Phase Changes and Phase Diagrams (Video Notes)

Learning Objectives

• Describe the different types of phase changes: melting, freezing, vaporization, condensation, sublimation, deposition.
• Explain phase changes using the concept of energy and intermolecular forces.
• Interpret a phase diagram and identify the triple point and critical point.
• Relate the importance of phase changes to real-life applications such as food preservation, climate systems, and industrial processes.

Phase Changes in Matter

• Phase change: process by which matter changes from one physical state or phase to another (solid, liquid, gas).
• The three most common phases: solid, liquid, gas.
• Examples:
  • Solid to liquid: melting
  • Liquid to gas: vaporization (boiling, evaporation)
• Transitions occur due to changes in energy input and/or environmental conditions (temperature, pressure).

Phase Changes Through Heating

• Heating generally increases the kinetic energy of molecules, leading to phase changes when bonds break or intermolecular forces are overcome.

Melting

• Phase change: solid → liquid due to an increase in temperature.
• Process: as solid is heated, molecular vibrations increase; bonds weaken; particles move more freely, turning solid into liquid.
• Example: melting popsicles.
• Key point: occurs at the solid–liquid boundary when sufficient energy is supplied to overcome rigid packing.
• Related concepts: melting point depends on pressure and impurities (and context of the system).

Boiling

• Phase change: liquid → gas due to an increase in temperature and pressure (overall energy input driving bulk conversion).
• Process: heating causes molecules to move rapidly and collide more; at sufficiently high temperature/pressure, liquid becomes gas throughout the volume (boiling).
• Example: boiling dry ice (CO₂) under appropriate conditions.

Evaporation

• Phase change: liquid → gas at a temperature below the boiling point, driven by energy supplied to surface molecules.
• Process: more energetic surface molecules escape into the air as gas while the liquid remains.
• Example: evaporation of water at room temperature.
• Key distinction: evaporation is a surface phenomenon and can occur at temperatures well below the boiling point.

Sublimation

• Phase change: solid → gas directly, skipping the liquid phase.
• Occurs when temperature and pressure are such that melting is unfavorable and the solid transitions straight to the gas.
• Can occur with heat input or pressure reduction (or both).
• Example: sublimation of dry ice (CO₂).

Phase Changes Through Cooling

• Cooling generally reduces kinetic energy, promoting phase transitions toward more ordered states.

Freezing

• Phase change: liquid → solid when the temperature drops below the freezing point.
• Process: molecules arrange into a fixed, rigid structure due to stronger intermolecular interactions at lower energy.
• Factors affecting freezing point: pressure and presence of impurities.
• Example: freezing of ice cubes.

Condensation

• Phase change: gas → liquid when the temperature drops below the dew point (or when pressure/temperature conditions favor liquid formation).
• Process: gas molecules come closer together and form a liquid as they lose enough energy to bind.
• Example: condensation of water droplets on a cool surface.

Deposition

• Phase change: gas → solid directly, skipping the liquid phase.
• Occurs when gas is cooled at a rate or pressure such that solid formation is favored below the sublimation point.
• Consequences: frost, snowflakes, and other ice crystals can form via deposition.

What Determines When and How a Substance Changes?

• Phase changes depend on two main variables:
  • Pressure
  • Temperature
• Under different P–T conditions, substances exist as solid, liquid, or gas; line boundaries separate phases on a phase diagram.
• The phase diagram shows regions for solid, liquid, gas, and a supercritical fluid region at high pressure/temperature.

Phase Diagram

• Phase diagram shows boundaries between phases:

  • Melting (solid ↔ liquid)
  • Freezing (liquid ↔ solid)
  • Vaporization (liquid ↔ gas)
  • Condensation (gas ↔ liquid)
  • Sublimation (solid ↔ gas)
  • Deposition (gas ↔ solid)

• Key features:

  • Solid, Liquid, Gas regions separated by phase boundaries.
  • Triple Point (Pt, Tt): all three phases coexist in equilibrium.
  • Critical Point (Pc, Tc): beyond this point, liquid and gas become indistinguishable; supercritical fluid region exists above this point.
  • Temperature and pressure axes illustrate how increasing either can drive phase transitions.

• Notation and points to remember:

  • Triple Point: Pt, Tt where all three phases coexist.
  • Critical Point: Pc, Tc where the liquid and gas phases are indistinguishable.
  • On a typical PvT diagram, sublimation occurs on the solid–gas boundary, deposition on the gas–solid boundary, melting/freezing on the solid–liquid boundary, and vaporization/condensation on the liquid–gas boundary.

