Question

Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) OCS; (b) NO.

Short answer

For OCS: O=C=S with all formal charges 0. For NO: N≡O with N at +1 and O at 0.

Step-by-step solution

Count Valence Electrons for OCS

Determine the total number of valence electrons by adding the valence electrons of oxygen (O), carbon (C), and sulfur (S).O: 6 valence electrons C: 4 valence electrons S: 6 valence electrons Total = 6 + 4 + 6 = 16 valence electrons.

Draw Initial Skeleton Structure for OCS

Place the least electronegative atom (carbon) in the center and connect it with oxygen and sulfur using single bonds: O–C–S.

Distribute Remaining Electrons for OCS

Place the remaining valence electrons around the atoms to complete their octets. There are 4 electrons used in the bonds, leaving 12 electrons (16-4 = 12):- Place 6 electrons (3 pairs) around O.- Place 6 electrons (3 pairs) around S.

Form Double Bonds to Minimize Formal Charges for OCS

To minimize formal charges, form double bonds. Move lone pair electrons from O to form O=C and from S to form C=S:O=C=S.

Calculate Formal Charges for OCS

Formal Charge = (valence electrons) - (non-bonding electrons) - (bonding electrons/2)- For O: (6) - (4) - (4/2) = 0- For C: (4) - (0) - (8/2) = 0- For S: (6) - (4) - (4/2) = 0All atoms have a formal charge of 0.

Count Valence Electrons for NO

Determine the total number of valence electrons by adding the valence electrons of nitrogen (N) and oxygen (O).N: 5 valence electrons O: 6 valence electrons Total = 5 + 6 = 11 valence electrons.

Draw Initial Skeleton Structure for NO

Connect nitrogen and oxygen with a single bond: N–O.

Distribute Remaining Electrons for NO

Place the remaining valence electrons around the atoms to complete their octets. There are 2 electrons used in the bond, leaving 9 electrons (11-2 = 9):- Place 6 electrons (3 pairs) on O.- Place the remaining 3 electrons on N.

Form Bond to Minimize Formal Charges for NO

To minimize formal charges and complete octets, form a double bond and place a lone electron on N. Move electrons to form N=O with a lone electron on N:N≡O.

Calculate Formal Charges for NO

Formal Charge = (valence electrons) - (non-bonding electrons) - (bonding electrons/2)- For N: (5) - (1) - (6/2) = +1- For O: (6) - (4) - (6/2) = 0The formal charges are: N = +1, O = 0.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom, and they play a crucial role in chemical bonding. For instance, they determine how an atom will bond with others to form molecules.
Each element in the periodic table has a specific number of valence electrons. For example:

• Oxygen (O): 6 valence electrons
• Carbon (C): 4 valence electrons
• Sulfur (S): 6 valence electrons

Knowing how to count valence electrons helps in drawing the Lewis structure of molecules, as it dictates the possible bonding scenarios and the arrangement of electrons around atoms.

Formal Charge

Formal charge is an important concept to ensure that Lewis structures reflect the lowest energy and most stable form of a molecule.
It is calculated using the formula:
\begin{array} {c} Formal \ charge = (valence \ electrons) - (non-\ bonding \ electrons) - (bonding \ electrons/2) \ \times \end{array}
Here's how it applies:
For an oxygen atom in OCS: \[\begin{equation} Formal \ charge = 6 - 4 - 4/2 = 0 \ \times \end{equation}\]
For a nitrogen atom in NO: \[\begin{equation} Formal \ charge = 5 - 1 - 6/2 = +1 \ \times \end{equation}\]
Calculating the formal charge helps confirm the most stable structure of a molecule by making sure the charges are as close to zero as possible.

Molecular Structure

The molecular structure of a compound refers to how the atoms are arranged and bonded together. It influences the molecule's properties and behavior in chemical reactions.
When constructing a molecular structure:

• The least electronegative atom usually occupies the central position.
• Bonds are formed using valence electrons.
• The remaining electrons are used to complete the octets of the surrounding atoms.

This process leads to the realization of different bond types, such as single, double, or triple bonds, to achieve the most stable configuration.

OCS Molecule

The OCS molecule consists of oxygen (O), carbon (C), and sulfur (S).
The steps to draw its Lewis structure are:

• Count valence electrons: 6 (O) + 4 (C) + 6 (S) = 16 valence electrons.
• Place the least electronegative atom (C) in the center and bond with O and S: O-C-S.
• Distribute remaining electrons to complete octets: 3 pairs around O, 3 pairs around S.
• Form double bonds to minimize formal charges: O=C=S.
• Calculate formal charges to confirm stability: All atoms have a formal charge of 0.

The final structure shows O, C, and S with no formal charge, indicating a stable molecule.

NO Molecule

The NO molecule is composed of nitrogen (N) and oxygen (O).
The steps to draw its Lewis structure are:

• Count valence electrons: 5 (N) + 6 (O) = 11 valence electrons.
• Connect N and O with a single bond: N-O.
• Distribute remaining electrons: 3 pairs around O, 1 pair and 1 lone electron on N.
• Form double bonds and adjust electrons: N≡O with a lone electron on N.
• Calculate formal charges: N = +1, O = 0.

The final structure has a triple bond between N and O to minimize formal charges and yield the most stable form.