BackMolecular Polarity and Intermolecular Forces (Chapter 7) – Study Notes
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Molecular Polarity & Intermolecular Forces
Introduction
This chapter explores how the polarity of molecules and the forces between them (intermolecular forces) determine many physical and chemical properties. Understanding these concepts is essential for predicting molecular behavior, solubility, boiling/melting points, and biological function.
Electronegativity and Bond Polarity
Electronegativity
Electronegativity is the tendency of an atom to attract a shared pair of electrons toward itself in a chemical bond.
Atoms with high electronegativity attract shared electrons more strongly (e.g., fluorine, oxygen).
Atoms with low electronegativity attract shared electrons less strongly (e.g., alkali metals).
Electronegativity generally increases across a period (left to right) and decreases down a group in the periodic table.
Bond Polarity
The difference in electronegativity between two bonded atoms determines the polarity of the bond.
If electrons are shared equally, the bond is nonpolar (e.g., H2).
If electrons are shared unequally, the bond is polar (e.g., HCl).
Types of Covalent Bonds
Nonpolar Covalent Bonds
Occur between nonmetals with equal or almost equal sharing of electrons.
Very small electronegativity difference (typically 0.0–0.4).
Atoms | Electronegativity Difference | Type of Bond |
|---|---|---|
N–N | 0.0 | Nonpolar covalent |
Cl–Br | 0.2 | Nonpolar covalent |
H–Si | 0.3 | Nonpolar covalent |
Polar Covalent Bonds
Occur between nonmetals with unequal sharing of electrons.
Moderate electronegativity difference (typically 0.5–1.8).
Atoms | Electronegativity Difference | Type of Bond |
|---|---|---|
O–Cl | 0.5 | Polar covalent |
Cl–C | 0.5 | Polar covalent |
O–S | 1.0 | Polar covalent |
Dipoles and Bond Polarity
A dipole is created when there is a separation of charge in a polar bond.
The positive and negative ends are indicated by the lowercase Greek letter delta (δ+ and δ−).
An arrow points from the positive to the negative end of the dipole.
Examples: C–O, N–O, Cl–F all have dipoles pointing toward the more electronegative atom.
Ionic Bonds
Occur between metal and nonmetal ions.
Result from electron transfer.
Large electronegativity difference (typically ≥ 1.8).
Atoms | Electronegativity Difference | Type of Bond |
|---|---|---|
Cl–K | 2.2 | Ionic |
N–Na | 2.1 | Ionic |
S–Cs | 1.8 | Ionic |
Summary Table: Electronegativity Differences and Bond Types
Electronegativity Difference | Bond Type | Electron Sharing |
|---|---|---|
0.0–0.4 | Nonpolar covalent | Electrons shared equally |
0.5–1.8 | Polar covalent | Electrons shared unequally |
>1.8 | Ionic | Electrons transferred |
Polarity of Molecules
Nonpolar Molecules
Molecules like H2, Cl2, and O2 are nonpolar because they contain only nonpolar bonds.
Molecules with polar bonds can still be nonpolar if the dipoles cancel due to symmetrical arrangement (e.g., CO2, CF4).
Polar Molecules
Polar molecules (e.g., HCl) have one end more negatively charged than the other.
The polar bonds do not cancel each other due to molecular geometry.
Electrons are shared unequally in the polar covalent bond.
Examples: Water and Ammonia
H2O: Two lone pairs and two bonds create a bent shape; dipoles do not cancel, making water polar.
NH3: One lone pair and three bonds create a trigonal pyramidal shape; dipoles do not cancel, making ammonia polar.
Practice Examples
Identify each molecule as polar or nonpolar: OF2, PBr3, HBr, CH4
Polarity of Large Molecules
Hydrocarbons
Composition: Only carbon (C) and hydrogen (H).
Bond polarity: C–H bonds are nonpolar (electronegativity difference = 0.4).
Molecular shape: Symmetrical, equal charge distribution.
Result: Hydrocarbons are nonpolar.
Example: Octane (C8H18), main component of gasoline.
Alcohols
Structure: Hydrocarbon chain with an –OH (hydroxyl) group.
Effect of –OH: The O–H bond is polar due to oxygen's high electronegativity.
Overall polarity:
Short chains (e.g., ethanol): Polar molecule.
Longer chains (e.g., pentanol): Both polar (–OH end) and nonpolar (hydrocarbon chain) regions.
Example: Pentanol (C5H11OH)
Nonpolar region: Hydrocarbon tail (C–C and C–H bonds).
Polar region: –OH group (can form hydrogen bonds with water).
Phospholipids
Structure: Large biomolecule with two long nonpolar hydrocarbon tails and one polar "head" group (phosphate, glycerol, choline).
Behavior:
Polar head is hydrophilic (water-attracting).
Nonpolar tails are hydrophobic (water-repelling).
Function: This dual nature allows phospholipids to form cell membranes—a bilayer with heads facing outward toward water and tails facing inward, away from water.
Practice Examples
Identify the nonpolar region(s) of the ibuprofen molecule.
Identify the polar region(s) of the cortisol molecule.
Intermolecular Forces (IMFs)
Definition and Importance
IMFs are attractive forces between molecules.
Weaker than chemical bonds within molecules, but crucial for determining physical properties (boiling/melting point, solubility, etc.).
Type | Acts Within / Between Molecules | Strength | Example |
|---|---|---|---|
Bonding forces (ionic, covalent) | Within molecules | Strong | O–H bond in water |
Intermolecular forces | Between molecules | Weaker | Attraction between two water molecules |
Types of Intermolecular Forces
1. Dipole-Dipole Forces
Occur between polar molecules.
Attractive forces between oppositely charged ends of polar molecules.
2. Hydrogen Bonds
Special type of dipole-dipole force.
Formed when hydrogen is bonded to F, O, or N and interacts with a lone pair on F, O, or N in another molecule.
Strongest intermolecular force; crucial in DNA structure and water's properties.
3. Dispersion Forces (London Forces)
Weak attractions between nonpolar molecules.
Caused by temporary dipoles that develop when molecules bump into each other.
Allow nonpolar molecules to form liquids and solids.
Comparison of Bonding and Attractive Forces
Type of Force | Particle Arrangement | Example | Strength |
|---|---|---|---|
Ionic bond | Na+Cl− | NaCl | Strong |
Covalent bond | Cl–Cl | Cl2 | Strong |
Hydrogen bond | H–F···H–F | H2O, NH3 | Moderate |
Dipole-dipole | δ+···δ− | HCl | Moderate |
Dispersion forces | Temporary dipoles | All molecules | Weak |
Melting Points and Attractive Forces
Melting points are related to the strength of attractive forces between molecules or compounds.
Lower melting points: Weak forces (dispersion).
Higher melting points: Stronger forces (hydrogen bonding).
Highest melting points: Ionic compounds (strong ionic forces).
Practice Examples
Identify the main type of attractive force in liquids of the following: NCl3, H2O, Br2, KCl, NH3.
For formaldehyde, besides London dispersion, which other intermolecular force is present?
Identify the strongest intermolecular force in HCN, CF4, and CH3CH2OH.
Select the correct drawing representing a hydrogen bond.
Additional info: All molecules experience London dispersion forces, but polar molecules may also experience dipole-dipole or hydrogen bonding, depending on their structure. The presence and strength of these forces explain many physical properties and biological functions.
