CHEM 210 Notes — Dispersion Forces (Part II) and Phase Transitions

1. Three Main Classes of IMFs

• Intermolecular forces (IMFs) are the forces between molecules, excluding covalent and ionic bonds.
• Major categories:
  • Dipole–dipole and charge–dipole interactions: arise from permanent dipoles (partial charges) in molecules.
  • Hydrogen bonds: a special, stronger case of dipole–dipole interactions involving large partial charges and small distances.
  • Dispersion forces: universal forces between all chemical species (molecules, ions, atoms) arising from instantaneous and induced dipoles.
• Core idea: polarity underpins IMFs; dispersion can exist even in nonpolar species.

2. Dispersion Forces — Mechanism and Key Concepts

• Dispersion forces exist in all atoms/molecules, including nonpolar ones, and even noble gases can condense due to these forces.
• How they arise:
  • Instantaneous dipoles: electron motion creates a temporary unequal distribution of charge, giving a fleeting b4+ b4-.
  • Induced dipoles: this instantaneous dipole can induce a dipole in a neighboring atom/molecule, leading to an attractive interaction during brief encounters.
• Important nuance:
  • Without instantaneous dipoles, there would be no dipole–induced dipole attraction.
  • Electron–electron repulsion in collisions is offset by the transient attractive phase when instantaneous dipoles align favorably.
• Textual representation:
  • Instantaneous dipole:
  • Symmetric, nonpolar → no permanent dipole, but transient dipole can occur.
  • Encounter of instantaneous dipoles can produce an attractive force between atoms/molecules.
• General statement: dispersion forces are present in all species but vary in strength.

3. Ne Atom as a Case Study — Why dispersion forces arise

• Ne atom configuration: $1s^22s^22p^6$
  • Fully filled valence shell; symmetric and nonpolar; no bonds.
• Question posed: What would have to happen for two Ne atoms to experience attractive interaction?
  • A process must create a non-symmetric, polar atom (or momentarily polarize both atoms) so that instantaneous dipoles can form and interact.
• Resulting picture:
  • Instantaneous dipoles can arise from electron motion.
  • When two Ne atoms collide, instantaneous dipoles (b4+ and b4−) can induce interactions that are attractive for a brief moment.

4. Induced Dipoles — How Two Ne Atoms Interact

• Two Ne atoms, without an instantaneous dipole, have no permanent dipole and would not attract via dipole–dipole forces.
• With an instantaneous dipole on one atom, the neighboring atom develops an induced dipole in response.
• In collisions, electron–electron repulsion is present, but the induced-dipole interaction can produce an overall attractive force during the encounter.
• Takeaway: dispersion forces arise from instantaneous dipoles and their induction of dipoles in neighbors; they can produce attraction even in symmetric, nonpolar species.

5. Instantaneous Dipoles, Induced Dipoles, and Repulsion/ Attraction Balance

• When a collision occurs, two atoms may have instantaneous dipoles of opposite orientations that yield an attractive interaction:
  $ext{δ}^+ ext{δ}^- o ext{attractive interaction}.$
• If only instantaneous dipoles are present without polarization, repulsion can offset attraction, but dispersion can still produce net attraction under suitable configurations.
• Summary: instantaneous dipoles enable short-lived polar interactions in otherwise nonpolar species.

6. Factors that Control Dispersion Strength

• Two main factors:
  • Polarizability: how easily the electron cloud can be distorted by external fields.
  • Contact surface area: how much surface is in contact between two particles.
• Polarizability details:
  • Describes how much a charge interaction can push electrons around in an atom or molecule.
  • Increases with larger atomic size; decreases with higher nuclear charge.
• Surface area: larger contact area yields more possible dispersion contacts, strengthening the interaction.
• Practical implication: stronger dispersion forces arise from greater polarizability and larger surface area; weaker dispersion forces arise from smaller size and less polarizable electron clouds.
• Qualitative trend: if polarizability increases, dispersion forces strengthen; larger surface area amplifies contacts and strengthens dispersion.

7. Polarizability Trend Across Diatomic Halogens

• Clace-to-face comparison: Cl, Br, I
  • Polarizability trend: Cl$2$ < Br$2$ < I$_2$
  • Consequently, dispersion forces are strongest in I$2$, weakest in Cl$2$ at STP (25°C, 1 atm).
• This trend helps explain why halogen diatomics exist in different phases at STP due to dispersion strength.

8. Surface Contact and Polarizability — Why Some Substances Have Stronger Dispersion

• Dispersion forces depend on contact between atoms/molecules.
• For substances with similar polarizability, larger surface area leads to stronger dispersion due to more contact points.
• Key takeaway: both polarizability and surface area govern dispersion strength; thus, heavier, more easily distorted electron clouds with larger contact areas show stronger dispersion.

