Intermolecular Forces, Polarity & Physical Properties

States of Matter & Why Physical Forms Differ

  • Matter exists mainly as solid, liquid, gas (plasma is 4th, rarely tested here).
  • Physical state at a given T/P is governed by the COMPETITION between:
    • Intramolecular forces (bonds within a molecule/lattice)
    • Intermolecular forces (IMF) (attractions between separate particles)
  • Key idea: Intramolecular » Intermolecular in strength, yet IMFs control melting point (m.p.), boiling point (b.p.), volatility, viscosity, surface tension, solubility.
  • Examples introduced in the clip:
    • \text{I}2 (solid @ RT) vs \text{Cl}2 (gas @ RT): heavier \text{I}_2 experiences much stronger London dispersion ⇒ higher m.p. and b.p.
    • \text{NaCl}: ionic lattice with enormous Coulombic attractions ⇒ very high m.p./b.p. (solid under ordinary conditions).

Hierarchy of Intermolecular Forces (weak → strong)

  1. London dispersion (LDF, instantaneous‐induced dipole)
    • Present in all atoms/molecules; sole IMF for non-polar species.
    • Strength ∝ polarizability ∝ molecular weight/size.
    • Explains rising b.p. series: \text{He} < \text{Ne} < \text{Ar} < \text{Kr} < \text{Xe}.
  2. Dipole–dipole
    • Requires permanent molecular dipole (polar but not necessarily H-bonding).
    • Orientation: partial + on one molecule aligns with partial – on neighbor.
  3. Hydrogen bonding (special case of dipole–dipole)
    • Donor: H covalently bonded to F, O, or N.
    • Acceptor: lone pair on F, O, or N.
    • Responsible for ice floating, protein secondary structure, DNA base pairing.
  4. Ion–dipole
    • Interaction between a full charge (ion) and a molecular dipole.
    • Strongest IMF; critical for dissolving salts in polar solvents.
    • Ex: \text{Na}^+ attracted to the O end (δ–) of \text{H}_2\text{O} whereas \text{Cl}^- aligns with the H end (δ+).

Visual: Hydration of Ions (Water Orientation)

  • Water is bent; O carries two lone pairs ⇒ δ– on O, δ+ on each H.
  • Around \text{Na}^+ the O atoms face inward; around \text{Cl}^- the H atoms point inward.

Quantitative Notes & Formulas

  • Coulomb’s Law for ionic lattice (simplified pairwise):
    F \;\propto\; \frac{q1 q2}{r^2} (large charges & small radii → strong force).
  • Typical bond energy order (kJ·mol^{-1}):
    • Ionic > Covalent \approx Metallic (intramolecular) ≫ Ion–dipole (≈50) > H-bond (10‒40) > Dipole–dipole (5‒25) > LDF (0.05‒40; highly size-dependent).
  • Electronegativity (EN) guidelines for bond type (rough):
    • \Delta EN > 1.7 → predominantly ionic.
    • 0 < \Delta EN < 1.7 → polar covalent.
    • \Delta EN \approx 0 → non-polar covalent.

London Dispersion & Molecular Weight Trend

  • Empirical graph (described by lecturer): y-axis = b.p. (K), x-axis = molar mass (g·mol^{-1}).
  • Straight-line or smooth upward curve for homologous series shows b.p. ↑ with mass.
  • Example data (diatomic halogens):
    • \text{F}_2: −188 °C, 38 g·mol^{-1}
    • \text{Cl}_2: −34.6 °C, 71 g·mol^{-1}
    • \text{Br}_2: 59 °C, 160 g·mol^{-1}
    • \text{I}_2: 184 °C, 254 g·mol^{-1}

Classifying Compounds: Ionic vs Covalent

  • Rule of thumb: metal + non-metal ⇒ ionic; non-metal + non-metal ⇒ covalent.
  • Examples from the clip:
    • \text{NCl}_3: N (non-metal) + Cl (non-metal) ⇒ covalent.
    • \text{Br}_2: same element ⇒ non-polar covalent; EN(Br)=2.8 ⇒ \Delta EN = 0.
    • \text{NO}_2^- (nitrite ion): polyatomic ion; bonding inside ion is covalent, but interactions with cations in lattice are ionic; resonance needed to see charge distribution.

Molecular Geometry & Polarity Pointers

  • Non-polar molecules CANNOT exhibit dipole–dipole interactions (they have no permanent dipole).
  • Criteria for non-polarity:
    • All peripheral atoms identical and molecular geometry is symmetrical (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral).
  • Example: \text{CO}_2 is linear & symmetrical ⇒ net dipole =0 even though C−O bonds are polar.
  • Hints given:
    • "What do we know about linear? Symmetric → dipole cancels."

Worked Question Snapshots (from lecture)

  1. "Which of these molecules have dipole–dipole forces?"
    • Strategy: test polarity. Non-polar? → only LDF. Polar? → dipole–dipole present (plus LDF).
  2. "Which substance has the highest boiling point?"
    • First eliminate non-polar options (lowest IMF).
    • Compare remaining based on IMF hierarchy & molar mass.
    • Example showdown: \text{CO} vs \text{CH}_3\text{OH}
      • \text{CO}: only dipole–dipole (and LDF), small MW.
      • \text{CH}_3\text{OH}: hydrogen bonding + heavier; hence higher b.p.

Connections to Other Lectures & Real-World Relevance

  • Previous session: atomic structure & periodic trends → now leveraged (EN, radii).
  • Future: solutions, colligative properties rely heavily on IMF strength and type.
  • Real-world examples:
    • Volatility of gasoline (mixture of non-polar hydrocarbons) due to weak LDF.
    • Anti-freeze (ethylene glycol) resists vaporization thanks to extensive H-bonding.
    • Protein folding: delicate balance of H-bonds, LDF (a.k.a. van der Waals), and ionic salt bridges.

Ethical & Practical Implications Mentioned

  • Safe handling of iodine (solid) vs chlorine (toxic gas) – must store Cl_2 as pressurized liquid.
  • Pharmaceutical design: modulating IMF (e.g., adding OH for H-bonding) can enhance water solubility.

Summary Cheat-List

  • Intramolecular > Intermolecular (memorize order of IMF strength).
  • Non-polar ⇒ only LDF.
  • Heavier / larger ⇒ stronger LDF ⇒ higher b.p.
  • H-bonding needs H–F, H–O, or H–N.
  • Ion dissolved in polar solvent ⇒ ion–dipole.
  • Symmetrical geometry cancels dipoles, even if bonds themselves are polar.