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

Verified step by step guidance

Step 1: Understand the concept of bond polarity. Bond polarity arises due to the difference in electronegativity between two atoms in a covalent 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: Identify the electronegativity values of the atoms involved in the bonds. For the H3C—NH2 bond, compare the electronegativity of carbon (C, approximately 2.55) and nitrogen (N, approximately 3.04). For the H3C—Cl bond, compare the electronegativity of carbon (C, approximately 2.55) and chlorine (Cl, approximately 3.16).

Step 3: Determine the direction of polarity for the H3C—NH2 bond. Since nitrogen (N) is more electronegative than carbon (C), the shared electrons will be pulled toward the nitrogen atom. Assign δ− to nitrogen and δ+ to carbon.

Step 4: Determine the direction of polarity for the H3C—Cl bond. Since chlorine (Cl) is more electronegative than carbon (C), the shared electrons will be pulled toward the chlorine atom. Assign δ− to chlorine and δ+ to carbon.

Step 5: Represent the polarity of each bond using the δ+ and δ− symbols. For H3C—NH2, write Cδ+—Nδ−. For H3C—Cl, write Cδ+—Clδ−. This notation indicates the direction of electron density shift in each bond.

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Key Concepts

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

Polarity of Bonds

Polarity in chemical bonds arises from the difference in electronegativity between the atoms involved. When two atoms form a bond, the atom with higher electronegativity attracts the shared electrons more strongly, resulting in a partial negative charge (δ−) on that atom and a partial positive charge (δ+) on the other. This creates a dipole moment, indicating the direction of electron density.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond. It varies across the periodic table, with elements like fluorine being highly electronegative. Understanding the relative electronegativities of the atoms in a bond helps predict the bond's polarity and the resulting charge distribution.

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 the context of polar bonds, the dipole moment indicates the direction of polarity, pointing from the positive end (δ+) to the negative end (δ−), which is crucial for understanding molecular interactions.

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