10.3 Phase Transitions

Matter changes states through phase transitions. These shifts between solid, liquid, and gas are driven by temperature and pressure changes. Understanding these transitions helps explain everyday phenomena like boiling water or why ice cubes melt.

Intermolecular forces play a crucial role in phase transitions. Stronger forces require more energy to break, leading to higher melting and boiling points. This explains why water, with its strong hydrogen bonds, boils at a higher temperature than many other substances.

Phase Transitions

Phase transitions and key temperatures

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• Phase transitions change matter between solid, liquid, and gas phases
  • Melting converts solid to liquid by overcoming intermolecular forces
    • Key melting points at 1 atm pressure: water (0°C), ethanol (-114°C), gold (1064°C)
  • Vaporization converts liquid to gas when vapor pressure equals atmospheric pressure
    • Boiling vaporizes throughout the liquid at the boiling point
    • Key boiling points at 1 atm pressure: water (100°C), ethanol (78°C), gold (2856°C)
  • Sublimation converts solid directly to gas, skipping the liquid phase
    • Occurs with substances such as dry ice (solid CO2), iodine, naphthalene
  • Condensation converts gas to liquid by removing heat energy
  • Freezing converts liquid to solid at the same temperature as the melting point for a given pressure
  • Deposition converts gas directly to solid, the reverse of sublimation

Intermolecular forces in phase transitions

• Intermolecular forces (IMFs) attract molecules to each other
  • Stronger IMFs lead to higher temperatures needed for melting and boiling
• IMFs increase in strength from London dispersion forces to dipole-dipole forces to hydrogen bonding
  • London dispersion forces (LDFs) exist between all molecules due to temporary dipoles, increasing with molecular size and surface area
  • Dipole-dipole forces occur between polar molecules with permanent dipoles from uneven charge distribution
  • Hydrogen bonding, the strongest IMF, occurs when H bonds to highly electronegative N, O, or F
• Substances with stronger IMFs require more energy to overcome attractions and change phase
  • Water has strong hydrogen bonding, resulting in high melting (0°C) and boiling (100°C) points compared to similar-sized nonpolar methane (-182°C and -161°C)

Interpretation of heating and cooling curves

• Heating curves show temperature change as heat is added at a constant rate
  • Plateaus represent phase transitions where heat is added but temperature remains constant
  • Heat added during a phase transition equals the enthalpy (H) of that transition
    • $H_{fus}$ represents the enthalpy of fusion for melting/freezing
    • $H_{vap}$ represents the enthalpy of vaporization for boiling/condensing
    • $H_{sub}$ represents the enthalpy of sublimation
• Cooling curves show temperature change as heat is removed at a constant rate, the reverse of heating curves
• Calculate heat flow (q) and enthalpy changes (ΔH) using:
  1. $q = mcΔT$ for temperature changes (m = mass, c = specific heat capacity, ΔT = temperature change)
  2. $q = mΔH_{fus}$ or $q = mΔH_{vap}$ for phase transitions (m = mass, ΔH = enthalpy of fusion or vaporization)
  3. Enthalpy change (ΔH) equals the heat flow (q) at constant pressure
  • The energy absorbed or released during a phase change without temperature change is called latent heat

Phase diagrams and critical points

• Phase diagrams graphically represent the relationship between temperature, pressure, and physical state of a substance
• Key features of a phase diagram include:
  • Triple point: where solid, liquid, and gas phases coexist in equilibrium
  • Critical point: the highest temperature and pressure at which liquid and gas phases can coexist
  • Vapor pressure curve: shows the pressure at which a liquid and its vapor are in equilibrium at a given temperature