Bond Types and Polarity Notes( AP chem unit 2)

Electronegativity, Bond Types, and Polarity

• Electronegativity (EN) differences dictate bond type and polarity. This is viewed as a continuum rather than sharp cut-offs, with approximate thresholds.

Key Concepts

• Electronegativity (EN): The ability of an atom's nucleus to attract shared electrons towards itself.
• Bond Types are determined by EN differences:
  • Ionic Bond: Large EN difference (\Delta\text{EN} > 1.7); electron transfer (metal to nonmetal).
  • Covalent Bond: Smaller EN difference (\Delta\text{EN} < 1.7); electrons are shared.
  • Polar Covalent Bond: Intermediate EN difference (typically 0.4 < \Delta\text{EN} < 1.7); unequal electron sharing.
• EN Trends:
  • Across a period (left to right): EN increases due to more protons (stronger nuclear attraction) with similar energy levels.
  • Down a group: EN decreases due to increased energy levels, placing valence electrons farther from the nucleus, reducing attraction.

Noble Gases and EN

• He, Ne, Ar generally lack EN values as they are stable. Kr, Xe, and Rn can have detectable EN values due to their ability to form bonds under specific conditions.

Predicting Bond Type by EN Difference $\Delta\text{EN}$

• Nonpolar Covalent: Equal sharing of electrons (\Delta\text{EN} < 0.4). Occurs between identical atoms (e.g., H–H, O=O).
• Polar Covalent: Unequal sharing ($0.4 \le \Delta\text{EN} \le 1.7$). Electrons spend more time near the more electronegative atom.
  • Dipole Concept:
  • Partial Charges: $\delta^+$ on less EN atom, $\delta^-$ on more EN atom due to unequal sharing.
  • Dipole Moment ($\mu$): A measure of bond/molecular polarity. For a diatomic, $\boldsymbol{\mu} = q \times d$, where $q$ is partial charge magnitude and $d$ is distance.
  • Dipole Arrows: Point towards the more EN atom ($\delta^-$), with the tail indicating the $\delta^+$-side (e.g., H→Cl).
• Ionic: Electron transfer (\Delta\text{EN} > 1.7).

Ionic Bonding

• Occurs between a metal (electron donor, forming cation) and a nonmetal (electron acceptor, forming anion).
• Characterized by electron transfer and strong Coulombic attraction between resulting ions.
• In aqueous solution, ionic compounds dissociate into ions (e.g., $\text{NaCl}<em>{\text{(s)}} \rightarrow \text{Na}^{+}</em>{\text{(aq)}} + \text{Cl}^{-}_{\text{(aq)}}$).

Covalent Bonding

• Formed between nonmetals where electrons are shared.
  • Nonpolar Covalent: Equal sharing (e.g., H–H, O=O).
  • Polar Covalent: Unequal sharing due to EN differences, leading to partial charges and dipole moments.
• Covalent substances generally do not form ions or dissociate in water; pure water is H$_2$O, not an aqueous solution.

Metallic Bonding

• Occurs between metal atoms.
• Characterized by a "sea" of delocalized electrons surrounding positive metal nuclei.
• This free movement of electrons explains high electrical conductivity and other metallic properties.