Question

How many nonbonding electron pairs are there in each of the following molecules: (a) \(\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3},(\mathbf{b}) \mathrm{CO},(\mathbf{c}) \mathrm{BF}_{3},\) (d) \(\mathrm{SO}_{2} ?\)

Short answer

In the given molecules, the number of nonbonding electron pairs are as follows: (a) N(CH3)3: 1 nonbonding electron pair, (b) CO: 2 nonbonding electron pairs, (c) BF3: 9 nonbonding electron pairs, and (d) SO2: 5 nonbonding electron pairs.

Step-by-step solution

Determine the total number of valence electrons for each molecule

To do this, we need to know the number of valence electrons for each element and then find the total number of valence electrons in each molecule. - Nitrogen (N) has 5 valence electrons - Carbon (C) has 4 valence electrons - Hydrogen (H) has 1 valence electron - Oxygen (O) has 6 valence electrons - Boron (B) has 3 valence electrons - Fluorine (F) has 7 valence electrons - Sulfur (S) has 6 valence electrons Now, let's find the total number of valence electrons for each given molecule: (a) N(CH3)3 has 1 Nitrogen, 3 Carbons, and 9 Hydrogens Total valence electrons = 5(1) + 4(3) + 1(9) = 5 + 12 + 9 = 26 (b) CO has 1 Carbon and 1 Oxygen Total valence electrons = 4(1) + 6(1) = 4 + 6 = 10 (c) BF3 has 1 Boron and 3 Fluorines Total valence electrons = 3(1) + 7(3) = 3 + 21 = 24 (d) SO2 has 1 Sulfur and 2 Oxygens Total valence electrons = 6(1) + 6(2) = 6 + 12 = 18

Draw the Lewis structure of each molecule

Now that we have the total number of valence electrons, we can draw the Lewis structure for each molecule. (a) N(CH3)3: - Draw a central Nitrogen atom surrounded by three CH3 groups. - Each Carbon atom is bonded to three Hydrogen atoms. - Nitrogen forms three single bonds with the Carbon atoms, leaving it with one lone electron pair. (b) CO: - Draw a central Carbon atom bonded to an Oxygen atom. - They form a triple bond with the remaining two electron pairs on Oxygen as lone pairs. (c) BF3: - Draw a central Boron atom surrounded by three Fluorine atoms. - Boron forms single bonds with each Fluorine atom, and each Fluorine atom has three lone pairs. (d) SO2: - Draw a central Sulfur atom bonded to two Oxygen atoms. - Sulfur forms a double bond with one Oxygen atom and a single bond with the other Oxygen atom. The Sulfur also has one lone pair of electrons, and the Oxygen atoms each have two lone pairs.

Identify the nonbonding electron pairs in each molecule

Count the number of nonbonding electron pairs for each molecule: (a) N(CH3)3: - There is one nonbonding electron pair on the Nitrogen atom, and none on the Carbon or Hydrogen atoms, so there is a total of 1 nonbonding electron pair. (b) CO: - There are two nonbonding electron pairs on the Oxygen atom, and none on the Carbon atom. So, there is a total of 2 nonbonding electron pairs. (c) BF3: - There are no nonbonding electron pairs on the Boron atom, and three nonbonding electron pairs on each Fluorine atom, for a total of 3 x 3 = 9 nonbonding electron pairs. (d) SO2: - There is one nonbonding electron pair on the Sulfur atom, and two nonbonding electron pairs on each Oxygen atom, for a total of 1 + 2(2) = 5 nonbonding electron pairs. In summary, for these molecules, there are: (a) N(CH3)3: 1 nonbonding electron pair (b) CO: 2 nonbonding electron pairs (c) BF3: 9 nonbonding electron pairs (d) SO2: 5 nonbonding electron pairs

Key concepts

Valence Electrons

Let's dive into the world of valence electrons! Valence electrons are the outermost electrons found in an atom. They're crucial because these are the electrons that participate in chemical bonding. Knowing the number of valence electrons helps predict how atoms will bond and what molecules will form. Each element in the periodic table has a characteristic number of valence electrons:

• Nitrogen (N) has 5 valence electrons.
• Carbon (C) has 4 valence electrons.
• Hydrogen (H) has 1 valence electron.
• Oxygen (O) has 6 valence electrons.
• Boron (B) has 3 valence electrons.
• Fluorine (F) has 7 valence electrons.
• Sulfur (S) has 6 valence electrons.

Understanding and counting these electrons is the first step when working on molecular structures like those found in N(CH3)3, CO, BF3, and SO2. It's their arrangements that dictate the chemical properties and structures of molecules.

Lewis Structure

The Lewis structure is a handy model for visualizing valence electrons in molecules and is a great tool for predicting the shape, bonding types, and lone pairs in a molecule. Here's how you create a Lewis structure: 1. Count the total valence electrons available in the molecule. 2. Position the atoms to reflect typical molecular connections with a central atom. 3. Distribute electrons to form bonds, typically starting with single connections. 4. Assign remaining electrons to form lone pairs, ensuring that each atom achieves a full outer shell (where applicable).
For example, in the case of BF3: - Place Boron in the center with Fluorine atoms surrounding it. - Connect each Fluorine with Boron using a single bond. - The remaining electrons become lone pairs, situated on Fluorine atoms. Visualizing molecules in this way allows us to better understand the makeup and function of various compounds.

Molecular Geometry

Molecular geometry is all about the 3D arrangement of atoms in a molecule. It tells you not just what atoms are present, but how they sit in space. This is critical for understanding the behavior and reactivity of molecules. Different atoms in a molecule repel each other due to electron cloud interactions, and this is what determines the molecule's shape. Some common shapes include:

• Linear
• Trigonal planar
• Tetrahedral
• Bent or angular
• Pyramidal

For instance, the geometry of CO is linear due to its two atoms and triple bond, while BF3 forms a trigonal planar shape because of the single bonded arrangement and absence of a lone pair on the central atom. The spatial arrangement provides insight into polarity, reactivity, and interaction with other molecules.

Electron Pair Geometry

The electron pair geometry considers both bonding and nonbonding electron pairs around the central atom of a molecule. Unlike molecular geometry, which only looks at atom positions, electron pair geometry accounts for all electron clouds. This comprehensive view is essential for interpreting molecular shapes and how they change bond angles:

• Tetrahedral: when there are four pairs (like in CH4)
• Trigonal planar: with three electron pairs (as in BF3)
• Linear: made by two electron pairs (found in CO)
• Bent: with more than two atom connections, including lone pairs (e.g., SO2)

The presence of nonbonding pairs affects the geometry by pushing bonded pairs closer together, altering the angles between them. For molecules such as SO2, the lone pairs on sulfur result in a bent shape instead of a linear one, despite its initial electron distribution suggesting a different formation.