Thermodynamics Quiz: Phases And Phase Changes
20 questions · exam conditions
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Phases And Phase ChangesQuestion 1 of 20

Steam initially at 300°C and 1 MPa is cooled at constant pressure until it reaches 150°C. During this cooling process, at what temperature will the first liquid droplet appear, and what is the significance of this point?

At 179.9°C when steam reaches saturated vapor state, marking the onset of condensation
At 150°C when the final temperature is reached, since cooling always ends in liquid phase
At 200°C as an intermediate point where phase change typically begins during cooling
Liquid will not appear since 150°C is above atmospheric boiling point temperature
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Thermodynamics Quiz

Thermodynamics Quiz: Phases And Phase Changes

Practice Phases And Phase Changes in Thermodynamics with focused quiz questions that help you check what you know, review explanations, and build confidence with test-style prompts.

What this quiz covers

This quiz focuses on Phases And Phase Changes, giving you a quick way to practice the rules, question types, and explanations that matter most for Thermodynamics.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

All questions

Question 1

Steam initially at 300°C and 1 MPa is cooled at constant pressure until it reaches 150°C. During this cooling process, at what temperature will the first liquid droplet appear, and what is the significance of this point?

  1. At 179.9°C when steam reaches saturated vapor state, marking the onset of condensation (correct answer)
  2. At 150°C when the final temperature is reached, since cooling always ends in liquid phase
  3. At 200°C as an intermediate point where phase change typically begins during cooling
  4. Liquid will not appear since 150°C is above atmospheric boiling point temperature
Explanation: At 1 MPa pressure, the saturation temperature is 179.9°C. Steam initially at 300°C is superheated. During constant pressure cooling, it first reaches the saturated vapor state at 179.9°C, which is where the first liquid droplet appears (dew point). This marks the beginning of the condensation process. Choice B confuses the final temperature with the condensation start point. Choice C uses an arbitrary temperature. Choice D incorrectly applies atmospheric pressure logic to a 1 MPa system.

Question 2

Which statement correctly explains why the critical point represents a unique state for a pure substance?

  1. It's the highest temperature at which the liquid phase can exist
  2. It's the point where all three phases coexist simultaneously
  3. It's where the distinction between liquid and vapor phases disappears (correct answer)
  4. It's the lowest pressure at which the vapor phase can exist
  5. It's where the substance has maximum density in any phase
Explanation: When analyzing phase behavior, the critical point represents one of the most fascinating phenomena in thermodynamics—a state where the boundary between liquid and vapor phases completely vanishes. Understanding this concept requires visualizing what happens as you approach critical conditions. At the critical point, liquid and vapor phases become indistinguishable because their densities converge to the same value. Below the critical temperature, you can clearly see the difference between liquid (high density) and vapor (low density) phases. But as temperature and pressure increase toward critical conditions, the density difference shrinks until it disappears entirely. At this unique state, properties like surface tension become zero, and the meniscus between phases vanishes. This is why choice C correctly captures the essence of the critical point. Let's examine why the other options miss the mark. Choice A incorrectly suggests the critical point is simply about maximum temperature for liquid existence, but this ignores the crucial aspect of phase distinction disappearing. Choice B confuses the critical point with the triple point—the triple point is where solid, liquid, and vapor coexist simultaneously, occurring at much lower conditions. Choice D mischaracterizes the critical point as defining minimum vapor pressure conditions, which isn't what makes this state special. For thermodynamics problems, remember that the critical point is fundamentally about phase boundary elimination, not just extreme conditions. When you see questions about critical phenomena, focus on the concept of phases becoming identical rather than thinking about temperature or pressure limits alone.

Question 3

Ice at 0°C0°C and 11 atm is in contact with liquid water at the same conditions. A small amount of heat is added to the system. What happens first?

