Covalent Bonding and Electronegativity

Covalent Bonds

  • Covalent bond defined:

    • Electrostatic attraction between a pair of electrons and two nuclei.
    • Electrons are more centered between the two elements.

  • Potential energy and distance between atoms:

    • As atoms get closer, potential energy decreases.
    • If atoms get too close, potential energy increases due to repulsive forces.

  • Forces involved:

    • Attractive forces: Van der Waals forces or London dispersion forces.
      • Nucleus attracts electrons of the other atom.
    • Repulsive forces:
      • Electrons repel each other.
      • Nuclei repel each other.

  • Net force and bond formation:

    • A sweet spot exists where net attractive force is balanced, leading to bond formation.
    • Bonds vibrate at a frequency, pulling atoms close and pushing them away.
    • A bond represents a low-energy state where attraction and repulsion are balanced.

  • Inert gases

    • Inert gases like Argon (Ar) typically do not form bonds due to their electron configuration and the balance of forces.

Bond Types: Single, Double, and Triple

  • Single, double, and triple bonds:

    • Single bonds are the longest and weakest.
    • Double bonds are shorter and stronger than single bonds.
    • Triple bonds are the shortest and strongest.

  • Data for carbon bonds:

    • Carbon-carbon single bond length is longer than a double or triple bond length.
    • Bond length is measured in nanometers or picometers.

  • Bond strength:

    • Measured in kilojoules per mole (\frac{kJ}{mol}).

Ionic vs. Covalent Bonds

  • Ionic Bond

    • Complete transfer of electrons (although, this is an oversimplification).
    • Typically occurs between a metal and a nonmetal.

  • Covalent Bond

    • Sharing of electrons between two atoms.

  • Polar Covalent Bond

    • Unequal sharing of electrons.
    • Electrons are pulled towards the more electronegative atom.

  • Delta Charges

    • Represent partial charges (positive (\delta^+) or negative (\delta^-)) due to unequal electron sharing.
    • The side with higher electron density has a partial negative charge.

Electronegativity and Bond Character

  • Electronegativity Difference

    • Used to determine whether a bond is ionic, polar covalent, or nonpolar covalent.

  • Sliding Scale

    • Bond character exists on a sliding scale from ionic to covalent.
    • Even in ionic compounds, there can be some covalent character.

  • Electronegativity Difference and Bond Character

    • The greater the difference, the more ionic the bond.
    • No difference results in a pure covalent bond.

  • Pure Covalent Bonds

    • Occur in diatomic molecules (e.g., H2, F2, O_2) where the electronegativity is the same for both atoms.

  • Polar Bonds

    • Arise from differences in electronegativity, resulting in bond dipoles.
    • Bond dipole: has magnitude and direction, pointing towards the negative atom.

Electronegativity Scales and Data

  • Pauling Scale

    • A common scale for electronegativity.

  • Ranges for bond type classification:

    • Pure covalent: 0 difference in electronegativity.
    • Nonpolar covalent: 0 to 0.4 difference.
    • Ionic: Greater than 1.8 difference (arbitrary cutoff).

Bonding Triangle

  • Bonding Triangle:

    • A more comprehensive model considers both the difference and the average of electronegativity values.

  • Scales:

    • One scale: covalent to ionic character.
    • The other scale: average electronegativity.

  • Percentage of Bond Character

    • The triangle can provide a percentage of bond character, indicating the degree of ionic or covalent nature.