Problem 51 - Chemistry The Central Science In Si Units

Question

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: \((\mathbf{a}) \mathrm{OCS},(\mathbf{b}) \mathrm{SOCl}_{2}(\mathrm{~S}\) is the central atom), (c) \(\mathrm{BrO}_{3}^{-}\) (d) \(\mathrm{HClO}_{2}(\mathrm{H}\) is bonded to \(\mathrm{O})\).

Step-by-Step Solution

Answer

OCS: O=C=S; SOCl2: O=S(Cl)Cl; BrO3^-: BrO2(O); HClO2: HO-Cl=O. Each satisfies the octet rule with formal charges zero or adjusted for charge.

Step 1: Write the Lewis Structure for OCS

First, count the total valence electrons: O (6) + C (4) + S (6) = 16 electrons. To satisfy the octet rule, you arrange them in this order: O=C=S. Both oxygen and sulfur have full octets while carbon has two single bonds. The Lewis structure has no lone pairs on carbon, two lone pairs on oxygen, and two lone pairs on sulfur.

Step 2: Assign Oxidation Numbers and Formal Charges for OCS

For OCS, the oxidation numbers are: O (-2), C (+4), S (-2); for formal charges: O (0), C (0), S (0). All atoms fulfill the octet rule with no formal charge.

Step 3: Write the Lewis Structure for SOCl2

Count total valence electrons: S (6) + O (6) + 2Cl (14) = 26 electrons. Place sulfur in the center, oxygen forms a double bond with sulfur, and each chlorine forms a single bond. Oxygen has two lone pairs, sulfur has no lone pairs, and each chlorine has three lone pairs. All atoms follow the octet rule.

Step 4: Assign Oxidation Numbers and Formal Charges for SOCl2

The oxidation numbers are: S (+4), O (-2), Cl (0); formal charges: S (0), O (0), Cl (0) for each Cl. The structure adheres to the octet rule with no non-zero formal charges.

Step 5: Write the Lewis Structure for BrO3^-

Count total valence electrons: Br (7) + 3O (18) + 1 (due to charge) = 26 electrons. Bromine central, each oxygen forms a single bond except one which forms a double bond. Two oxygens have three lone pairs, and one has two lone pairs. Add an extra electron as this is an anion.

Step 6: Assign Oxidation Numbers and Formal Charges for BrO3^-

Assign oxidation numbers: Br (+5), O (-2) each. Assign formal charges: Br (0), two single-bonded O have a formal charge of -1, the double-bonded O has 0, making the total charge -1.

Step 7: Write the Lewis Structure for HClO2

Calculate total valence electrons for the molecule: H (1) + Cl (7) + 2O (12) = 20 electrons. Bond H to one O, which then bonds to Cl and Cl also bonds to the other O (double bond preferred for Cl–O). Each oxygen maintains its required octet.

Step 8: Assign Oxidation Numbers and Formal Charges for HClO2

Assign oxidation numbers: H (+1), Cl (+3), each O (-2). For formal charges: H (0), Cl (0), double-bonded O (0), single-bonded O (-1). All atoms fulfill the octet rule in this structure.

Key Concepts

Octet Rule

The Octet Rule is a guiding principle in chemistry that helps us predict the bonding behavior of atoms. It states that atoms tend to form bonds until they are surrounded by eight electrons, mimicking the electron configuration of noble gases. This rule is based on stability; atoms achieve greater stability when their outermost shell is full.

For example, when writing the Lewis structure for OCS:

• Oxygen, carbon, and sulfur all seek to complete their outer shell with eight electrons.
• Oxygen and sulfur each hold lone pairs to maintain their octet, while carbon does not have lone pairs but achieves its octet through bonding.

Sometimes, exceptions occur, particularly with elements that have access to d-orbitals, allowing for more than eight electrons. However, the Octet Rule is still a valuable tool for understanding molecular structures.

Oxidation Numbers

Oxidation Numbers are a bookkeeping method used in chemistry to keep track of electrons in chemical compounds. They help determine how electrons are distributed among atoms in molecules, especially when it comes to redox reactions.

For OCS, the oxidation numbers are:

• Oxygen is assigned -2 (since it generally gains electrons due to its high electronegativity).
• Carbon is assigned +4 (as it donates electrons to both oxygen and sulfur).
• Sulfur also has an oxidation number of -2.

Understanding oxidation numbers helps in balancing chemical equations and understanding the structure of the compounds. They are not always equal to formal charges but are often used together for deeper insight into molecule behavior.

Formal Charges

Formal Charges are calculated to determine the most stable arrangement of atoms in a molecule while writing Lewis structures. The formula is:\[\text{Formal Charge} = \text{Valence Electrons} - \text{Non-Bonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons}\] This calculation helps ensure that atoms are placed in optimal positions in a molecule.

For example, in the molecule SOCl\_2\:

• Each chlorine has a formal charge of 0, as does the central sulfur.
• The structure optimally shares valence electrons without any atom bearing unnecessary charges.

Utilizing formal charges allows chemists to identify the most stable structure by visualizing potential distribution of electrons more accurately.

Valence Electrons

Valence Electrons are the electrons in the outermost shell of an atom that are available for forming bonds with other atoms. They are crucial in determining how atoms interact and connect to form molecules.Identifying the correct number of valence electrons is the first step in drawing a Lewis structure. Here's how it works:

• Oxygen has 6 valence electrons, carbon has 4, sulfur also has 6.
• In molecules like \( ext{BrO}\_3^{-}\), the total count includes an extra electron due to the negative charge, making it 26 rather than 25 for cumulative bonding.

These electrons are either shared (in the case of covalent bonds) or transferred (in ionic bonds) to achieve stability as per the Octet Rule. Understanding valence electrons is fundamental when predicting how a molecule will form and behave.