  1. The ice temperature increases above 0°C0°C before any melting occurs
  2. The liquid water temperature increases above 0°C0°C before ice melts
  3. Some ice melts while both phases remain at 0°C0°C (correct answer)
  4. Both ice and liquid water temperatures increase simultaneously
  5. The liquid water begins to vaporize before ice melts
Explanation: When you encounter a phase equilibrium problem, focus on what happens at the phase transition temperature. At 0°C0°C and 11 atm, ice and liquid water exist in thermal equilibrium—both phases are at exactly the same temperature and can coexist indefinitely. When heat is added to this two-phase system, the energy doesn't immediately raise the temperature of either phase. Instead, it provides the latent heat of fusion needed to break the hydrogen bonds holding the ice crystal structure together. This energy goes directly into the phase change process, converting solid water molecules to liquid while both phases remain at 0°C0°C. Only after all the ice melts would additional heat raise the liquid water's temperature above 0°C0°C. Option A is wrong because ice cannot exceed 0°C0°C while liquid water is present at 11 atm—any energy added goes into melting instead of temperature increase. Option B fails for the same thermodynamic reason: the liquid water cannot heat above 0°C0°C while ice remains, as the system maintains thermal equilibrium at the phase transition temperature. Option D incorrectly suggests both phases can simultaneously increase in temperature, but this violates the principle that energy added during a phase change goes into breaking intermolecular bonds, not increasing kinetic energy (temperature). Remember this key insight: during any phase transition at constant pressure, temperature remains constant while the phase change occurs. Look for this pattern whenever you see problems involving melting, freezing, boiling, or condensation at equilibrium conditions.

Question 4

Water at 25°C25°C and 11 atm is slowly heated at constant pressure until it reaches 150°C150°C. During this process, which sequence of phases does the water experience?

  1. Liquid → vapor only
  2. Liquid → liquid-vapor mixture → vapor (correct answer)
  3. Liquid → solid → liquid-vapor mixture → vapor
  4. Liquid → supercritical fluid → vapor
  5. Liquid → vapor → supercritical fluid
Explanation: When analyzing phase transitions at constant pressure, you need to consider the specific conditions and the normal boiling point of the substance. Water has a boiling point of 100°C100°C at 11 atm, which falls right in the middle of the temperature range given. Starting at 25°C25°C and 11 atm, water exists as a liquid since this temperature is well below its boiling point. As you heat the water at constant pressure, it remains liquid until it reaches 100°C100°C. At this point, the water begins to boil, creating a liquid-vapor mixture where both phases coexist in equilibrium. The temperature stays constant at 100°C100°C during this phase change as energy goes into breaking intermolecular bonds rather than increasing temperature. Once all the liquid has vaporized, continued heating raises the vapor temperature from 100°C100°C to 150°C150°C. This gives us the sequence: liquid → liquid-vapor mixture → vapor. Choice A is incorrect because it ignores the coexistence region where liquid and vapor are both present during boiling. Choice C is wrong because cooling to solid phase isn't involved—we're heating throughout, and the temperatures are all above water's freezing point. Choice D is incorrect because supercritical conditions require both high temperature (above 374°C374°C) and high pressure (above 221221 atm) simultaneously, neither of which applies here. Remember: at constant pressure, any pure substance will have a temperature plateau during phase transitions where both phases coexist. Always check if the given conditions cross a phase boundary.

Question 5

A substance exists as a liquid-vapor mixture at 200°C200°C and 1515 bar. If the pressure is increased to 3030 bar while temperature remains constant, what phase change occurs?

  1. The mixture becomes entirely vapor due to pressure increase
  2. The mixture becomes entirely liquid due to pressure increase
  3. The mixture composition shifts toward more vapor content
  4. The mixture composition shifts toward more liquid content (correct answer)
  5. No phase change occurs; mixture composition remains unchanged
Explanation: When you encounter liquid-vapor mixture problems, the key principle is understanding how pressure affects phase equilibrium at constant temperature. At any given temperature, there's a specific saturation pressure where liquid and vapor can coexist in equilibrium. At 200°C200°C, the saturation pressure is approximately 15.515.5 bar. Since the initial condition (1515 bar) is very close to this saturation pressure, a liquid-vapor mixture exists. When pressure increases to 3030 bar while temperature stays constant at 200°C200°C, you're moving further away from the saturation conditions into the compressed liquid region. Higher pressure at constant temperature favors the denser phase—liquid. According to Le Chatelier's principle, the system responds to increased pressure by shifting toward the phase that occupies less volume per unit mass. Since liquid is much denser than vapor, increasing pressure drives the equilibrium toward more liquid content. Answer A incorrectly suggests vapor formation increases with pressure, which contradicts basic phase behavior. Answer B overstates the effect—while pressure favors liquid, it won't instantly convert the entire mixture to liquid at this moderate pressure increase. Answer C reverses the actual relationship, incorrectly claiming higher pressure produces more vapor. Answer D correctly identifies that higher pressure shifts the mixture composition toward more liquid content, as the system responds to pressure by favoring the denser liquid phase. Study tip: Remember that pressure and temperature have opposite effects on liquid-vapor equilibrium. Higher pressure favors liquid; higher temperature favors vapor. Always consider which phase is denser when predicting pressure effects.

