Chemistry – Chapter 7 Notes: Phase Changes, Energetics & Intro Gas Concepts

Chapter Scope and Learning Goals

• Chapter 7 focuses on:
  • Changes of state (phase transitions) in matter.
  • Energy transfer accompanying each transition.
  • Properties of gases and an introduction to the gas laws (to be covered later).
  • Gas mixtures and Dalton’s law of partial pressures (preview only).
• Key skills to master:
  • Identify and classify six fundamental phase changes.
  • Relate kinetic energy (KE), intermolecular forces (IMF), and temperature to the physical state of a substance.
  • Distinguish endothermic vs. exothermic processes.
  • Use thermodynamic quantities (heat of fusion, heat of vaporization, specific heat) as conversion factors.
  • Interpret and sketch heating curves and cooling curves.

Three States of Matter – Quick Review

• Solid
  • Particles tightly packed, fixed positions, vibrational motion only.
  • Highest IMF, lowest KE.
• Liquid
  • Particles close but able to slide/flow.
  • Intermediate IMF and KE.
• Gas
  • Particles far apart, move independently.
  • Lowest IMF, highest KE.

Six Fundamental Phase Changes

• Melting (Fusion)
  • Solid → Liquid.
  • Example: Ice cube melts on a hot day.
• Freezing (Solidification)
  • Liquid → Solid.
  • Example: Water freezing in a freezer.
• Vaporization / Evaporation / Boiling
  • Liquid → Gas.
  • Example: Puddle water evaporates in summer.
• Condensation
  • Gas → Liquid.
  • Example: Water vapor condenses on a cold glass.​
• Sublimation
  • Solid → Gas (skips liquid phase).
  • Example: Dry ice (solid CO₂) turning into vapor cloud.
• Deposition
  • Gas → Solid (skips liquid phase).
  • Example: Frost or snowfall (water vapor → ice crystals).

Molecular-Level Explanation of Phase Changes

• Every particle (atom, molecule, ion) possesses kinetic energy $\bigl(KE\,=\,\tfrac12 m v^2\bigr)$ that increases with temperature.
• Simultaneously, particles attract one another via intermolecular forces (potential energy, PE).
• Physical state is a balance:
  • Low KE + strong IMF ⇒ solid.
  • Moderate KE + moderate IMF ⇒ liquid.
  • High KE + weak IMF ⇒ gas.
• During a phase change
  • Temperature remains constant because energy is used to disrupt (or form) IMF, not to increase KE.
  • Covalent (intramolecular) bonds are NOT broken → phase change is a physical, not chemical, transformation.

Types of Intermolecular Forces (IMF)

• Dispersion (London) Forces
  • Present in all substances; only force in non-polar molecules.
  • Weakest; strength increases with molar mass/size.
• Dipole–Dipole Forces
  • Require permanent molecular dipole (polar molecules).
  • Intermediate strength.
• Hydrogen Bonding (H-bonding)
  • Special strong dipole force when H is directly bonded to N, O, or F.
  • Strongest of the three.

IMF vs. Physical Properties – Examples

• Ethanol (polar, O–H present)
  • Exhibits hydrogen bonding → stronger IMF.
  • Liquid at room temperature.
• Propane (non-polar, only dispersion)
  • Much weaker IMF.
  • Gas at room temperature.
• General trend:
  • Stronger IMF ⇒ higher melting point (m.p.) and boiling point (b.p.).
  • Non-polar gases (e.g., N₂, O₂) have very low m.p./b.p.; polar substances (e.g., H₂O, NH₃) have higher values.

Energy Classification of Phase Changes

• Endothermic (heat absorbed from surroundings)
  • Melting, Vaporization, Sublimation.
  • \text{q} > 0 for the system; surroundings cool down.
• Exothermic (heat released to surroundings)
  • Freezing, Condensation, Deposition.
  • \text{q} < 0 for the system; surroundings warm up.

Compressibility of Gases

• Gases are highly compressible.
• Applying external pressure can force gas particles closer, enabling condensation (gas → liquid) without cooling.
• Practical relevance: liquefied petroleum gas (LPG), aerosol cans, refrigeration cycles.

Characteristic Temperatures (at 1 atm)

• Melting Point (m.p.)
  • Temperature where solid ↔ liquid equilibrium exists.
• Boiling Point (b.p.)
  • Temperature where vapor pressure equals external pressure (1 atm ⇒ normal b.p.).
• Both are physical (not chemical) properties; useful for identification and purity checks.

Heating Curve – Water Example

• Axes: heat added (x) vs. temperature (y).
• Regions:
  • A (Solid Ice): $-30\,^{\circ}\text{C} \to 0\,^{\circ}\text{C}$, temp rises.
  • B (Melting Plateau): $0\,^{\circ}\text{C}$, temp constant, ice → water.
  • C (Liquid Water): $0\,^{\circ}\text{C} \to 100\,^{\circ}\text{C}$, temp rises.
  • D (Boiling Plateau): $100\,^{\circ}\text{C}$, temp constant, water → steam.
  • E (Superheated Steam): >$100\,^{\circ}\text{C}$, temp rises again.
• Key insight: Horizontal segments = phase changes (temperature constant while heat disrupts or forms IMF).

