Phase Transitions and Intermolecular Forces

• Phase transitions and phase diagrams are being examined
• Homework is HW26/28 – Phase Diagrams and Heating/Cooling Curves
• Lab 4 Paper Chromatography is scheduled for Thursday, March 13

Intermolecular Forces and Vapor Pressure

• Vapor pressure decreases as intermolecular forces increase
• Stronger IMFs reduce the tendency for a liquid to evaporate, causing slower evaporation and a higher boiling point

Boiling vs Evaporation

• At the boiling point, a liquid's vapor pressure equals atmospheric pressure
• Evaporation occurs when the vapor pressure of a liquid is less than the atmospheric pressure
• Evaporation refers to vaporization at temperatures below the boiling point

Intermolecular Forces and Phase Changes

• Strong intermolecular forces, such as hydrogen bonds in water, hold liquid molecules tightly, hindering their escape into the gas phase

Evaporation

• Liquids with lower vapor pressure evaporate more slowly
• Fewer molecules have enough energy to overcome strong intermolecular attractions and escape into the gas phase

Boiling Point

• Liquids with low vapor pressure need more energy and also a higher temperature to reach the boiling point

Examples

• Water has strong hydrogen bonds, exhibiting a low vapor pressure and slow evaporation
• Diethyl ether has weaker forces, resulting in a higher vapor pressure and quick evaporation

Vapor Pressure and Alcohols

• All alcohols exhibit hydrogen bonding that makes overcoming the intermolecular forces more difficult, resulting in lower vapor pressures
• As the molecule size increases from methanol to butanol, dispersion forces also increase
• The increase in molecule size causes vapor pressures to decrease
  • P methanol > P ethanol > P propanol > P butanol

Vapor Pressure and Temperature

• The vapor pressure of a liquid increases with an increase in temperature
• For example, water has a vapor pressure of 24 mm Hg at 25 °C and 92 mm Hg at 50 °C

Boiling Points

• The boiling point is reached when the vapor pressure of a liquid equals the external pressure
• A liquid's boiling temperature depends on the pressure above it
• The normal boiling point occurs when the pressure above the liquid is 1 atm
• Temperature remains constant during the boiling process

Boiling Point and External Pressure

• Variations in atmospheric pressure will change the boiling point
• At higher altitudes, where atmospheric pressure is lower, the boiling point is also lower

Enthalpy of Vaporization

• Vaporization absorbs energy
• The reverse of an endothermic process is exothermic, therefore condensation releases energy
  • H2O(l) → H2O(g) : ΔΗvap = 44.01 kJ/mol
  • H₂O(g) → H2O(l) : ΔΗcon = -ΔΗvap = -44.01kJ/mol

Enthalpy of Fusion

• Fusion(melting) absorbs energy
  • H2O(s) → H2O(1) : ΔΗfus = 6.01 kJ/mol
• The reciprocal process of freezing releases energy
  • H2O(l)→H2O(s) : ΔΗfrz = -ΔΗfus = -6.01kJ/mol

Enthalpy of Sublimation

• Like vaporization, sublimation requires energy input
  • CO₂(s) → CO2(g) : ΔΗsub = 26.1 kJ/mol
• Deposition releases energy
  • CO2(g) → CO₂(s) : ΔΗdep Hdep = -ΔΗsub = -26.1kJ/mol
• Converting a solid into a liquid requires that these attractions be only partially overcome
• Transitioning to the gaseous state requires that they be completely overcome
• Enthalpy of fusion is less than its enthalpy of vaporization

Heating and Cooling Curves

• The chapter on thermochemistry introduced the equation q=mc∆T
  • "m" is the mass of the substance
  • "c" is its specific heat
  • "q" is the amount of heat absorbed or released by a substance
  • "∆T" is the change in temperature
• It applies to matter being heated or cooled, but not undergoing a change in state
• Further gain or loss of heat results in diminishing or enhancing intermolecular attractions instead of increasing or decreasing molecular kinetic energies
• A substance's temperature remains constant while undergoing a change in state

Heating Curve Key Points

• A heating curve is divided into solid, melting (solid-liquid transition), liquid, boiling (liquid-gas transition), and gas phases
• During the solid and liquid phases, the temperature increases as heat is added
• Temperature remains constant during phase changes at the melting and boiling points

Phase Changes and Temperature

• Phase changes are when the temperature remains constant
• The intermolecular forces are broken (for energy being absorbed)
• Intermolecular forces are formed (for energy being released)

Phase Diagrams

• A phase diagram graphically represents the physical state of a substance and its phase-transition temperatures
• A phase diagram shows solid phase favored at low temperature and high pressure and the gas phase favored at high temperature and low pressure

Phase Diagram Key Points

• The x-axis represents temperature, and the y-axis represents pressure
• The lines indicate the equilibrium points between different phases (solid-liquid, liquid-gas, solid-gas)
• All three phases (solid, liquid, gas) coexist in equilibrium at the Triple Point
• A supercritical fluid exists beyond the Critical Point

Understanding Critical Temperature

• All substances have a critical temperature above which the gas can no longer be liquefied, regardless of pressure
• Critical temperature (Tc) is the highest temperature at which a substance can exist as a liquid
• Molecules have too much kinetic energy to be held together in a separate liquid phase if above critical temperature