Intermolecular Forces & Properties of Liquids – Chapter 11 Review

States of Matter & Overview of Intermolecular Forces

• The chapter contrasts the behaviour of gases (well-described by the ideal–gas law PV = nRT) with liquids/solids, which are called “condensed phases” because particles are far closer together.
• In condensed phases attractive forces cannot be ignored. Evidence: every gas liquefies if sufficiently cooled (e.g. \text{N}_2 becomes liquid at -196\,^{\circ}\text{C}).
• Microscopic comparison (Fig. 11.1)
  • 300 mL liquid \text{N}_2 → > 200 L gas at 25\,^{\circ}\text{C} (huge volume change).
  • Same sample frozen shows almost no volume change between liquid/solid benzene, illustrating close packing.
• Electrostatic origins: Coulomb’s law
  F = -k\frac{q1q2}{d^2}
  (negative sign ⇒ attractive when charges opposite).
• Ionic solids: strong Coulombic lattice ∴ high BPs/MPs.
• Molecular substances: much weaker attractions yet still vital → grouped under van der Waals forces:
  1. Dipole–dipole (permanent–permanent)
  2. Dipole–induced-dipole (Debye)
  3. Induced-dipole–induced-dipole (London/dispersion).

Ion–Dipole Interactions (Section 11.2)

• Occur between ions and polar molecules; strength << ion–ion but >> molecule–molecule.
• Factors controlling magnitude (from Coulomb’s law):
  1. d, distance ion ↔ dipole (smaller d ⇒ stronger).
  2. |q_{\text{ion}}| (higher charge ⇒ stronger).
  3. Dipole moment \mu of molecule.
• Hydration (solvation in water) enthalpies quantify interaction. Example (alkali cations, Table 11.1): Li⁺ (78 pm) \Delta_\text{hyd}H^{\circ} = -515\,\text{kJ mol}^{-1} vs. K⁺ (133 pm) -321\,\text{kJ mol}^{-1}.
  • Smaller radius → shorter d → more exothermic ✔
  • Higher charge dramatically increases magnitude: Mg²⁺ (79 pm) -1922\,\text{kJ mol}^{-1}.
• Example 11.1 rationalises order Mg²⁺ > Na⁺ > Cs⁺.
• Hydrated salts: water molecules can occupy lattice voids or bind directly to cation (e.g. \ce{CrCl3·6H2O} is really [\ce{Cr(H2O)4Cl2}]Cl·2H2O).

Dipole–Dipole Forces & Regular Polarity (Section 11.3)

• Permanent dipoles attract: + end toward - end of neighbour.
• Influence:
  • \Delta{\text{vap}}H ↑ with dipole strength (Table 11.2): compare \text{Br}2 vs. iso-mass ICl.
  • Boiling point ↑ with polarity for similar molar mass.
  • Solubility rule: “like dissolves like.” Polar ethanol miscible with water; non-polar gasoline not.
• Energy diagram of evaporation (Fig. 11.2): \Delta_{\text{vap}}H^{\circ} >0 endothermic; condensation releases same magnitude.

Hydrogen Bonding – a Special Dipole–Dipole

• Requirements: H covalently bonded to highly electronegative N, O, or F (X–H) interacting with lone-pair bearing N/O/F (Y).
• Typical energy 5–30 kJ mol⁻¹.
• Explains anomalously high BPs of \ce{HF, H2O, NH3} compared with group trends (Fig. 11.4).
• Example 11.2: ethanol (has O–H) boils at 78.3\,^{\circ}\text{C}; dimethyl ether (no O–H) -24.8\,^{\circ}\text{C}.

Hydrogen Bonding & Water’s Anomalies

• Each \ce{H2O} can form 4 H-bonds (2 donors + 2 acceptors) → open tetrahedral “ice” lattice (Fig. 11.5) ⇒ large empty space.
• Consequences:
  • Ice density ≈ 0.917 g cm⁻³ < liquid water ⇒ ice floats.
  • Maximum liquid density at 4\,^{\circ}\text{C} (Fig. 11.6). Lakes freeze top-down; overturn at 4 °C brings O₂/nutrients.
  • Enormous specific heat 4.184\,\text{J g}^{-1}\,\text{K}^{-1}: moderates climate.

Biological Significance (box)

• DNA base pairing relies on precise hydrogen bonds: adenine···thymine (two H-bonds); guanine···cytosine (three). Determines double-helix stability.

Forces Involving Non-Polar Molecules (Section 11.4)

Dipole–Induced-Dipole (Debye) Forces

• Polar molecule polarises electron cloud of non-polar partner (O₂ in water; I₂ in ethanol). Depend on polarizability α.
• Larger atoms/molecules (many electrons) → larger α → stronger interaction & greater solubility in polar solvents (Table 11.3 shows gas solubility increases \ce{H2} < N2 < O2).