9. Identifying IMFs in Pure Liquids — Examples

• Exercise: identify IMFs in liquids:
  • (a) CCl$_4$(l): Dispersion only (nonpolar, symmetric)
  • (b) CH$_3$COOH(l): Dispersion + Dipole–Dipole + H-bonding (carboxylic acid has strong H-bonding)
  • (c) CH$3$COCH$3$(l): Dispersion + Dipole–Dipole (no H-bonding donor/acceptor sites)
  • (d) H$_2$S(l): Dispersion + Dipole–Dipole; H-bonding is weak if present at all (S less capable of H-bonding than O)
• Note: In other slides, you may see nuanced distinctions about H-bond donor/acceptor roles; here, hot spots indicate H-bonding would be weak for H$2$S compared to H$2$O.

10. Ranking IMFs (General Guidance)

• Ranking of IMFs by overall strength (considering a given molecule pair):
  • H-bonding generally strongest when present (dominant in phase behavior)
  • Dipole–dipole interactions are medium strength
  • Dispersion forces are weak individually but always present; when molecules are similar in size/shape/composition, dispersion contributes notably
• Practical rule: compare all IMFs present; if H-bonding exists, it often determines the phase behavior; otherwise, dipole–dipole and dispersion compete.

11. Dispersion Forces and Phase Transitions — Overview

• All atoms and molecules experience dispersion forces, but their strengths differ due to polarizability and surface area.
• Phase transitions involve changes in the number of IMFs, not changes in chemical bonding or molecular structure.
• Phases (gas, liquid, solid) and transitions between them are driven by temperature and enthalpy changes.

12. Phase Transitions — Conceptual Framework

• Phases: gases, liquids, and solids; transitions occur when substances move from one phase to another.
• Key features:
  • The molecular (or atomic) structure does not change during phase transitions.
  • The number of IMFs changes as the phase changes.
• How transitions are induced: add heat (enthalpy, H).
• Temperature increase reflects rising average kinetic energy, enabling molecules to overcome IMFs and change phase.

13. Enthalpy, Enthalpy of Formation, and Phase Change Diagrams

• Enthalpy is denoted H and is a state function: it depends only on the current state, not on how the system arrived there.
• Enthalpy changes (ΔH) correspond to differences between final and initial states.
• Phase diagrams and enthalpy diagrams illustrate phase transitions for H$_2$O: solid, liquid, gas variations with heat input.
• Average kinetic energy per mole in an ideal gas:
  $KE= frac{3}{2}RT$
  (per molecule: $KE= frac{3}{2}k_B T$; per mole: $KE= frac{3}{2}RT$)

14. Enthalpies for Water: Formation and Vaporization

• Standard enthalpies of formation (H°f):
  • for H$2$O(g): ext{H}2 ext{O(g)}: \n \Delta H_f^\u209a = -241.8 ext{ kJ/mol}
  • for H$2$O(l): ext{H}2 ext{O(l)}: \Delta H_f^\u209a = -285.8 ext{ kJ/mol}
  • for elements H$2$(g) and O$2$(g): ext{Δ}H_f^\u209a = 0 ext{ kJ/mol}
• Vaporization enthalpy of water:
  • Direct difference: ext{Δ}H{ ext{vap}}^\u209a = ext{Δ}Hf^\u209a( ext{H}2 ext{O(g)}) - ext{Δ}Hf^\u209a( ext{H}_2 ext{O(l)}) = -241.8 - (-285.8) = +44.0 ext{ kJ/mol}
  • This is often cited as the standard enthalpy of vaporization: ext{Δ}H_{ ext{vap}}^\u209a = +44.0 ext{ kJ/mol}
• Hess’s Law (Ch. 9): The enthalpy change of a process that can be written as the sum of stepwise processes equals the sum of their enthalpy changes:
  • For the formation of H$2$O(g) from H$2$O(l) via the steps:
  • Step 1: H$2$O(l) → H$2$(g) + 1/2 O$2$(g): ext{Δ}H^\u209a{ ext{rxn}} = +285.8 ext{ kJ/mol}
  • Step 2: H$2$(g) + 1/2 O$2$(g) → H$2$O(g): ext{Δ}H^\u209a{ ext{rxn}} = -241.8 ext{ kJ/mol}
  • Sum: ext{Δ}H^\u209a_{ ext{vap}} = +285.8 + (-241.8) = +44.0 ext{ kJ/mol}
• Equivalent direct statement: \Delta H°_vap from Hess’s Law also yields the same +44.0 kJ/mol when combining the formation and vaporization steps.
• Key conceptual takeaways:
  • Enthalpy is a state function; the path taken to reach the final state does not alter ΔH.
  • Phase transitions (solid → liquid, liquid → gas) are endothermic for vaporization and fusion, and exothermic for condensation and freezing; thus their ΔH values are positive for endothermic processes and negative for exothermic processes.

15. Looking Ahead

• Next topics include:
  • Vapor pressure, the Clausius–Clapeyron equation, and how ΔH_vap relates to temperature dependence of vapor pressure (Ch. 10.3).
  • A brief introduction to solid-state concepts (Ch. 10.5), phase diagrams (Ch. 10.4), and heating/cooling curves (Ch. 10.3).
  • These topics are presented in the order above for coherence.