Critical Temperatures and Pressures of Some Simple Substances

• Data table (Tc in °C, Pc in atm):

  • NH₃: Tc = 132.4^\circ\ ext{C}, \, Pc = 113.5\,\text{atm}
  • CO₂: Tc = 31.0^\circ\text{C}, \, Pc = 73.8\,\text{atm}
  • CH₃CH₂OH (ethanol): Tc = 240.9^\circ\text{C}, \, Pc = 61.4\,\text{atm}
  • He: Tc = -267.96^\circ\text{C}, \, Pc = 2.27\,\text{atm}
  • Hg: Tc = 1477^\circ\text{C}, \, Pc = 1587\,\text{atm}
  • CH₄: Tc = -82.6^\circ\text{C}, \, Pc = 46.0\,\text{atm}
  • N₂: Tc = -146.9^\circ\text{C}, \, Pc = 33.9\,\text{atm}
  • H₂O: Tc = 374.0^\circ\text{C}, \, Pc = 217.7\,\text{atm}

• These values illustrate how different substances have very different phase behavior and phase boundaries.

Water Phase Diagram (Illustrative Features)

• Water (H₂O) phase diagram highlights:

  • Triple Point: Tt = 0.01^\circ\text{C}, \, Pt \approx 0.611\,\text{kPa}
  • Critical Point: Tc = 374.0^\circ\text{C}, \, Pc = 217.7\,\text{atm}
  • At pressures and temperatures below the triple point, deposition or sublimation can occur.
  • At temperatures above the critical point, the liquid and gas phases become indistinguishable (supercritical fluid region).

• Figure references (descriptive):

  • On a PvT diagram, water shows the solid (ice), liquid (water), and gas (water vapor) regions with the triple point near the left-lower corner and the critical point far to the right/up.

Example Phase Scenarios (Practice Questions)

• Example 1: -10 °C and 50 kPa

  • Determine whether the substance would be solid, liquid, or gas using the phase diagram (practice problem from the slide set).

• Example 2: 25 °C and 90 kPa

  • Determine phase state and possible phase transitions nearby.

• Example 3: 50 °C and 40 kPa

  • Determine phase and note any expected phase change if conditions move along a boundary.

• Example 4: 80 °C and 5 kPa

  • Likely gas, given low pressure and high temperature for many substances; verify against a substance’s PvT data.

• Example 5: -10 °C and 0.3 kPa

  • Likely solid or deposition boundary depending on the substance; consult phase data.

• These scenarios illustrate how phase state depends on both temperature and pressure and why phase diagrams are useful tools in predicting matter behavior.

Phase Changes in Real Life: Applications and Implications

• Food preservation: controlling phase states (freezing, dehydration via sublimation) to extend shelf life and maintain quality.
• Climate systems: phase changes of water (evaporation, condensation, freezing, melting) regulate weather, humidity, and energy transfer in Earth's climate.
• Industrial processes: phase transitions are exploited in distillation, crystal growth, freeze-drying, and materials synthesis.
• Thermodynamics and energy efficiency: understanding phase changes leads to better thermal management and energy storage strategies.

Energy Storage and Environment

• Battery and energy storage chemistry: phase behavior can influence the performance and safety of energy storage materials.
• Energy storage concepts: thermal storage materials store and release heat for climate control, industrial processes, and renewable energy integration.
• Environmental considerations: optimizing phase changes can reduce energy consumption, minimize emissions, and improve sustainability in industrial settings.

Final Reflection: How Chemistry Knowledge Helps Care for Others and the Environment

• Apply phase-change concepts to improve food security (preservation, safety, quality).
• Enhance climate resilience and understanding of environmental systems by modeling phase changes in water and atmospheric gases.
• Design safer, more efficient industrial processes that minimize waste and energy usage.
• Consider ethical and practical implications of energy storage technologies on access to clean energy and environmental justice.

Quick References and Key Terms

• Phase changes: melting, freezing, vaporization, condensation, sublimation, deposition.
• Phase diagram: boundaries between solid, liquid, gas; triple point; critical point; supercritical fluid region.
• Sublimation: solid → gas.
• Deposition: gas → solid.
• Evaporation: liquid → gas at temperatures below boiling point (surface phenomenon).
• Boiling: liquid → gas throughout the liquid at sufficient temperature/pressure.
• Freezing point and boiling/freezing relations depend on pressure and impurities.
• Critical point coordinates: Tc, Pc; Triple point coordinates: Tt, Pt.