Thermodynamic Quantities for Phase Change

• Heat of Fusion ($\Delta H_{fus}$)
  • Energy required to melt 1 g (or 1 mol) at the m.p.
  • For H₂O: $\Delta H_{fus}=80\,\text{cal g}^{-1}$ (common reference value).
• Heat of Vaporization ($\Delta H_{vap}$)
  • Energy required to vaporize 1 g (or 1 mol) at the b.p.
  • For H₂O: $\Delta H_{vap}=539\,\text{cal g}^{-1}$.
• $\Delta H<em>{vap} > \Delta H</em>{fus}$ because all IMF must be overcome to enter gas phase.

Using $\Delta H<em>{fus}$ and $\Delta H</em>{vap}$ as Conversion Factors

• Format: $\Delta H_{x}\bigl(=\frac{\text{cal}}{\text{g}}\bigr)$ or $\frac{\text{kJ}}{\text{mol}}$.
• Example Problem: Heat to vaporize 10 g H₂O at $100\,^{\circ}\text{C}$
  • $q = 10\,\text{g} \times 539\,\tfrac{\text{cal}}{\text{g}} = 5.39\times10^{3}\,\text{cal}$ (5390 cal).

Temperature Change (No Phase Change) – Specific Heat

• Specific Heat ($c$): heat needed to raise 1 g of substance by $1\,^{\circ}\text{C}$.
• Common units: $\text{cal g}^{-1}\,^{\circ}\text{C}^{-1}$ or $\text{J g}^{-1}\,^{\circ}\text{C}^{-1}$.
• Heat equation ("q-equation"): $q = m\,c\,\Delta T$ where:
  • $q$ = heat (cal or J).
  • $m$ = mass (g).
  • $c$ = specific heat.
  • $\Delta T = T<em>{final}-T</em>{initial}$ $\,(^{\circ}\text{C})$.

Worked Examples

1. Paraffin Wax
   • $m = 10\,\text{g},\; c = 0.60\,\tfrac{\text{cal}}{\text{g}\,^{\circ}\text{C}},\; \Delta T = 30-20 = 10\,^{\circ}\text{C}$
   • $q = 10 \times 0.60 \times 10 = 60\,\text{cal}$.
2. Liquid Water
   • $m = 45\,\text{g},\; c = 1.00\,\tfrac{\text{cal}}{\text{g}\,^{\circ}\text{C}},\; \Delta T = 32-23 = 9\,^{\circ}\text{C}$
   • $q = 45 \times 1.00 \times 9 = 405\,\text{cal}$.

Typical Conceptual & Multiple-Choice Insights

• Greatest potential energy (strongest IMF) exists in the solid state.
• During vaporization, IMF are disrupted; covalent bonds remain intact.
• Endothermic vs. exothermic:
  • Deposition, condensation, freezing are all exothermic.
• Correct identification of $\Delta H_{vap}$:
  • It is the heat added when a substance evaporates at its boiling point (not when it melts, condenses, or sublimes).
• Heating curve recognition:
  • Phase changes occur where temperature line is flat (constant).

Ethical & Real-World Relevance

• Understanding phase change energetics underpins refrigeration, climate science (snowfall vs. rain), medicinal cryopreservation, and energy budgeting in chemical engineering.
• Comprehending gas compressibility is critical for safe cylinder storage and aerosol technology.

Key Equations & Numbers to Memorize

• Heat equation (temperature change): $q = m\,c\,\Delta T$.
• Energy for melting: $q = m\,\Delta H_{fus}$.
• Energy for vaporization: $q = m\,\Delta H_{vap}$.
• Typical water values (1 atm):
  • $m.p. = 0\,^{\circ}\text{C},\; b.p. = 100\,^{\circ}\text{C}$.
  • $c_{(l)} = 1.00\,\tfrac{\text{cal}}{\text{g}\,^{\circ}\text{C}} \;\bigl(4.184\,\tfrac{\text{J}}{\text{g}\,^{\circ}\text{C}}\bigr)$.
  • $\Delta H<em>{fus} = 80\,\text{cal g}^{-1},\; \Delta H</em>{vap} = 539\,\text{cal g}^{-1}$.

Study Tips & Connections

• Always specify the phase and temperature before applying $q = m\,\Delta H_{x}$ or $q = m c \Delta T$.
• Combine steps for multi-stage problems: cooling, freezing, further cooling ⇒ sum of individual $q$ values.
• Link back to Chapter 1 (matter & IMF) and upcoming gas-law chapters (relationship between P, V, T for gases and Dalton’s law of partial pressures).
• Practice drawing heating/cooling curves for different substances, labeling each plateau with correct $\Delta H$ value.