Phase & Intermolecular Forces Flashcards
(16 cards)
List the six fundamental phase transitions in order of increasing entropy change.
Freezing < Condensation < Deposition < Fusion < Vaporization < Sublimation.
First three are exothermic; last three endothermic.
Why are intermolecular forces (IMFs) central to phase-change energetics?
Energy is required to overcome or released upon forming IMFs; stronger IMFs correspond to larger enthalpy changes.
Explain why ‘breaking bonds’ is an improper phrase for melting ice.
Melting disrupts hydrogen-bond-based intermolecular attractions, not covalent O–H intramolecular bonds.
How does phase favorability change with temperature in terms of Gibbs free energy?
ΔG phase shifts sign when TΔS equals ΔH; at this point, the two phases are in equilibrium (e.g., melting point).
Identify the three intermolecular force types emphasized in the deck.
London dispersion forces, dipole–dipole interactions, and hydrogen bonding.
What molecular property principally determines dipole–dipole strength?
The magnitude of the permanent molecular dipole moment, which is governed by bond polarity and molecular geometry.
Specify the requisite atomic attachments for hydrogen bonding.
Hydrogen must be directly bonded to fluorine, oxygen, or nitrogen.
Describe how molecular size influences London dispersion forces.
Larger electron clouds are more polarizable, yielding stronger instantaneous dipoles and thus stronger dispersion forces.
Relate polarity to boiling-point trends among homologous series.
More polar molecules (or those capable of hydrogen bonding) exhibit higher boiling points due to stronger intermolecular attractions that demand greater energy input to vaporize.
Why do highly polar molecules release more energy upon condensation?
They establish stronger intermolecular attractions, lowering the potential energy substantially and liberating heat to the surroundings.
Distinguish between intermolecular and intramolecular forces in terms of energy scale.
Intramolecular (bonding) energies are typically one to two orders of magnitude larger than intermolecular energies.
How do IMFs influence solubility patterns (‘like dissolves like’)?
Solvation is favored when solvent–solute interactions are comparable in strength to solvent–solvent and solute–solute interactions, minimizing the free-energy penalty for mixing.
Classify solid↔liquid↔gas transitions of H2O as physical or chemical changes.
They are physical changes because molecular identity and covalent bonding remain intact.
Define temperature in kinetic-molecular terms.
Temperature is the proportional measure of the average translational kinetic energy of molecules.
Identify which water-phase conversions involve temperature change versus potential-energy change.
Heating or cooling within a single phase alters temperature (kinetic energy); phase transitions at constant temperature alter potential energy (intermolecular).
Explain why vaporization occurs at a constant temperature for a pure substance at fixed pressure.
The heat supplied is consumed entirely as latent heat to overcome IMFs; thus kinetic energy (temperature) remains steady until the phase transition is complete.
