Select The Polar Bonds In The Compounds Below

How to Identify Polar Bonds in Chemical Compounds: A Step-by-Step Guide

Understanding which bonds within a molecule are polar is a fundamental skill in chemistry that unlocks insights into a substance’s physical properties, reactivity, and behavior in biological systems. A polar bond arises from an unequal sharing of electrons between two atoms, creating a slight electrical imbalance. This guide provides a clear, systematic method to select polar bonds in any compound, moving from basic principles to practical application with common examples.

The Core Principle: Electronegativity

The sole determinant of bond polarity is the difference in electronegativity between the two bonded atoms. Electronegativity is a measure of an atom’s ability to attract and hold onto shared electrons in a covalent bond. The greater the difference, the more polar the bond.

• Nonpolar Covalent Bond: Electronegativity difference ≈ 0.0 to 0.4. Electrons are shared almost equally. (e.g., H-H, C-H, Cl-Cl).
• Polar Covalent Bond: Electronegativity difference ≈ 0.5 to 1.7. Electrons are shared unequally, spending more time near the more electronegative atom. (e.g., H-Cl, C-O, N-H).
• Ionic Bond: Electronegativity difference > ~1.7. Electrons are essentially transferred, forming ions. (e.g., Na-Cl). While ionic compounds have complete charge separation, the bond itself is often discussed separately from covalent polar bonds.

Key Takeaway: You must compare the electronegativity values of the two atoms in each specific bond within a molecule. A molecule can contain both polar and nonpolar bonds.

A Step-by-Step Method to Select Polar Bonds

Follow this checklist for any compound:

Step 1: Draw the Correct Lewis Structure

Before analyzing bonds, you must know the exact connectivity of atoms. A correct Lewis structure shows all atoms, bonds (single, double, triple), and lone pairs. Misplaced atoms will lead to incorrect bond identification.

Step 2: List All Bonds in the Molecule

Identify every covalent bond. Do not consider ionic interactions between separate molecules or ions at this stage. For example, in sodium sulfate (Na₂SO₄), you only analyze the bonds within the sulfate ion (S-O bonds), not the ionic attraction between Na⁺ and SO₄²⁻.

Step 3: Consult the Electronegativity Scale

Use a standard Pauling electronegativity chart. Key values to memorize:

• Fluorine (F): 4.0 (most electronegative)
• Oxygen (O): 3.5
• Nitrogen (N): 3.0
• Chlorine (Cl): 3.0
• Carbon (C): 2.5
• Hydrogen (H): 2.1
• Sodium (Na): 0.9
• Metals generally have low electronegativity (<1.8).

Step 4: Calculate the Difference for Each Bond

For each bond (A-B), subtract the smaller electronegativity from the larger. ΔEN = |EN_A - EN_B|.

Step 5: Apply the Polarity Threshold

If 0.5 ≤ ΔEN ≤ 1.7, the bond is polar covalent. The atom with the higher electronegativity carries a partial negative charge (δ⁻), and the other carries a partial positive charge (δ⁺). This creates a bond dipole—a vector pointing from δ⁺ to δ⁻.

If ΔEN < 0.5, the bond is nonpolar covalent.

Step 6: Consider Symmetry for the Whole Molecule (Important Distinction!)

This step is crucial to avoid a common mistake. Bond polarity and molecular polarity are different concepts. You are only asked to select the polar bonds, not determine if the entire molecule is polar.

• A molecule can have polar bonds but be nonpolar overall if its geometry is symmetrical, causing bond dipoles to cancel out (e.g., CO₂, CCl₄).
• A molecule with polar bonds can be polar if its geometry is asymmetrical, preventing dipole cancellation (e.g., H₂O, NH₃). For this task, you stop at Step 5 for each individual bond. Symmetry only matters if the question asks about the molecule’s polarity.

Practical Application: Analyzing Common Compounds

Let’s apply the method.

Example 1: Water (H₂O)

1. Lewis Structure: H-O-H, with two lone pairs on oxygen.
2. Bonds: Two O-H bonds.
3. Electronegativity: O (3.5), H (2.1). ΔEN = 1.4.
4. Verdict: Both O-H bonds are polar covalent. Oxygen is δ⁻, Hydrogen is δ⁺.

Example 2: Carbon Dioxide (CO₂)

1. Lewis Structure: O=C=O (linear).
2. Bonds: Two C=O double bonds. (Each double bond counts as one bond unit for polarity analysis).
3. Electronegativity: C (2.5), O (3.5). ΔEN = 1.0.
4. Verdict: Both C=O bonds are polar covalent. Oxygen is δ⁻, Carbon is δ⁺. (Note: The molecule is nonpolar due to linear symmetry canceling the two bond dipoles, but the bonds themselves are polar.)

Example 3: Methane (CH₄)

1. Lewis Structure: Tetrahedral C with four H atoms.
2. Bonds: Four C-H bonds.
3. Electronegativity: C (2.5), H (2.1). ΔEN = 0.4.
4. Verdict: All C-H bonds are nonpolar covalent (difference is exactly at the borderline, generally considered nonpolar). The molecule is nonpolar.

Example 4: Ammonia (NH₃)

1. Lewis Structure: Trigonal pyramidal N with three H atoms, one lone