Condensation is the phase change in which a gas becomes a liquid, releasing energy equal in magnitude to the enthalpy of vaporization (q = -nΔHvap). On the AP Chem exam it shows up in thermochemistry (6.5), vapor-liquid equilibrium (7.1), and deviations from ideal gas behavior (3.6).
Condensation is the gas-to-liquid phase transition. When gas particles slow down enough (lower temperature) or get pushed close enough together (higher pressure), their intermolecular attractions win, and they stick together as a liquid. Because the particles are forming new attractions, the system releases energy. That makes condensation exothermic from the system's perspective, with q = -nΔHvap (EK 6.5.A.1 and 6.5.A.2). The energy released when one mole of a gas condenses is exactly the same magnitude as the energy absorbed when one mole of that liquid vaporizes. Just flip the sign.
Condensation is also one of the CED's go-to examples of a reversible physical process (EK 7.1.A.1). In a closed container, water evaporates and condenses at the same time. When the rates of those two opposing processes become equal, the system reaches dynamic equilibrium. Nothing looks like it's changing, but molecules are constantly trading places between the liquid and the gas. One more place it matters is real gas behavior. Near the conditions where a gas is about to condense, intermolecular attractions become significant and the gas stops behaving ideally (EK 3.6.A).
Condensation is rare in that it directly anchors essential knowledge in three different units. In Unit 6, LO 6.5.A asks you to calculate the heat released during a phase transition using q = nΔH, and condensation is the classic 'don't forget the negative sign' case. In Unit 7, LO 7.1.A uses evaporation and condensation as the textbook example of a reversible process that establishes dynamic equilibrium, which is the conceptual foundation for everything else in Units 7 and 8. In Unit 3, LO 3.6.A says gases deviate most from ideal behavior at conditions close to condensation, because that's exactly when attractive forces between particles matter most. If you understand why a gas condenses, you understand why PV = nRT breaks down.
Keep studying AP Chemistry Unit 3
Vaporization (Unit 6)
Vaporization and condensation are the same process run in opposite directions. The energy absorbed to vaporize a sample equals the energy released when it condenses, so you only need one ΔHvap value. Just watch the sign.
Chemical Equilibrium (Unit 7)
A sealed container of water at constant temperature reaches vapor-liquid equilibrium when the rate of condensation equals the rate of evaporation. This is the CED's favorite picture of dynamic equilibrium, and it works for physical changes exactly the way it works for chemical reactions.
Deviation from Ideal Gas Law (Unit 3)
A real gas behaves least ideally right before it condenses. Low temperature and high pressure are exactly the conditions where attractive forces pull particles together, which is the 'a' correction in the van der Waals equation.
Enthalpy Change (Unit 6)
Condensation has a negative ΔH because the system loses energy as particles form intermolecular attractions. Meanwhile the temperature of the pure substance stays constant during the entire phase change, a detail MCQs love to test.
Condensation usually shows up in calculations and conceptual MCQs rather than as a vocabulary word. Expect q = nΔH problems where a sample condenses and then freezes, so you add the heat released in both steps (using ΔHvaporization and ΔHfusion) and keep the signs straight. Calorimetry questions are another favorite, like steam bubbled into cooler water, where you have to recognize that the condensation of the steam (not the cooling of the resulting hot water) releases the most energy, because ΔHvap is large. In Unit 3, MCQs connect condensation to the van der Waals 'a' constant and ask why real gases deviate from ideal behavior near condensing conditions. In Unit 7, you might have to explain why a closed flask of liquid and vapor at constant pressure is at dynamic equilibrium even though nothing appears to change. No released FRQ has hinged on the word 'condensation' itself, but phase change energetics and particle-level explanations of equilibrium are standard FRQ territory.
They're opposite phase changes. Vaporization is liquid to gas and absorbs energy (endothermic, +ΔHvap); condensation is gas to liquid and releases energy (exothermic, -ΔHvap). The magnitude of the enthalpy change is identical for both. Students most often lose points by using a positive q for condensation. If a gas is condensing, the system is releasing heat, so q for the system is negative.
Condensation is the gas-to-liquid phase change, and the system releases energy because particles form intermolecular attractions.
The heat released during condensation is q = -nΔHvap, the same magnitude as vaporization but with the opposite sign.
The temperature of a pure substance stays constant while it condenses, since the released energy comes from forming attractions, not from cooling.
In a closed system, condensation and evaporation occur simultaneously, and dynamic equilibrium is reached when their rates are equal.
Real gases deviate most from ideal behavior near the conditions where they condense, because attractive forces between particles become significant.
In calorimetry problems with steam, the condensation step usually releases far more energy than the cooling step, since ΔHvap is large.
Condensation is the phase change from gas to liquid. It's exothermic, releasing heat equal to nΔHvap, and it's one of the CED's main examples of a reversible process that can reach dynamic equilibrium (EK 7.1.A.1).
Exothermic. The system releases energy as gas particles slow down and form intermolecular attractions, so q for the system is negative. Vaporization, the reverse process, is endothermic with the same magnitude of ΔH.
No. For a pure substance, temperature stays constant during the entire phase change. The energy released comes from forming intermolecular attractions, not from a drop in particle kinetic energy. This is why heating and cooling curves have flat plateaus.
They're exact opposites. Vaporization turns liquid into gas and absorbs +ΔHvap per mole, while condensation turns gas into liquid and releases that same amount of energy. In a closed container at equilibrium, both happen at equal rates simultaneously.
Near condensing conditions (low temperature, high pressure), attractive forces between gas particles become significant, which the ideal gas law ignores. The van der Waals equation corrects for this with the 'a' constant, which accounts for interparticle attractions.
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