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Ch. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 2)
Bruice - Organic Chemistry 8th Edition
Bruice8th EditionOrganic ChemistryISBN: 9780135213711Not the one you use?Change textbook
All textbooksBruice 8th EditionCh. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 2)Problem 11g,h
Chapter 2, Problem 11g,h

Use the symbols δ+ and δ− to show the direction of the polarity of the indicated bond in each of the following compounds:
g. I—Cl
h. H2N—OH

Verified step by step guidance
1
Step 1: Understand the concept of bond polarity. Bond polarity arises due to the difference in electronegativity between two atoms in a bond. The atom with higher electronegativity attracts the shared electrons more strongly, acquiring a partial negative charge (δ−), while the less electronegative atom acquires a partial positive charge (δ+).
Step 2: For the bond I—Cl, compare the electronegativities of iodine (I) and chlorine (Cl). Chlorine is more electronegative than iodine. Therefore, chlorine will have a δ− charge, and iodine will have a δ+ charge. Represent this as: I−Cl, where I has δ+ and Cl has δ−.
Step 3: For the bond H2N—OH, analyze the polarity of the N—O bond. Oxygen (O) is more electronegative than nitrogen (N), so oxygen will have a δ− charge, and nitrogen will have a δ+ charge. Represent this as: H2N−OH, where N has δ+ and O has δ−.
Step 4: Additionally, consider the polarity of the O—H bond in H2N—OH. Oxygen is more electronegative than hydrogen, so oxygen will have a δ− charge, and hydrogen will have a δ+ charge. Represent this as: O−H, where O has δ− and H has δ+.
Step 5: Summarize the polarity directions for each bond in the compounds. For I—Cl, the bond is polarized with δ+ on iodine and δ− on chlorine. For H2N—OH, the N—O bond is polarized with δ+ on nitrogen and δ− on oxygen, and the O—H bond is polarized with δ+ on hydrogen and δ− on oxygen.

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Key Concepts

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

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond. In a bond between two different atoms, the atom with higher electronegativity will attract the shared electrons more strongly, resulting in a polar bond. This difference in electronegativity is crucial for determining the polarity of bonds, as it dictates the distribution of electron density.
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Polarity of Bonds

Polarity in chemical bonds arises when there is an unequal sharing of electrons between two atoms due to differences in their electronegativities. A polar bond is characterized by a dipole moment, where one end of the bond becomes partially negative (δ−) and the other end becomes partially positive (δ+). Understanding bond polarity is essential for predicting molecular behavior, reactivity, and interactions.
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Dipole Moment

A dipole moment is a vector quantity that represents the separation of positive and negative charges in a molecule. It is calculated as the product of the charge and the distance between the charges. In polar bonds, the dipole moment indicates the direction of electron density, with the δ− end pointing towards the more electronegative atom and the δ+ end towards the less electronegative atom, providing insight into molecular polarity.
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Related Practice
Textbook Question

Use the symbols δ+ and δ− to show the direction of the polarity of the indicated bond in each of the following compounds:

e. HO—Br

f. H3C—Li

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Textbook Question

Which of the following has

a. the most polar bond?

b. the least polar bond?

NaI LiBr Cl2 KCl

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Textbook Question

Use the symbols δ+ and δ− to show the direction of the polarity of the indicated bond in each of the following compounds:

c. H3C—NH2

d. H3C—Cl

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Textbook Question

After examining the potential maps for LiH, HF, and H2, answer the following questions:

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a. Which compounds are polar?

b. Why does LiH have the largest hydrogen?

c. Which compound has the hydrogen that would be most apt to attract a negatively charged molecule?

Textbook Question

Explain why HCl has a smaller dipole moment than HF, even though the H—Cl bond is longer than the H—F bond.

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Textbook Question

Draw the Lewis structure for each of the following:

f. NaOH

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