Chapter 16: Liquids and Solids

Chapter 16: Liquids and Solids

16.1 Intermolecular Forces

  • Definition: The attractive forces (electrostatic in nature) that exist between molecules, atoms, and ions.

  • Types of Molecular Forces:

    • Intermolecular Forces: Weaker forces that occur between molecules. ( extbf{Strength:} 2 - 20 kJ/mol)

    • Intramolecular Forces: Stronger forces that occur within molecules. ( extbf{Strength:} 100's - 1000's kJ/mol)

  • Example of Energy Comparison:

    • It is easier to vaporize water (overcome intermolecular forces) than to break O-H bonds (overcome intramolecular forces).

    • For 1 mole:

    • $ ext{ΔH}{vap}( ext{H}2 ext{O}) = 40.7 ext{ kJ}$

    • $ ext{D(O-H)} = 2(467 ext{ kJ})$

16.2 The Liquid State

  • Properties of the Liquid State:

    • Intermediate phase between solids and gases characterized by:

    1. Relatively strong intermolecular forces, leading to short-range order.

      • Example: Each ext{H2O} molecule hydrogen-bonds with 2-3 others in a tetrahedral arrangement.

      • Only about 10% of hydrogen-bonding interactions are broken during the phase change from ext{H}2 ext{O}(s) to ext{H}2 ext{O}(l).

      • Hydrogen bonds are continuously broken and reformed as liquid molecules move.

  • Comparison of Liquid and Solid States:

    • Liquids:

    • Assume shape of the portion of container they occupy.

    • Flow and mix, nearly incompressible, dense.

    • Solids:

    • Keep their own shape, virtually incompressible, mix only slowly, dense.

16.3 Properties Related to Intermolecular Forces

  • Many physical properties of solids, liquids, and gases can be explained through their intermolecular forces, including:

    • Melting point (m.p.), boiling point (b.p.),

    • Enthalpy of fusion ($ ext{ΔH}{fusion}$), enthalpy of vaporization ($ ext{ΔH}{vaporization}$)

    • Surface tension, Viscosity, Vapour pressure

  • Changes of State:

    • Caused by variations in intermolecular forces between molecules.

    • The molecules themselves remain intact.

16.4 Types of Intermolecular Forces

  • Van der Waals forces include:

    1. Dipole-dipole forces

    2. Hydrogen bonding

    3. London dispersion forces

A. Dipole-Dipole Forces

  • Occur between polar molecules.

  • Concept:

    • Polar bonds generate polar molecules with a net dipole moment (Σμ ≠ 0).

    • Example: $PF_3$ shows uneven electron distribution leading to partial positive ($ ext{δ}^+$) and negative ($ ext{δ}^-$) charges.

  • Polar Molecules:

    • Attracted to each other through their dipoles during the condensed phases, leading to semi-organized structures.

  • Strength of Dipole-Dipole Forces:

    • Depends on the magnitude of the dipole moment ($μ = rac{Q}{r}$) and shape of the molecule.

    • Generally operates over shorter distances than ion-ion interactions.

    • Example: Non-polar molecules like $CCl_4$ do not exhibit dipole-dipole forces despite having polar bonds.

B. Hydrogen Bonding

  • Definition: A special case of dipole-dipole forces occurring only in compounds where hydrogen is directly bonded to a highly electronegative atom such as fluorine (F), nitrogen (N), or oxygen (O).

  • Polar Bonds: Highly polar leading to strong intermolecular attractions.

  • Examples of Compounds: $H2O$, $NH3$, $CH_3COOH$.

  • Factors Affecting Strength of Hydrogen Bonds:

    1. Increasing the number of hydrogen-bonding centers in the molecule (e.g., Glycerol vs. Ethanol).

    2. Increasing the electronegativity of atoms bonded to hydrogen.

  • Comparison of Hydrogen Bonding with Regular Dipole-Dipole Forces:

  • Hydrogen bonding is generally stronger due to the small size of hydrogen, leading to closer intermolecular approaches, and minimal electron shielding on H.

C. London Dispersion Forces

  • Definition: All molecules, atoms, and ions possess dispersion forces due to the ability of electron clouds to distort creating temporary dipoles.

  • Types:

    1. Dipole-induced Dipole: Occurs when a polar molecule induces a dipole in a nearby non-polar molecule.

    2. Instantaneous Dipole-induced Dipole: Momentary distortion can induce a dipole moment in one molecule affecting a nearby molecule.

  • Strength Factors of Dispersion Forces:

    1. Increases with molar mass, size, and number of electrons (larger species are more polarizable).

    2. Greater for linear chains of atoms compared to branched structures (e.g., n-pentane vs. 2,2-dimethylpropane).

16.5 Properties of Bulk Matter

  • Compressibility: Liquids are only slightly compressible due to strong intermolecular forces that keep molecules in close proximity.

