Electronegativity and Molecular Polarity

Electronegativity and Valence Electrons

  • Valence electrons, located in the outermost energy level of an atom, determine an element's electronegativity.
  • Electronegativity is related to the distance between valence electrons and the positively charged nucleus.

Trends in Electronegativity

  • Within a Group: Electronegativity decreases from top to bottom.
    • Atoms at the top are smaller, placing valence electrons closer to the nucleus, leading to greater attraction.
    • Atoms at the bottom are larger, increasing the distance between valence electrons and the nucleus, resulting in weaker attraction.
  • Within a Period: Electronegativity increases from left to right due to the increasing number of protons in the nucleus.
    • Example: Fluorine > Oxygen > Nitrogen in electronegativity.

Covalent Bonds and Electronegativity Difference

  • The difference in electronegativity between two atoms forming a covalent bond determines if the bond is polar or nonpolar.
  • Example 1: HF (Hydrogen Fluoride)
    • Hydrogen and fluorine share electrons, forming a covalent bond.
    • Fluorine is more electronegative than hydrogen.
    • Fluorine attracts the shared electrons, causing them to spend more time closer to fluorine.
    • Fluorine gains a partial negative charge ($\delta^-$).
    • Hydrogen loses partial control of its electron, gaining a partial positive charge ($\delta^+$).
  • Example 2: H2 (Hydrogen Molecule)
    • Two hydrogen atoms share electrons.
    • Since both atoms have equal electronegativity, the electrons are shared equally.
    • There is no charge separation.

Dipoles and Polar Bonds

  • Dipole: A force running from the positive side of a molecule to the negative side, represented by an arrow.
  • Polar Bond: A bond with a dipole.
  • Nonpolar Bond: A bond without a dipole.

Uneven Electron Distribution

  • Occurs when a high electronegativity element bonds with a low electronegativity element.
  • Leads to the formation of partial charges (dipole).
  • The high electronegativity element gets a partial negative charge.
  • The low electronegativity element gets a partial positive charge.

Representing Dipoles

  • A dipole is represented by an arrow.
  • The tip of the arrow points to the negative charge.
  • The base of the arrow (often with a plus sign) points to the positive charge.

Nonpolar Covalent Bonds

  • Form between two nonmetal atoms with the same or very similar electronegativity values.
  • Example: bond between hydrogen and hydrogen or between fluorine and fluorine.
  • Electrons are evenly distributed.
  • No partial charges or bond dipoles.
  • Also found in carbon-carbon and carbon-hydrogen bonds in hydrocarbons (e.g., propene).
  • Example: Carbon electro negativity value is 2.5 and hydrogen is 2.1.

Bond Continuum

  • Bonds exist on a continuum.
  • Nonpolar Covalent Bond: Equally shared electrons, negligible electronegativity difference.
  • Polar Covalent Bond: Unequally shared electrons, significant electronegativity difference.
  • Ionic Bond: Transfer of electrons, large electronegativity difference.

Determining Partial Charges

  • If the difference of electronegativity is anywhere from 0.5 to 2.0, so the bond will be polar.
  • Identify the more electronegative atom.
  • Place a partial negative charge ($\delta^-$) on it.
  • Place a partial positive charge ($\delta^+$) on the less electronegative atom.
  • Draw Bond Dipole: Arrow from positive to negative charge.

Molecular Polarity

  • Molecules with multiple bonds can be polar or nonpolar overall.
  • Polar molecules have a separation of charge.
  • Nonpolar molecules have even electron distribution.
  • Polarity influences physical and chemical properties.
  • Example: Polarity of a drug molecule affects its ability to cross the blood-brain barrier.

Determining Molecular Polarity

  • Consider the polarity of all bonds and the molecular geometry.
  • Nonpolar Molecules:
    • All bonds within the molecule are nonpolar (e.g., methane, CH_4).
    • All bonds in the molecule are polar and identical; the central atom has no nonbonding electrons.
    • Symmetrical molecules (linear, trigonal planar, tetrahedral) with identical polar bonds (e.g., CO2, BF3, CCl_4).
  • CO_2 has dipoles that cancel each other out, due to symmetry.
  • Polar Molecules:
    • Only one polar bond (e.g., HCl).
    • Nonsymmetrical molecular geometry and more than one polar bond (e.g., bent molecules such as H_2O).

Rules for Polarity

  • If all bonds are nonpolar, the molecule is nonpolar.
  • For molecules with polar bonds, consider geometry:
    • Linear, trigonal planar, and tetrahedral geometries are symmetrical; identical polar bonds will cancel out.

Examples

Symmetrical Shapes: Linear, Trigonal Planar, and Tetrahedral.

  • Symmetrical Shapes: Dipoles cancel each other out in these shapes (if the molecule has identical polar bonds).
  • Molecules with only one polar bond are always polar (e.g. HF).
  • Molecules with non-symmetrical molecular geometry are polar (e.g. Bent shape molecules).
  • Molecules with a symmetrical geometry but non-identical polar bonds are polar (e.g.CH2Cl2).

Polarity Problem

  • H_2: symmetrical molecule with the same electronegativity value, so non-polar.
  • CH_2ClO: a symmetrical molecule but the dipoles are non-identical, so polar.
  • IF: One bond and fluorine has a high, electronegativity, iodine is pretty low, so polar.
  • CBr_4: tetrahedral structure, which means dipoles cancel each other out, so nonpolar.
  • CS_2: Carbon and sulfur has pretty similar electronegativity value, so no dipoles, so nonpolar.