Molecular Polarity and Intermolecular Forces

Molecular Polarity

Definition and Importance

Molecular polarity refers to the distribution of electrical charge over the atoms joined by the bond. It determines many physical properties of substances, such as solubility, melting point, and boiling point.

• Polar molecules have regions of partial positive and negative charge due to unequal sharing of electrons.

• Nonpolar molecules have an even distribution of charge, with electrons shared equally.

• Polarity affects how molecules interact with each other and with solvents (e.g., water).

Types of Covalent Bonds

Covalent bonds can be classified based on how equally electrons are shared between atoms.

• Nonpolar covalent bond: Electrons are shared equally between two atoms of similar or identical electronegativity (e.g., Cl2).

• Polar covalent bond: Electrons are shared unequally due to a difference in electronegativity, resulting in partial charges (e.g., H–Cl).

Electronegativity and Bond Polarity

Electronegativity (EN) is a measure of an atom's ability to attract shared electrons in a chemical bond. The greater the difference in EN between two atoms, the more polar the bond.

• ΔEN (Delta EN): The difference in electronegativity between two bonded atoms.

• As ΔEN increases, bond polarity increases.

Example Table: Electronegativity Values

Element	Electronegativity
H	2.1
O	3.5
C	2.5
P	2.1
F	4.0
S	2.5

Calculating ΔEN for Bonds:

• H–O: ΔEN = |3.5 – 2.1| = 1.4

• H–H: ΔEN = |2.1 – 2.1| = 0

• S–O: ΔEN = |3.5 – 2.5| = 1.0

• P–F: ΔEN = |4.0 – 2.1| = 1.9

• S–P: ΔEN = |2.5 – 2.1| = 0.4

Ranking Bond Polarity (from least to most polar):

1. H–H (0)

2. S–P (0.4)

3. S–O (1.0)

4. H–O (1.4)

5. P–F (1.9)

Determining Molecular Polarity

To determine if a molecule is polar, consider both the polarity of its bonds and its molecular geometry (shape).

• Polar bonds are necessary but not sufficient for a molecule to be polar.

• Asymmetrical shape (e.g., bent or trigonal pyramidal) leads to a polar molecule if polar bonds are present.

• Symmetrical shape (e.g., linear, tetrahedral with identical atoms) can result in a nonpolar molecule even if it contains polar bonds, because the dipoles cancel out.

Examples:

• H2O (water): Polar molecule due to bent shape and polar O–H bonds.

• CO2 (carbon dioxide): Nonpolar molecule due to linear shape; dipoles cancel.

• NH3 (ammonia): Polar molecule due to trigonal pyramidal shape and polar N–H bonds.

• CH4 (methane): Nonpolar molecule due to tetrahedral symmetry; dipoles cancel.

Intermolecular Forces

Types of Intermolecular Forces

Intermolecular forces are attractions between molecules, influencing physical properties like boiling and melting points.

• London Dispersion Forces: Weak, temporary attractions due to momentary uneven electron distribution. Present in all molecules, but dominant in nonpolar molecules. Strength increases with molecular size and surface area.

• Dipole-Dipole Interactions: Attractions between the positive end of one polar molecule and the negative end of another. Only present in polar molecules.

• Hydrogen Bonds: Strong dipole-dipole interactions occurring when hydrogen is bonded to highly electronegative atoms (N, O, or F). Responsible for unique properties of water.

• Ionic Bonds: Electrostatic attractions between fully charged ions (not technically intermolecular, but important for comparison).

Summary Table: Types of Intermolecular Forces

Type	Occurs Between	Relative Strength
London Dispersion	All molecules (especially nonpolar)	Weakest
Dipole-Dipole	Polar molecules	Intermediate
Hydrogen Bond	Molecules with N–H, O–H, or F–H bonds	Strongest (of intermolecular)
Ionic Bond	Fully charged ions	Very strong (intramolecular)

Ranking Intermolecular Forces

• 1 (Weakest): London Dispersion

• 2: Dipole-Dipole

• 3 (Strongest): Hydrogen Bond

Example: Water (H2O) molecules form hydrogen bonds with each other, leading to high boiling point and surface tension.

Polarity and Solubility

Polarity affects how substances dissolve:

• "Like dissolves like": Polar substances dissolve in polar solvents (e.g., salt in water); nonpolar substances dissolve in nonpolar solvents (e.g., oil in hexane).

• Polar and nonpolar substances generally do not mix well.

Hydrophobic vs. Hydrophilic

• Hydrophobic: "Water-fearing"; nonpolar molecules that do not interact favorably with water (e.g., oils, fats).

• Hydrophilic: "Water-loving"; polar or charged molecules that interact well with water (e.g., sugars, salts).

Soap Molecules: Amphipathic Nature

Soap molecules have both polar (hydrophilic) and nonpolar (hydrophobic) regions, allowing them to interact with both water and oils.

• Polar head: Interacts with water (hydrophilic).

• Nonpolar tail: Interacts with oils and grease (hydrophobic).

Key Definitions and Concepts

• Electronegativity (EN): The tendency of an atom to attract electrons in a bond.

• Partial charge (δ+, δ–): Slight positive or negative charge due to unequal electron sharing.

• Symmetry: Molecular shape affects whether dipoles cancel (nonpolar) or reinforce (polar).

• Intermolecular force: Attraction between molecules, not within them.

Summary Table: Bond and Molecular Polarity

Bond Type	Electron Sharing	Polarity
Nonpolar Covalent	Equal	Nonpolar
Polar Covalent	Unequal	Polar
Ionic	Electron transfer	Very polar (full charges)

Example: In water (H2O), the O–H bonds are polar, and the bent shape makes the molecule polar overall.

Additional info: The notes also reference the "tug-of-war" analogy for electronegativity, which helps visualize how atoms with higher EN pull electrons more strongly, creating partial charges and molecular polarity.