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

There are two nonbonding electron pairs in \((\mathrm{CH}_{3})_{2} \mathrm{~S}\), one nonbonding electron pair in \(\mathrm{HCN}\), one nonbonding electron pair in \(\mathrm{H}_{2} \mathrm{C}_{2}\), and three nonbonding electron pairs in \(\mathrm{CH}_{3} \mathrm{~F}\).

Step-by-step solution

(a) Find total valence electrons for \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S}\)

First, we need to find the total number of valence electrons in this molecule. The molecule is (CH3)2S, which means there are two CH3 groups attached to a sulfur atom. Each carbon has 4 valence electrons, and each hydrogen has 1 valence electron. Sulfur, being in group 16, has 6 valence electrons. Therefore, the total number of valence electrons for this molecule is: 2(Carbon) + 6(Hydrogen) + 1(Sulfur) = 2(4) + 6(1) + 6 = 20 electrons.

(a) Calculate nonbonding electron pairs for \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S}\)

Next, we need to find how many electrons are involved in bonding. Each C-H bond requires 2 electrons, and there are 6 C-H bonds in total. The Sulfur atom forms two bonds with the two carbons, with each bond requiring 2 electrons. So, the total number of bonding electrons is: 6(C-H bonds) + 2(S-C bonds) = 6(2) + 2(2) = 16 electrons. Now, we can determine the nonbonding electron pairs: Total valence electrons - Bonding electrons = Nonbonding electrons 20 - 16 = 4 electrons Thus, there are two nonbonding electron pairs for \((\mathrm{CH}_{3})_{2} \mathrm{~S}\), as each pair has 2 electrons.

(b) Find total valence electrons for \(\mathrm{HCN}\)

For HCN, the total number of valence electrons is: 1(Hydrogen) + 4(Carbon) + 5(Nitrogen) = 1(1) + 4(1) + 5(1) = 10 electrons.

(b) Calculate nonbonding electron pairs for \(\mathrm{HCN}\)

The bonding in HCN involves a single H-C bond and a triple C≡N bond. Therefore, there are: 1(H-C bond) + 3(C≡N bonds) = 1(2) + 3(2) = 8 bonding electrons. Nonbonding electron pairs can be calculated as follows: Total valence electrons - Bonding electrons = Nonbonding electrons 10 - 8 = 2 electrons Thus, there is one nonbonding electron pair for HCN.

(c) Find total valence electrons for \(\mathrm{H}_{2} \mathrm{C}_{2}\)

For H2C2, the total valence electrons are: 2(Hydrogen) + 2(Carbon) = 2(1) + 2(4) = 10 electrons.

(c) Calculate nonbonding electron pairs for \(\mathrm{H}_{2} \mathrm{C}_{2}\)

The bonding in H2C2 involves two H-C single bonds and a C≡C triple bond. There are: 2(H-C bonds) + 3(C≡C bonds) = 2(2) + 3(2) = 10 bonding electrons. However, there are 4 shared electrons in a triple bond: Total valence electrons - Bonding electrons + shared electrons = Nonbonding electrons 10 - 10 + 2 = 2 electrons Thus, there is one nonbonding electron pair for H2C2.

(d) Find total valence electrons for \(\mathrm{CH}_{3} \mathrm{~F}\)

For CH3F, the total valence electrons are: 3(Hydrogen) + 1(Carbon) + 7(Fluorine) = 3(1) + 1(4) + 1(7) = 14 electrons.

(d) Calculate nonbonding electron pairs for \(\mathrm{CH}_{3} \mathrm{~F}\)

The bonding in CH3F involves three H-C single bonds and a C-F single bond. There are: 3(H-C bonds) + 1(C-F bond) = 3(2) + 1(2) = 8 bonding electrons. Nonbonding electron pairs can be calculated as follows: Total valence electrons - Bonding electrons = Nonbonding electrons 14 - 8 = 6 electrons Thus, there are three nonbonding electron pairs for CH3F.

Key concepts

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. They are crucial in determining how an atom interacts with others, especially in the formation of chemical bonds. Typically, elements aim to achieve a stable electron configuration similar to the nearest noble gas, which often involves having eight electrons in their valence shell—a concept known as the octet rule.

When counting valence electrons for a molecule, the total is found by adding the valence electrons of each atom present in the molecule. For instance, in the molecule \((\mathrm{CH}_3)_2 \mathrm{~S}\), each carbon atom contributes 4 electrons, each hydrogen contributes 1 electron, and sulfur brings in 6 electrons. Adding them gives the total valence electrons for the molecule. Understanding the distribution of these electrons is essential for drawing correct molecular structures and predicting the chemical behavior of compounds.

Chemical Bonding

Chemical bonding relates to the force that holds atoms together in a molecule. There are several types of chemical bonding, including covalent, ionic, and metallic bonds. Covalent bonds, which involve the sharing of electrons between atoms, are most frequently encountered in organic compounds.

In the context of our molecules from the exercise, such as \(\mathrm{HCN}\) and \(\mathrm{CH}_3 \mathrm{~F}\), covalent bonds are predominantly present. The number of shared electrons determines whether the bond is a single, double, or triple bond. For example, in \(\mathrm{HCN}\), we have a triple bond between carbon and nitrogen indicating three pairs of electrons being shared. Understanding these bonds helps in knowing the molecule's stability, orientation, and reactivity in different environments.

Lewis Structures

Lewis structures are graphical representations that show how valence electrons are arranged among atoms in a molecule. They provide visual insights into the bonding between atoms and the presence of any nonbonding electron pairs, often depicted as dots surrounding the atomic symbols.

For instance, in \(\mathrm{CH}_3\mathrm{~F}\), the carbon atom forms bonds with three hydrogens and one fluorine. The fluorine atom, with seven valence electrons, forms a bond using one of its electrons to pair with carbon's electron. The remaining six electrons appear as nonbonding pairs, visible as three pairs around fluorine in the Lewis structure.

• Nonbonding pairs, also known as lone pairs, can affect the geometric shape of a molecule and its polarity.
• Understanding Lewis structures aids in predicting molecule behavior during chemical reactions.

Creating accurate Lewis structures is a fundamental skill in chemistry, as they aid greatly in understanding molecular stability and reactivity.