2.3 Chemical Bonding & Electronegativity

Chemical Bonding and Electronegativity

Introduction to Electronegativity

• Electronegativity: Defined as the ability of an atom to attract bonding electrons in a chemical bond.

• Developed by Linus Pauling in 1922; assigned values to elements influencing their bond characteristics.

• Highest electronegativity: Fluorine (4.0), lowest: Francium (0.7).

• Periodic Trend: Increases from left to right on the periodic table and decreases from top to bottom.

Understanding Electronegativity Trends

• Atomic Radius:

  • As atomic radius increases (going down a group), the ability to attract bonding electrons decreases, hence electronegativity decreases.

• Electronegativity is calculated from physical properties like ionization energy; it cannot be measured directly.

• Common electronegativity values (Table 1):

  • H: 2.2, C: 2.6, N: 3.0, O: 3.4, F: 4.0, Na: 0.9, Mg: 1.3, Cl: 3.2, Ca: 1.0.

Classifying Compounds: Ionic vs Molecular

General Rules for Bond Formation

• Ionic Compounds: Formed by the combination of metals and non-metals.

• Molecular Compounds: Formed by the combination of non-metals.

Role of Electronegativity in Bond Classification

• Electronegativity Difference (ΔEN): Used to predict bond nature. Calculated as the absolute difference between two atoms’ electronegativities.

  • Ionic Bond: ΔEN ≥ 1.7 (greater likelihood of electron transfer).

  • Covalent Bond: ΔEN < 1.7 (electron sharing).

Examples of Electronegativity Differences

• Example: Sodium (Na) and Chlorine (Cl):

  • Na: 0.9, Cl: 3.2, thus ΔEN = 3.2 - 0.9 = 2.3 (ionic bond).

• Identical atoms (e.g., H-H): ΔEN = 0 (non-polar covalent bond as electrons are shared equally).

Types of Covalent Bonds

Non-Polar Covalent Bonds

• Occur when ΔEN = 0 (identical atoms, e.g., H2 or Cl2).

• Electrons are shared equally between atoms.

• Characterized by simple equal sharing of bonding electrons.

Polar Covalent Bonds

• Occur when there is a significant electronegativity difference (0 < ΔEN < 1.7).

• Example: H and Cl (ΔEN = 0.9).

  • Cl has higher electronegativity, pulls electrons closer, resulting in localized positive (H) and negative (Cl) poles.

• Represented with delta symbols (δ+ for positive and δ- for negative charges).

Summary of Key Concepts

• Electronegativity indicates an element’s ability to attract bonding electrons.

• The electronegativity difference aids in classifying bonds:

  • Ionic (ΔEN ≥ 1.7)

  • Covalent (ΔEN < 1.7)

    • Non-Polar Covalent: ΔEN = 0

    • Polar Covalent: 0 < ΔEN < 1.7

Review Questions

1. Discuss electronegativity trends in relation to atomic radius.

2. Arrange elements (K, Cs, Br, Fe, Ca, F, Cl) in increasing order of electronegativity.

3. Explain how electronegativity values predict bonding types.

4. Differentiate between polar and non-polar covalent bonds.

5. Given various bonds, compute the electronegativity difference and determine the bond type.

6. Identify the more polar bond in provided pairs.

7. Predict bond types in listed compounds.

8. Discuss the most ionic bond and its properties based on electronegativities.

9. Summarize Linus Pauling’s achievements and contributions to science.

10. Research Neil Bartlett’s discovery regarding xenon and fluorine compounds.