Question

How many nonbonding electron pairs are there in each of the following molecules: (a) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S} ;\) (b) \(\mathrm{HCN}\); (c) \(\mathrm{H}_{2} \mathrm{C}_{2} ;\) (d) \(\mathrm{CH}_{3} \mathrm{~F}\) ?

Short answer

The number of nonbonding electron pairs in each molecule is as follows: (a) \((\mathrm{CH}_{3})_{2} \mathrm{~S}\): 2 non-bonding electron pairs; (b) \(\mathrm{HCN}\): 1 non-bonding electron pair; (c) \(\mathrm{H}_{2} \mathrm{C}_{2}\): 0 non-bonding electron pairs; (d) \(\mathrm{CH}_{3} \mathrm{~F}\): 3 non-bonding electron pairs.

Step-by-step solution

Lewis Structure of the Molecules

Draw the Lewis structure for each molecule. Represent each atom with its element symbol and valence electrons around it. Connect the atoms using single, double, or triple bonds according to the valence electrons of each atom.

Count Non-Bonding Electron Pairs

After drawing the Lewis structures, count the non-bonding electron pairs for each atom in the molecule and find the total number of nonbonding electron pairs in each molecule. (a) \((\mathrm{CH}_{3})_{2} \mathrm{~S}\)

Lewis Structure

The Lewis structure for \((\mathrm{CH}_{3})_{2} \mathrm{~S}\) is: H H \ / H-C-S-C-H \ / H H

Count Non-Bonding Electron Pairs

In this molecule, the central Sulfur atom has two non-bonding electron pairs, and all other atoms have no non-bonding electron pairs. Therefore, there are 2 non-bonding electron pairs. (b) \(\mathrm{HCN}\)

Lewis Structure

The Lewis structure for \(\mathrm{HCN}\) is: H-C ≡ N

Count Non-Bonding Electron Pairs

In this molecule, the H and C atoms have no non-bonding electron pairs. The N atom has one non-bonding electron pair. Therefore, there is 1 non-bonding electron pair. (c) \(\mathrm{H}_{2} \mathrm{C}_{2}\)

Lewis Structure

The Lewis structure for \(\mathrm{H}_{2} \mathrm{C}_{2}\) is: H-C ≡ C-H

Count Non-Bonding Electron Pairs

In this molecule, all the atoms have no non-bonding electron pairs. Therefore, there are 0 non-bonding electron pairs. (d) \(\mathrm{CH}_{3} \mathrm{~F}\)

Lewis Structure

The Lewis structure for \(\mathrm{CH}_{3} \mathrm{~F}\) is: H H \ / H-C-F \ H

Count Non-Bonding Electron Pairs

In this molecule, the central F atom has three non-bonding electron pairs, and all other atoms have no non-bonding electron pairs. Therefore, there are 3 non-bonding electron pairs.

Key concepts

Non-bonding Electron Pairs

Non-bonding electron pairs, also known as lone pairs, are pairs of valence electrons that are not shared with another atom in a chemical bond. They belong to a single atom and play a crucial role in the shape and reactivity of molecules. For instance, in the molecule \((\text{CH}_3)_2 \text{S}\), the sulfur atom has two non-bonding electron pairs. These lone pairs are crucial because they can affect the geometry of the molecule by repelling bonding pairs, which influences the overall shape of the molecule.
In the molecule \(\text{HCN}\), nitrogen, an atom with a high electronegativity, possesses one non-bonding pair. Such lone pairs can also be involved in hydrogen bonding and influence a molecule's physical properties. Meanwhile, in \(\text{H}_2\text{C}_2\), there are no lone pairs on either hydrogen or carbon, which results in a straight molecular geometry.
Lastly, in \(\text{CH}_3\text{F}\), the fluorine atom holds three lone pairs. These pairs impact both the molecule's polarity and its three-dimensional configuration. Understanding the role of lone pairs helps in predicting and explaining the behavior and properties of molecules.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. These electrons are essential because they determine how an atom interacts or forms bonds with others. Drawing Lewis structures helps to visualize these valence electrons and determine how many are available for bonding.
For example, sulfur in \((\text{CH}_3)_2 \text{S}\) has six valence electrons. It uses two for bonding with carbon atoms, leaving four as non-bonding electron pairs or lone pairs. Similarly, nitrogen in \(\text{HCN}\) has five valence electrons; it shares three with carbon in a triple bond and retains two as a lone pair.
Each atom aims to achieve stability by fulfilling the octet rule, which signifies having eight valence electrons. Although hydrogen is an exception with just two electrons needed, the octet rule generally guides bond formation, ensuring that atoms achieve a stable electronic arrangement.

Chemical Bonds

Chemical bonds are forces that hold atoms together in molecules, resulting from the sharing or transfer of valence electrons. In essence, bonds form when atoms share electrons to achieve a more stable electron configuration.
In \((\text{CH}_3)_2 \text{S}\), chemical bonds are formed between sulfur and carbon atoms. Each carbon shares its valence electrons with hydrogen to form \(\text{CH}_3\). In \(\text{HCN}\), the chemical bond between carbon and nitrogen is a triple bond, which is a robust bond type, involving the sharing of three pairs of electrons.
The \(\text{H}_2\text{C}_2\) molecule features a strong triple bond between two carbon atoms. This multiple bonding significantly affects the chemical and physical properties of the compound. In \(\text{CH}_3\text{F}\), the chemical bond between carbon and fluorine results from sharing electrons. Although each type of bond varies in strength and length, the essence of all bonds is to attain a stable state for the atoms involved, often resembling the noble gas configuration.