text{ClO}_3^{-}\), the total count includes chlorine's valence electrons, the contribution from each oxygen, and the additional electron due to the molecule's negative charge. Calculating valence electrons helps ensure that all atoms in the molecule can achieve a stable configuration.

octet rule

The octet rule is a guideline in chemistry where atoms tend to bond in a manner that gives each of them eight valence electrons, resembling the electron configuration of a noble gas. This rule explains why -

• Atoms bond together, either by sharing (covalent bonding) or transferring (ionic bonding) valence electrons.
• The octet rule makes atoms reach a more stable state, reducing their energy.
• Elements like hydrogen are exceptions, aiming for two valence electrons to achieve stability similar to helium.

In Lewis structures, following the octet rule is essential for most atoms, like oxygen and chlorine in \(\text{ClO}_3^{-}\). However, beware that molecules with odd numbers of electrons and some transition metals may not always adhere to the octet rule. Always start by arranging the bonds to satisfy this rule for as many atoms as possible while fine-tuning as needed to accommodate all formal charge requirements for the whole structure.