U3 Electronegativity & Polarity

Introduction to Ionic and Covalent Compounds

  • Ionic Compounds

    • Formation involves a metal transferring electrons to a non-metal.

    • Resulting Charges:

    • Non-metal: Gains electrons → negative charge.

    • Metal: Loses electrons → positive charge.

    • Bonding mechanism: Attraction of opposite charges (positive and negative).

Comparison with Covalent Compounds

  • Covalent Compounds

    • Involve sharing of electrons between atoms.

    • Formation of a neutral molecule (no net charge).

    • Atoms bonded by an octet rule (stable electron configuration).

Types of Covalent Bonds

  • Comparison of H2 (Hydrogen gas) and HF (Hydrogen Fluoride):

    • Each hydrogen molecule has one proton, sharing electrons results in no charge (neutral).

    • Fluorine: Atomic number 9 → 9 protons; stronger attractive force due to higher positive charge compared to hydrogen's single proton.

    • Electron Sharing in:

    • H2: Electrons shared equally.

    • HF: Electrons shared unequally; closer to fluorine, creating partial charges.

      • Fluorine → partial negative charge.

      • Hydrogen → partial positive charge.

Covalent vs Polar Covalent Bonds

  • Nonpolar Covalent Bonds:

    • Equal sharing of electrons.

    • Example: H2 (hydrogen gas), Cl2 (chlorine gas).

  • Polar Covalent Bonds:

    • Unequal sharing of electrons.

    • Statistics reflect charge distribution (dipole formed).

    • Measured using delta (Δ) symbol for partial charges:

    • ∆+ = partial positive charge.

    • ∆- = partial negative charge.

Electronegativity

  • Definition: The ability of an atom to attract bonding electrons towards itself in a covalent bond.

  • Characteristics:

    • Influenced by atomic number and size of the nucleus.

    • Fluorine is the most electronegative element (highest in periodic table).

    • Electronegativity scale: Ranges from approximately 0 to 4 (with 4 being the most electronegative).

    • Example values near fluorine:

    • Oxygen: 3.5

    • Nitrogen: 3.0

    • Carbon: 2.5

    • Hydrogen: 2.1

Criteria for Classifying Bonds

  • Nonpolar Covalent Bonds:

    • Electronegativity difference between 0 and 0.4.

    • Example: Bond between two identical atoms (N2, O2).

  • Polar Covalent Bonds:

    • Electronegativity difference between 0.5 and 1.7.

    • Non-metals with varying electronegativities lead to charge separation.

    • Example: Water (H2O) with O and H.

  • Ionic Bonds:

    • Significant electronegativity difference (greater than 1.8).

    • Characterized by complete transfer of electrons, forming distinct ions.

    • Example: NaCl (Sodium Chloride).

Summary of Bonding Types

  • Ionic Bonds: Occur between metals and non-metals.

    • Qualitative Differentiation of Electronegativity: Metal has low electronegativity, non-metal has high.

  • Covalent Bonds:

    • Between non-metals with slight differences in electronegativity:

    • Nonpolar (0.0-0.4) → Equally shared electrons.

    • Polar (0.5-1.7) → Unequally shared electrons.

Impact of Polarity on Properties

  • Molecule Attraction:

    • Polar molecules attract each other; stronger forces than non-polar interactions.

    • Solubility: Polar molecules dissolve in polar solvents (like water); non-polar in non-polar solvents (like oils).

Guidelines for Determining Molecule Polarity

  • Nonpolar Molecules:

    • Even distribution of electrons; symmetrical.

  • Polar Molecules:

    • Uneven distribution; lone pairs of electrons or different central bonds.

Conclusion

  • Understanding bond types (ionic, polar, nonpolar) is crucial for predicting molecular behavior and interactions, particularly in solubility and chemical reactivity.