Bond Polarity

Bond Polarity

Polar and Nonpolar Molecules

Electronegativity

  • Definition: Electronegativity is the degree of attraction of electrons from another atom in a chemical bond.

  • Characteristics of Nonmetals:

    • High electronegative values

    • Strong attraction for bonding electrons

  • Characteristics of Metals:

    • Low electronegative values

    • Low attraction for bonding electrons

Review: What are Covalent Bonds?

  • Definition: A covalent bond is formed when nonmetal atoms share valence electrons to achieve a full electron shell.

  • Types of Covalent Bonds:

    1. Polar covalent

    2. Nonpolar covalent

What is a Nonpolar Molecule?

  • Definition: Nonpolar molecules formed when nonmetal atoms equally share valence electrons.

  • Characteristics:

    • The atoms in nonpolar molecules are not charged.

    • Typically have similar electronegativity values.

    • Electronegativity Difference (EV): Between 0 to 0.49.

  • Examples:

    • H-H: 2.1 - 2.1 = 0

    • C-H: 2.5 - 2.1 = 0.4

What is a Polar Covalent Bond?

  • Definition: A polar covalent bond exists between nonmetal atoms that unequally share electrons.

  • Key Feature: One atom attracts/pulls the bonded electrons to itself, creating a bond dipole.

  • Electronegativity Difference (EV): A range from 0.5 to 1.9.

  • Example: H-Cl is a polar covalent bond.

    • Calculation: Cl-H: 3.16 - 2.10 = 1.06

Characteristics of the H-Cl Polar Covalent Bond

  • Chlorine's Role:

    • Chlorine attracts the shared electrons to itself due to having a higher electronegativity value.

  • Bond Dipole:

    • Represented by an arrow pointing towards the more electronegative atom (chlorine).

    • Electronegativity Values: Cl = 3.16, H = 2.10

Polar Covalent Molecules and Charge Distribution

  • Partial Charges:

    • Exist because the bonded electrons are pulled toward one atom, creating regions of partial positive and negative charges.

    • Chlorine attains a slight negative charge (δ-)

    • Hydrogen attains a slight positive charge (δ+)

Visualization of Charges in Polar Bonds

  • Charge Representation:

    • Partially positive charge represented as δ+

    • Partially negative charge represented as δ-

  • Regions Characteristics:

    • H: Electron-poor region, slight positive charge

    • F: Electron-rich region, slight negative charge is not explicitly mentioned but assumed in the same context

Molecular Dipole

  • Bond Dipole (Black Arrows):

    • Indicates the direction the bonded electrons are being pulled/attracted between two atoms.

  • Molecular Dipole (Blue Arrow):

    • Represents the overall direction of all the molecule’s bonded electrons being attracted.

Electronegativity Difference and Bonding Types

  • Electronegativity Difference (EV Diff.) and Bond Types:

    • 0 to 0.49: Nonpolar covalent

    • Charges on Atoms: No Charge

    • 0.5 to 1.9: Polar covalent

    • Charges on Atoms: Partial charges (δ+/δ-)

    • 2.0 to 3.3: Ionic Bond

    • Charges on Atoms: Full charges (+/-)

    • Note: In ionic bonds, electrons are completely transferred, resulting in fully charged ions.

Summary Chart for Bonding Types

  • Types of Bonds:

    • Covalent: No charge on atoms, electrons shared equally, electronegativity range (0 - 0.4 ΔEN)

    • Polar Covalent: Partial charge on atoms, electrons shared unequally, electronegativity range (0.4 – 2.0 ΔEN)

    • Ionic: Full charge on ions, electrons transferred, electronegativity greater than 2.0

  • Example Calculation: K+ as a full positive ion demonstrates transfer of electrons in ionic bonding.

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