  • Lack of Rigidity: High kinetic energy of molecules offset the short-range order enforced by intermolecular forces.

  • High Densities: Results from strong intermolecular forces; for example, water density:

    • $d{H2O(s)} = 0.9168$ g/cm³ (solid),

    • $d{H2O(l)} = 0.9971$ g/cm³ (liquid),

    • $d{H2O(g)} = 3.26 imes 10^{-4}$ g/cm³ (gas).

  • Properties Related to State Changes:

    • Enthalpies of fusion ($ ext{ΔH}{fus}$) and vaporization ($ ext{ΔH}{vap}$) are reflective of intermolecular force strength.

  • Viscosity: Resistance to flow, increasing with strength of intermolecular forces and molecular complexity (e.g., higher for long-chain hydrocarbons).

  • Surface Tension: Tendency of liquids to minimize surface area, increases with stronger intermolecular forces and lower temperatures.

  • Capillary Action: Ability of liquids to rise in a narrow space, influenced by adhesive vs cohesive forces.

16.6 Vapour Pressure

  • Definition: The pressure exerted by a gas in dynamic equilibrium with its liquid (or solid) at a specific temperature in a closed container.

  • Vaporization: Endothermic process of molecules escaping from liquid to gas.

  • Condensation: Gas molecules re-entering the liquid state are part of the equilibrium process, keeping vapor pressure constant.

  • Properties of Vapor Pressure:

    1. Intensive property: independent of amount present. For example, $ ext{Pvap(H}_{2} ext{O(l)}) = 23.756 ext{ torr @ 25°C}$.

    2. Inversely proportional to strength of intermolecular forces; weaker forces lead to higher vapor pressures.

  • Liquid Types:

    • Volatile liquids: Easily evaporate (weak IMF, high Pvap).

    • Non-volatile solids: Generally exhibit low Pvap.

16.7 Phase Diagrams

  • Definition: A graphical representation showing the conditions (temperature, pressure) under which distinct states of matter (solid, liquid, gas) exist.

  • Common Features:

    1. Regions delineate states (solid, liquid, gas).

    2. Boundary lines indicate pressure and temperature at which phase transitions occur (e.g., melting point).

  • Notable Points in Phase Diagram for H₂O:

    • Normal Melting Point (Tm): $0°C$, dynamic equilibrium between solid and liquid at $1.00 ext{ atm}$.

    • Normal Boiling Point (Tb): $100°C$, dynamic equilibrium between liquid and gas at $1.00 ext{ atm}$.

    • Triple Point (P3/T3): $P3 = 0.0060 ext{ atm}$, $T3 = 0.0098 ext{ °C}$.

    • Critical Point: Defines the temperature above which the vapor cannot be liquefied regardless of pressure.

16.8 Energy and Changes of State

  • Enthalpy of Fusion ($ ext{ΔH}_{fus}$): Amount of heat to convert 1 g of solid to liquid.

  • Enthalpy of Solidification ($ ext{ΔH}_{sol}$): Heat released when turning liquid to solid.

  • Enthalpy of Vaporization ($ ext{ΔH}_{vap}$): Heat required to convert liquid to gas.

  • Energy Considerations in Heating Curves:

    • Horizontal sections correspond to phase changes (temperature constant).

    • Sloped sections correspond to state where temperature increases (temperature variable).

  • Supercooling and Superheating:

    • Supercooling occurs when a liquid's temperature falls below its freezing point without forming solid.

    • Superheating occurs when the temperature exceeds the boiling point without vaporizing.

16.9 Temperature Dependence of the Vapour Pressure

  • The vapor pressure increases with temperature due to more molecules acquiring sufficient kinetic energy to overcome intermolecular forces.

  • Clausius-Clapeyron Equation: Describes the relationship between vapor pressure and temperature:
    extln(P<em>vap)=−racextΔH</em>vapRrac1T+Cext{ln}(P<em>{vap}) = - rac{ ext{ΔH}</em>{vap}}{R} rac{1}{T} + C

  • Practical Example: Knowing $ ext{ΔH}{vap}$ for water is $40.7 ext{ kJ/mol}$, calculate $P{vap}$ at $37°C$.

  • Boiling Point: Temperature at which vapor pressure equals external atmospheric pressure, varies with atmospheric conditions.

16.10 Conclusion

  • Understanding intermolecular forces and their implications allows predictions about physical properties (e.g., boiling points, surface tension).

  • Comparing substances based on intermolecular forces provides insights into their behavior in different states (solid, liquid, gas).