1. Draw the Lewis structure.
2. Show how the bond dipole moments (and those of any nonbonding pairs of electrons) contribute to the molecular dipole moment.
3. Estimate whether the compound will have a large, small, or zero dipole moment.
g. HCN
h. CH₃CHO
i. H₂C=NH

Verified step by step guidance

Step 1: Begin by drawing the Lewis structure for each molecule. For HCN, identify the central carbon atom bonded to hydrogen and nitrogen. For CH3CHO, recognize the aldehyde group with a carbonyl bond. For H2C=NH, note the double bond between carbon and nitrogen.

Step 2: Assign formal charges and check for resonance structures if applicable. Ensure that each atom satisfies the octet rule where possible, and adjust for any formal charges.

Step 3: Identify the bond dipole moments. For HCN, consider the electronegativity difference between hydrogen, carbon, and nitrogen. For CH3CHO, focus on the polar C=O bond. For H2C=NH, evaluate the C=N bond polarity.

Step 4: Analyze how these bond dipole moments contribute to the overall molecular dipole moment. Consider the molecular geometry and symmetry to determine if dipoles cancel out or reinforce each other.

Step 5: Estimate the magnitude of the dipole moment. For HCN, expect a significant dipole due to the linear structure and polar bonds. For CH3CHO, anticipate a moderate dipole due to the polar C=O bond. For H2C=NH, predict a small dipole moment due to the geometry and bond polarities.

Verified video answer for a similar problem:

This video solution was recommended by our tutors as helpful for the problem above.

Video duration:

Was this helpful?

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 of a molecule and the lone pairs of electrons that may exist. They are essential for visualizing the arrangement of atoms, the distribution of electrons, and the connectivity within a molecule, which is crucial for predicting molecular geometry and reactivity.

Bond Dipole Moments

Bond dipole moments occur due to differences in electronegativity between bonded atoms, resulting in a separation of charge. This concept is important for understanding how individual bond polarities contribute to the overall molecular dipole moment, which affects the molecule's physical properties and interactions with other molecules.

Molecular Dipole Moment

The molecular dipole moment is a vector sum of all individual bond dipoles and lone pair contributions in a molecule. It determines the molecule's polarity, influencing its solubility, boiling point, and interaction with electric fields. Estimating the dipole moment involves considering both the magnitude and direction of each bond dipole within the molecular geometry.

Recommended video:

Guided course

01:46

How dipole-dipole forces work.

Related Practice

Textbook Question

Two isomers of 1,2-dichloroethene are known. One has a dipole moment of 2.4 D; the other has zero dipole moment. Draw the two isomers, and explain why one has zero dipole moment.

CHCl=CHCl 1,2-dichloroethene

The C=O double bond has a dipole moment of about 2.4 D and a bond length of about 1.23 Å.

b. Use this information to evaluate the relative importance of the following two resonance contributors:

3100

views

Textbook Question

The N—F bond is more polar than the N—H bond, but NF₃ has a smaller dipole moment than NH₃. Explain this curious result.

NF₃