10.3 Phase Transitions

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, melting ice cubes, and even why frost forms on cold mornings.

Intermolecular forces play a crucial role here. Stronger forces require more energy to break, which leads to higher melting and boiling points. Water, for example, has strong hydrogen bonds, so it boils at a much higher temperature than you'd expect for such a small molecule.

Phase Transitions and Key Temperatures

Each phase transition involves either adding or removing energy to overcome (or give in to) the attractions between molecules. Here are the six transitions you need to know:

• 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. Boiling is a specific type of vaporization that happens throughout the liquid when vapor pressure equals atmospheric pressure.
  • 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 entirely. Dry ice (solid $CO_2$) is the classic example. Iodine and naphthalene (mothballs) also sublime.
• Condensation converts gas to liquid by removing heat energy. Think of water droplets forming on the outside of a cold glass.
• Freezing converts liquid to solid. This happens at the same temperature as melting for a given substance at a given pressure.
• Deposition converts gas directly to solid, the reverse of sublimation. Frost forming on a window is deposition of water vapor.

Notice the pattern: melting/freezing, vaporization/condensation, and sublimation/deposition are reverse pairs.

Intermolecular Forces in Phase Transitions

Intermolecular forces (IMFs) are the attractions between molecules (not the bonds within a molecule). The stronger these attractions, the more energy you need to pull molecules apart and change phase.

IMFs increase in strength in this order:

• London dispersion forces (LDFs) exist between all molecules. They arise from temporary, instantaneous dipoles. Larger molecules with more surface area have stronger LDFs because their electron clouds are bigger and more easily distorted.
• Dipole-dipole forces occur between polar molecules that have permanent partial charges from uneven electron distribution.
• Hydrogen bonding is the strongest common IMF. It occurs when hydrogen is bonded to a highly electronegative atom: nitrogen, oxygen, or fluorine. The H atom carries a strong partial positive charge that attracts lone pairs on N, O, or F of neighboring molecules.

A quick comparison shows why IMF strength matters so much: water ($H_2O$) has strong hydrogen bonding and boils at 100°C. Methane ($CH_4$) is a similar-sized molecule but is nonpolar with only LDFs, so it boils way down at −161°C. That 261°C difference comes entirely from the difference in intermolecular forces.

Heating and Cooling Curves

A heating curve plots temperature on the y-axis against heat added on the x-axis. It tells you exactly what happens as you steadily pump energy into a substance.

The curve has two types of regions:

• Sloped sections where the temperature rises. Here, added heat increases the kinetic energy of the molecules. You calculate heat in these sections using: $q = mc\Delta T$ where $m$ = mass, $c$ = specific heat capacity, and $\Delta T$ = temperature change.
• Flat plateaus where the temperature stays constant even though heat is still being added. All the energy goes into breaking intermolecular forces (changing phase), not raising temperature. This energy is called latent heat. You calculate it using: $q = m\Delta H_{fus}$ for melting/freezing, or $q = m\Delta H_{vap}$ for boiling/condensing.

Key enthalpy values to know:

• $\Delta H_{fus}$ = enthalpy of fusion (melting or freezing)
• $\Delta H_{vap}$ = enthalpy of vaporization (boiling or condensing)
• $\Delta H_{sub}$ = enthalpy of sublimation

A cooling curve is the reverse: temperature drops in the sloped sections, and plateaus appear where the substance releases energy as it changes phase. At constant pressure, the enthalpy change $\Delta H$ equals the heat flow $q$.

Phase Diagrams and Critical Points

A phase diagram is a graph with pressure on the y-axis and temperature on the x-axis. It maps out which phase a substance exists in at any combination of temperature and pressure.

Key features to identify:

• Triple point: the single temperature and pressure where solid, liquid, and gas all coexist in equilibrium. Every pure substance has exactly one triple point.
• Critical point: the highest temperature and pressure at which distinct liquid and gas phases can exist. Beyond this point, the substance becomes a supercritical fluid where liquid and gas properties merge.
• Phase boundary lines separate the regions on the diagram. The line between liquid and gas is the vapor pressure curve, showing the pressure at which liquid and vapor are in equilibrium at each temperature. The line between solid and liquid shows how melting point changes with pressure.

To use a phase diagram, find your temperature and pressure on the axes. Where those values intersect tells you the phase of the substance under those conditions.