Electronegativity, Molecular Polarity, and Intermolecular Forces: Study Guide

Electronegativity

Definition and Trends

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a fundamental concept in understanding chemical bonding and molecular polarity.

• Definition: Electronegativity quantifies how strongly an atom attracts electrons in a covalent bond.

• Trend Down a Group: Electronegativity decreases as you move down a group in the periodic table because atomic radius increases, and the outer electrons are farther from the nucleus and more shielded by inner electrons.

• Trend Across a Period: Electronegativity increases as you move from left to right across a period due to increasing nuclear charge and decreasing atomic radius, which enhances the nucleus's pull on bonding electrons.

• Example: Fluorine is the most electronegative element, while cesium is among the least.

Electronegativity Order of Elements

Arranging Elements by Electronegativity

Elements can be arranged in order of increasing electronegativity using periodic trends.

• Order Example: For the elements H, S, N, Al: Al < H < S < N

• Application: This order helps predict bond polarity and molecular properties.

Molecular Polarity

Determining Polarity

Molecular polarity depends on both the electronegativity difference between atoms and the geometry of the molecule.

• Polar Molecule: Has an uneven distribution of electron density, resulting in partial positive and negative charges.

• Nonpolar Molecule: Has an even distribution of electron density, so no significant partial charges exist.

• Example: Water (H2O) is polar; carbon dioxide (CO2) is nonpolar.

Dipole and Partial Charges

Dipole Definition and Drawing Lewis Structures

A dipole is a separation of charge within a molecule due to differences in electronegativity between bonded atoms.

• Dipole Moment: A vector quantity representing the direction and magnitude of charge separation.

• Partial Charges: Indicated by δ+ (partial positive) and δ− (partial negative) on atoms in a polar bond.

• Lewis Structure: Shows the arrangement of atoms, bonds, and lone pairs in a molecule.

Intermolecular Forces

Types and Effects

Intermolecular forces are attractions between molecules that affect physical properties such as boiling and melting points.

• London Dispersion Forces: Weak, temporary attractions due to momentary dipoles in all molecules, especially nonpolar ones.

• Dipole-Dipole Forces: Attractions between permanent dipoles in polar molecules.

• Hydrogen Bonds: Strong dipole-dipole interactions involving H bonded to N, O, or F.

• Ionic Bonds: Electrostatic attractions between oppositely charged ions.

Comparison Table of Intermolecular Forces

Type of Force	Strength	Occurs Between	Example
London Dispersion	Weakest	All molecules (especially nonpolar)	CH4
Dipole-Dipole	Intermediate	Polar molecules	HCl
Hydrogen Bond	Strong	H bonded to N, O, or F	H2O
Ionic Bond	Strongest	Ions	NaCl

Lewis Structures, Shapes, and Polarity

Analysis of Example Molecules

Drawing Lewis structures helps determine molecular shape and polarity, which in turn affects intermolecular forces.

Molecule	Shape	Polar or Nonpolar	Intermolecular Forces
H2O	Bent	Polar	Hydrogen Bond, Dipole-Dipole, London Dispersion
CH4	Tetrahedral	Nonpolar	London Dispersion
CO2	Linear	Nonpolar	London Dispersion
NH3	Trigonal Pyramidal	Polar	Hydrogen Bond, Dipole-Dipole, London Dispersion

Boiling Point and Intermolecular Forces

Relationship

The strength of intermolecular forces directly affects boiling points. Molecules with stronger forces (e.g., hydrogen bonds) have higher boiling points.

• Example: H2O has a higher boiling point than CH4 due to hydrogen bonding.

Key Equations

Electronegativity Difference and Bond Type

• Bond Polarity: The difference in electronegativity ($riangle EN$) determines bond type:

  • Nonpolar Covalent: $riangle EN < 0.4$

  • Polar Covalent: $0.4 ext{ to } 1.7$

  • Ionic: $riangle EN > 1.7$

Additional info:

• London dispersion forces are generally weaker than dipole-dipole and hydrogen bonds because they arise from temporary fluctuations in electron distribution.

• Polar molecules can induce dipoles in nonpolar molecules, leading to dipole-induced dipole interactions.