Ch. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 1)

Bruice8th EditionOrganic ChemistryISBN: 9780135213711Not the one you use?Change textbook

Bruice 8th Edition

Ch. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 1)

Problem 17d

Chapter 1, Problem 17d

Draw the Lewis structure for each of the following:
d. ⁺C2H5

Verified step by step guidance

Step 1: Determine the total number of valence electrons in the molecule. Carbon (C) has 4 valence electrons, hydrogen (H) has 1 valence electron, and the positive charge (+) indicates the loss of one electron. For C2H5+, calculate the total valence electrons as follows: \( (2 \times 4) + (5 \times 1) - 1 \).

Step 2: Arrange the atoms in the molecule. Carbon is less electronegative than hydrogen, so the two carbons will form the backbone of the structure. Attach the hydrogens to the carbons, ensuring that the structure reflects the molecular formula \( C_2H_5^+ \).

Step 3: Form a single bond between the two carbon atoms. Then, attach the hydrogens to the carbons using single bonds. Distribute the hydrogens such that one carbon has three hydrogens (\( CH_3 \)) and the other carbon has two hydrogens (\( CH_2 \)).

Step 4: Adjust the structure to account for the positive charge. Since the molecule is a cation, one of the carbons will have an incomplete octet. Remove one electron from the carbon atom that is bonded to fewer hydrogens (the \( CH_2 \) group), resulting in a carbocation.

Step 5: Verify the structure by checking that the total number of valence electrons matches the calculated total from Step 1 and that the positive charge is correctly represented. Ensure that all hydrogens have a single bond and that the carbocation is clearly indicated.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They use dots to represent electrons and lines to represent bonds between atoms. Understanding how to draw Lewis structures is essential for visualizing molecular geometry and predicting reactivity.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial in determining how atoms bond with each other. The number of valence electrons influences the atom's ability to form bonds, either by sharing electrons (covalent bonds) or transferring them (ionic bonds). Knowing the valence electron count for each atom helps in accurately constructing Lewis structures.

Formal Charge

Formal charge is a concept used to determine the charge distribution within a molecule, calculated based on the number of valence electrons, the number of bonds, and the number of lone pair electrons on an atom. It helps in assessing the stability of a Lewis structure; structures with formal charges closest to zero are generally more stable. Understanding formal charge is vital for evaluating different resonance structures and their contributions to the overall molecule.

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Textbook Question

Give each atom the appropriate formal charge:

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Textbook Question

An atom with a formal charge does not necessarily have more or less electron density than the atoms in the molecule without formal charges. We can see this by examining the potential maps for H₂O, H₃O⁺, and HO⁻.

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c. Which atom bears the formal positive charge in the hydronium ion?

d. Which atom has the least electron density in the hydronium ion?

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Textbook Question

Which of the atoms in the molecular models in [Problem 20] have

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a. three lone pairs?

b. two lone pairs?

c. one lone pair?

d. no lone pairs?