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Chapter 1: Problem 8

Each of the following compounds has at least one multiple bond. Draw a Lewis structure for each molecule. Use lines to indicate electrons in bonds and dots to indicate nonbonding electrons. (a) \(\mathrm{F}_{2} \mathrm{CCF}_{2}\) *(b) \(\mathrm{H}_{3} \mathrm{CCN}\) (c) \(\mathrm{H}_{2} \mathrm{CO}\) (d) \(\mathrm{H}_{2} \mathrm{CCO}\) (e) \(\mathrm{H}_{2} \mathrm{CCHCHCH}_{2}\) (f) \(\mathrm{H}_{3} \mathrm{CNO}\) (g) \(\mathrm{H}_{3} \mathrm{COCOH}\)

Short Answer

Expert verified
Draw Lewis structures ensuring that all atoms satisfy the octet rule by counting valence electrons and adding necessary multiple bonds.

Step by step solution

01

- Understanding Multiple Bonds

Identify the types of bonds (single, double, or triple) between atoms in each molecule. A multiple bond is any bond that is a double or triple bond.
02

- Counting Valence Electrons

Sum the valence electrons for all atoms in each molecule. For instance, carbon has 4 valence electrons, hydrogen has 1 valence electron, etc.
03

- Drawing the Skeleton Structure

Draw the skeleton structure of each molecule, using single bonds to connect atoms initially.
04

- Adding Multiple Bonds

Identify where double or triple bonds are needed in the skeleton structure to satisfy the octet rule for each atom.
05

- Placing Nonbonding Electrons

Distribute remaining valence electrons as dots around atoms to complete their octets, starting with the most electronegative atoms.
06

- Example for \(\mathrm{F}_{2}\mathrm{CCF}_{2}\)

1. Valence Electrons: 28 (2F x 7, 2C x 4)2. Skeleton: \[ F-C=C-F \] 3. Add Double Bond: \[ F-C=C-F \] 4. Complete Electrons: Each C has 8 electrons around it and each F has 8 electrons. Final:\[{F}::{C}={C}::{F} \]
07

- Example for \(\mathrm{H}_{3}\mathrm{CCN}\)

1. Valence Electrons: 14 (3H x 1, 2C x 4, 1N x 5)2. Skeleton: \[ H_{3}C-C=N \] 3. Add Triple Bond between C and N: \[ H_{3}C-C\equiv N \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{3}C-C\equiv N \]
08

- Example for \(\mathrm{H}_{2}\mathrm{CO}\)

1. Valence Electrons: 12 (2H x 1, 1C x 4, 1O x 6)2. Skeleton: \[ H_{2}C-O \] 3. Add Double Bond between C and O: \[ H_{2}C=O \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{2}C=O \]
09

- Example for \(\mathrm{H}_{2}\mathrm{CCO}\)

1. Valence Electrons: 16 (2H x 1, 2C x 4, 1O x 6)2. Skeleton: \[ H_{2}C-C-O \] 3. Add Double Bond between C and O: \[ H_{2}C-C=O \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{2}C-C=O \]
10

- Example for \(\mathrm{H}_{2}\mathrm{CCHCHCH_{2}}\)

1. Valence Electrons: 24 (4H x 1, 4C x 4)2. Skeleton: \[ H_{2}C-CH-CH-CH_{2} \] 3. Add Double Bonds: \[ H_{2}C=CH-CH=CH_{2} \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{2}C=CH-CH=CH_{2} \]
11

- Example for \(\mathrm{H}_{3}\mathrm{CNO}\)

1. Valence Electrons: 18 (3H x 1, 1C x 4, 1N x 5, 1O x 6)2. Skeleton: \[ H_{3}C-N-O \] 3. Add Double Bond between N and O: \[ H_{3}C-N=O \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{3}C-N=O \]
12

- Example for \(\mathrm{H}_{3}\mathrm{COCOH}\)