London (Induced-Induced) Dispersion Forces

• Arise from instantaneous correlated electron motions creating temporary dipoles (Fig. 11.7).
• Universal (exist even in polar species) and often dominant (energy partition Fig. 11.8).
• Magnitude grows with molar mass & “contact area” (e.g. \text{I}_2 solid at RT).

Comparative Summary (Table 11.5)

Gecko Example (sidebar)

• Gecko foot setae & spatulae (≈200 nm) exploit massive cumulative London forces to adhere to walls/ceilings – can support 45 lb per dime-sized area.

Properties of Liquids (Section 11.6)

Vaporization & Condensation

• Endothermic vaporization \Delta_{\text{vap}} H^{\circ} >0 ; reverse exothermic.
• Example 11.5: 0.925 L water at 100 °C needs 2.0×10^{3}\,\text{kJ} to evaporate (using \Delta_{\text{vap}}H^{\circ}=40.7\,\text{kJ mol}^{-1}).

Equilibrium Vapor Pressure & Volatility

• Dynamic equilibrium when rate${\text{evap}}$ = rate${\text{cond}}$ inside closed vessel (Fig. 11.11).
• Vapor pressure ↑ with T (more molecules surpass energy threshold, Fig. 11.12).
• Example 11.6: in sealed 4.25×10⁴ L room at 25 °C only ≈ 1 L of a 2 L pan of water must evaporate to reach P_{\text{eq}}=23.8 mm Hg.

Clausius–Clapeyron Equation

• Linear relation:
  \ln P = -\dfrac{\Delta{\text{vap}} H}{R}\,\dfrac{1}{T} + C or two-point form \ln!\left(\dfrac{P2}{P1}\right)= -\dfrac{\Delta{\text{vap}} H}{R}\left(\dfrac{1}{T2}-\dfrac{1}{T1}\right)
• Slope of \ln P vs 1/T gives \Delta_{\text{vap}}H; for water slope -4.90×10^{3}\,\text{K} → 40.7\,\text{kJ mol}^{-1}.

Boiling Point

• Defined where P{\text{vap}} = P{\text{ext}}.
• Normal BP: P_{\text{ext}}=760 mm Hg.
• Lower external pressure (high altitude) ⇒ lower BP (e.g. Denver ~95 °C water) ⇒ longer cooking.

Critical Temperature Tc & Pressure Pc

• Beyond critical point no separate liquid/vapor phases – supercritical fluid (density ~ liquid, viscosity ~ gas).
• \text{CO}2: Tc = 30.99\,^{\circ}\text{C},\; P_c = 72.8 atm (Fig. 11.15). Used industrially:
  • Caffeine extraction (decaf coffee)
  • Hop-oil extraction, algae processing.
• Table 11.7 lists Tc, Pc for common substances.

Surface Tension

• Energy to increase surface area by 1 m² (units \text{J m}^{-2} or \text{N m}^{-1}).
• Origin: surface molecules experience net inward cohesive force (Fig. 11.16). Minimised by forming spheres (dew drops).

Capillary Action & Meniscus Formation

• Adhesive forces (liquid–wall) vs. cohesive (liquid–liquid) decide shape.
  • Water + glass: strong adhesion ⇒ concave meniscus & rise up capillary (Fig. 11.17).
  • Hg + glass: cohesion dominates ⇒ convex meniscus.
• Drives movement of water in paper, soils, plants.

Viscosity (Resistance to Flow)

• Greater intermolecular forces/chain length ⇒ higher viscosity.
  • Olive oil ≈ 70× viscous vs ethanol.
  • Honey viscous due to extensive H-bond network of sugars.
• Demonstrated by steel-ball drop or motor-oil comparison (Fig. 11.18).

Real-World/Applied Boxes

• Chromatography separation relies on differential intermolecular interactions between mobile & stationary phases (partition chromatography example with C-18 column; GC elution order CH₅ → C₈ explained by dispersion strength).
• Pet-food crisis (2007): melamine + cyanuric acid formed insoluble hydrogen-bonded melamine cyanurate → kidney stones in pets and infants. Highlights analytical need & toxicity via H-bonding networks.

Key Equations & Numerical References

• Coulomb: F = -k\dfrac{q1q2}{d^2}
• Hydration energy trend explained via \Delta H \propto \dfrac{q_{\text{ion}}}{d}.
• Clausius–Clapeyron (above).
• Vaporization enthalpy definition: \text{liquid} \xrightarrow[]{\Delta{\text{vap}}H^{\circ}} \text{vapor} ; reverse -\Delta{\text{vap}}H^{\circ}.

Using these notes you can now predict physical properties (BP, solubility, viscosity, surface behaviour) from molecular structure, rationalise environmental/biological phenomena, and perform quantitative vapor-pressure/enthalpy calculations with confidence.