Question

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

Short answer

(a) OCS Lewis Structure: O=C=S Formal Charges: Oxygen: +1, Carbon: +2, Sulfur: +1 Oxidation Numbers: Oxygen: -2, Carbon: +4, Sulfur: -2 (b) SOCl$_2$ Lewis Structure: ``` Cl | O = S-Cl ``` Formal Charges: Oxygen: +1, Sulfur: +2, Chlorine: 0 Oxidation Numbers: Oxygen: -2, Sulfur: +6, Chlorine: -1

Step-by-step solution

1. Total Number of Valence Electrons

Oxygen has 6 valence electrons, carbon has 4, and sulfur has 6. So the total number of valence electrons is 6 + 4 + 6 = 16.

2. Skeleton of the Molecule

Arrange the atoms and form skeletal bonds. Carbon is the central atom. Connect the two atoms in order: O-C-S.

3. Distribute Valence Electrons

We will now distribute the remaining valence electrons around the non-central oxygen and sulfur atoms: Oxygen has 6 - 2 = 4 electrons (two lone pairs) remaining. Sulfur has 6 - 2 = 4 electrons (two lone pairs) remaining. The Lewis structure for OCS is now: ``` O=C=S ```

4. Assign Formal Charges and Oxidation Numbers

To calculate the formal charge for each atom, use the formula: \[Formal\_Charge = Valence\_Electrons - (Bonds + Lone\_Electrons)\] Oxygen: 6 - (1 + 4) = 1 Carbon: 4 - (2 + 0) = 2 Sulfur: 6 - (1 + 4) = 1 Oxidation number: Oxygen: -2 Carbon: +4 Sulfur: -2 Now, let's move on to the second compound. (b) SOCl\(_2\) (S is the central atom)

1. Total Number of Valence Electrons

Sulfur has 6 valence electrons, oxygen has 6, and chlorine has 7. As there are two chlorine atoms, the total number of valence electrons is 6 + 6 + 7(2) = 26.

2. Skeleton of the Molecule

Sulfur is the central atom. Connect the central S atom to O and the two Cl atoms: S-O and S-Cl.

3. Distribute Valence Electrons

Distribute the remaining valence electrons around the non-central oxygen and chlorine atoms: Oxygen has 6 - 2 = 4 electrons (two lone pairs) remaining. Each chlorine atom has 7 - 1 = 6 electrons (three lone pairs) remaining. The Lewis structure for SOCl\(_2\) is now: ``` Cl | O = S-Cl ```

4. Assign Formal Charges and Oxidation Numbers

To calculate the formal charge for each atom, use the formula: Oxygen: 6 - (1 + 4) = 1 Sulfur: 6 - (4 + 0) = 2 Both Chlorine atoms: 7 - (1 + 6) = 0 (since the formal charges are the same for both Cl atoms) Oxidation number: Oxygen: -2 Sulfur: +6 Chlorine: -1 Complete the remaining two Lewis structures, formal charges, and oxidation numbers for the other compounds following the same steps.

Key concepts

Octet Rule

The octet rule is a guideline used in chemistry to predict the structure of molecules. It suggests that atoms are most stable when they have eight valence electrons in their outer shell, similar to the electron configuration of noble gases.
To achieve this stable configuration, atoms can share, donate, or receive electrons to form chemical bonds. This is the basis for creating Lewis structures, which illustrate how atoms connect in a molecule to satisfy the octet rule.
For example, in the molecule Oxygen Carbon Sulfide (OCS), carbon forms a double bond with both oxygen and sulfur, so all atoms can achieve an octet configuration. However, there are exceptions to the octet rule, such as hydrogen, which only requires two electrons, and certain transition metals that can have more than eight.

Formal Charge

Formal charge is a hypothetical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally. It's a tool used to predict the most stable structure for a molecule.
The formula for calculating formal charge is:\[\text{Formal Charge} = \text{Valence Electrons} - (\text{Bonds} + \text{Lone Electrons})\]
By calculating the formal charges, we can evaluate the stability of a Lewis structure. The best structures usually bear the lowest possible formal charges. For instance, in the OCS molecule, the formal charges help determine that the O=C=S configuration is reasonable.

• Carbon, with four valence electrons, would ideally have no formal charge when it shares four bonds.
• Oxygen and sulfur, both with six valence electrons, have formal charges that balance as close to zero as possible for stability.

Oxidation Numbers

Oxidation numbers, similar yet distinct from formal charges, reflect the hypothetical charge an atom would have if all bonds were 100% ionic. They are useful for understanding redox reactions and balancing chemical equations.
Assigning oxidation numbers involves applying a set of rules, including that the sum of oxidation numbers in a neutral molecule must be zero. In ions, they must equal the ionic charge.
For OCS, the oxidation numbers help confirm the molecule's electron distribution:

• Oxygen usually has an oxidation number of -2, since it is highly electronegative and prefers to gain electrons.
• Carbon, less electronegative, is given an oxidation number of +4 when it loses its four electrons for bonding.
• Sulfur typically has an oxidation number of -2, similar to oxygen in this context.

These numbers reassure that electron sharing is consistent throughout the molecule.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in chemical bonding. They determine how an atom interacts with others to form a molecule.
When constructing Lewis structures, knowing the number of valence electrons for each atom helps predict the molecule's formation. For example, in constructing OCS, you total the valence electrons to determine the bonding possibilities:

• Oxygen: 6 valence electrons, aiming to reach an octet.
• Carbon: 4 valence electrons, forms four bonds to fill its outer shell.
• Sulfur: 6 valence electrons, sharing and creating bonds to meet the octet rule.

This understanding of valence electrons guides how atoms bond within molecules, supporting the depiction of accurate Lewis structures.