1. Valence Electrons: 20 (4H x 1, 2C x 4, 2O x 6)2. Skeleton: \[ H_{3}C-O-C-OH \] 3. Add Double Bond between C and O: \[ H_{3}C-O-C=O-H \] 4. Complete Electrons: Each atom has its octet complete. Final:\[H_{3}C-O-C=O-H \]

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Multiple Bonds
Multiple bonds are essential in understanding Lewis structures and molecular geometry. When two atoms share more than one pair of electrons, they form multiple bonds. These can either be double bonds, involving two pairs of shared electrons, or triple bonds, involving three pairs of shared electrons. Multiple bonds are stronger than single bonds and significantly affect the molecule's properties and shape. For example, in \(C_2H_2\), the triple bond between the two carbon atoms contributes to the molecule's linear shape and strong bond strength.
Valence Electrons
Valence electrons are the outermost electrons of an atom and are crucial in forming bonds with other atoms. For example, carbon has four valence electrons, while oxygen has six. To determine the total number of valence electrons in a molecule, sum up the valence electrons of each atom. This sum helps in creating the initial skeleton structure and in fulfilling the octet rule during the Lewis structure formation.
Octet Rule
The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, resembling a noble gas configuration. For example, in the molecule \(H_2CO\), the double bond between carbon and oxygen ensures that both atoms follow the octet rule. Hydrogen is an exception as it only needs two electrons. Ensuring all atoms (except hydrogen) reach an octet is critical for drawing correct Lewis structures.
Skeleton Structure
A skeleton structure is the initial arrangement of atoms in a molecule connected by single bonds. It serves as the framework for adding multiple bonds and nonbonding electrons. For instance, in \(H_3CCN\), the initial skeleton would be \(H_3C-C-N\). This step simplifies determining where additional bonds or electron pairs are needed. The skeleton structure is always the starting point in drawing a complete Lewis structure.
Nonbonding Electrons
Nonbonding electrons, also known as lone pairs, are valence electrons not involved in bonding. They are depicted as dots around the atoms in Lewis structures. Nonbonding electrons play a significant role in the molecule's properties and reactivity. In \(H_2CO\), the remaining valence electrons after satisfying the bonds are placed around oxygen as lone pairs. This allocation ensures each element adheres to the octet rule, where applicable, enhancing molecule stability.

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Most popular questions from this chapter

Write the electronic configurations for the atoms in the fourth row of the periodic table, \({ }_{19} \mathrm{~K}\) through \({ }_{36} \mathrm{Kr}\).

Add electrons to complete the following Lewis structures. In each case, the charge is placed as close as possible to the charged atom. (a) \({ }^{+} \mathrm{CH}_{2}\) (b) \(-\mathrm{CH}_{2} \mathrm{CH}_{3}\) (c) \(\mathrm{HC}=\mathrm{CH}_{2}\) (d) \({ }^{+} \mathrm{OH}_{3}\) (e) \({ }^{-} \mathrm{OH}\) (f) \({ }^{+} \mathrm{NH}_{2}\) (g) \(-\mathrm{NH}_{2}\) (h) \(\mathrm{CH}_{3}-\mathrm{C} \equiv \mathrm{N}-\mathrm{H}\)

Write the electronic configurations for \({ }_{11} \mathrm{Na},{ }_{13} \mathrm{Al},{ }_{15} \mathrm{P},{ }_{16} \mathrm{~S}\), and \({ }_{18} \mathrm{Ar}\).

Construct Lewis structures for the following neutral molecules: *(a) \(\mathrm{BF}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{Be}\) (c) \(\mathrm{SiH}_{4}\) (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (e) \(\mathrm{HOCH}_{3}\) (f) \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{NH}_{2}\)

Draw a Lewis structure for nitric acid \(\left(\mathrm{HO}-\mathrm{NO}_{2}\right)\), and verify that the nitrogen is positive and one of the oxygens is negative
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