Question 6

Dry ice (solid CO2CO_2) at −78°C-78°C and 11 atm is placed in a sealed container and heated to 25°C25°C. Assuming the container doesn't rupture, what phase is most likely present at equilibrium?

  1. Solid CO2CO_2 only, unchanged from initial state
  2. Liquid CO2CO_2 only, after melting from solid
  3. Gaseous CO2CO_2 only, after sublimation from solid (correct answer)
  4. Liquid-gas mixture of CO2CO_2 in equilibrium
  5. Solid-liquid mixture of CO2CO_2 in equilibrium
Explanation: When you encounter phase change problems, you need to consider both temperature and pressure conditions, along with the specific properties of the substance. For CO2CO_2, understanding its phase diagram is crucial because it behaves differently from water. Carbon dioxide has a triple point at −56.6°C-56.6°C and 5.175.17 atm. Below this pressure, CO2CO_2 cannot exist as a liquid regardless of temperature—it can only exist as solid or gas. Since your container starts at 11 atm (below the triple point pressure) and is sealed, the pressure is determined by the vapor pressure of CO2CO_2 at the final temperature. When heated from −78°C-78°C to 25°C25°C, the solid CO2CO_2 will sublime directly to gas because liquid CO2CO_2 cannot exist at 11 atm. The final pressure will be the vapor pressure of CO2CO_2 at 25°C25°C, which is much higher than 11 atm but still below what's needed for liquid formation. Answer A is wrong because heating will definitely cause a phase change—the solid won't remain unchanged. Answer B is incorrect because CO2CO_2 cannot exist as liquid at the initial pressure of 11 atm; it would need to be above 5.175.17 atm. Answer D is wrong because a liquid-gas mixture requires conditions near the triple point or higher pressures, which don't apply here. Study tip: Always check if the pressure is above or below the triple point pressure when predicting CO2CO_2 phase changes. Below 5.175.17 atm, CO2CO_2 skips the liquid phase entirely.

Question 7

A liquid at its bubble point undergoes an infinitesimal increase in temperature at constant pressure. What immediately occurs?

  1. The liquid becomes superheated above its boiling point
  2. The first vapor bubbles begin to form within the liquid (correct answer)
  3. The liquid completely vaporizes instantaneously
  4. The liquid density decreases but remains single-phase
  5. The liquid transforms to a supercritical fluid
Explanation: When you encounter questions about phase transitions, focus on the precise definitions of key thermodynamic states. The bubble point is the temperature at which a liquid is saturated and in equilibrium with its vapor at a given pressure. At the bubble point, the liquid exists at the threshold of vaporization. Any infinitesimal temperature increase while maintaining constant pressure provides just enough energy to overcome intermolecular forces and begin the phase transition. This triggers nucleation - the formation of the very first vapor bubbles within the liquid bulk. The system transitions from pure liquid to a two-phase mixture where liquid and vapor coexist. Option A is incorrect because superheating refers to heating a liquid above its boiling point without vaporization occurring, which requires special conditions like perfectly smooth containers. Here, we start exactly at the bubble point where vaporization readily occurs. Option C represents a fundamental misunderstanding of phase transitions. Complete vaporization requires significant energy input (the latent heat of vaporization) and cannot happen instantaneously from an infinitesimal temperature change. Option D incorrectly suggests the system remains single-phase. While liquid density does change with temperature, once you exceed the bubble point, vapor formation is thermodynamically inevitable, creating a two-phase system. Remember that bubble point and dew point questions test your understanding of saturation states. The bubble point marks the onset of vaporization (first bubble), while the dew point marks the onset of condensation (first droplet). These are precise thermodynamic boundaries, not gradual transitions.

Question 8

Which statement correctly describes the difference between saturated liquid and compressed liquid for a pure substance?

  1. Saturated liquid exists at higher temperature than compressed liquid at the same pressure
  2. Compressed liquid exists at higher pressure than saturated liquid at the same temperature (correct answer)
  3. Saturated liquid has lower density than compressed liquid at the same conditions
  4. Compressed liquid is always at its bubble point, while saturated liquid is not
  5. Saturated liquid can coexist with vapor, while compressed liquid cannot under any conditions
Explanation: When you encounter questions about phase states of pure substances, focus on the relationship between pressure, temperature, and the saturation curve on a T-P diagram. Compressed liquid (also called subcooled liquid) exists at conditions where the substance remains liquid even though it could potentially vaporize. This occurs when the liquid is either at a temperature below its saturation temperature for a given pressure, or at a pressure above its saturation pressure for a given temperature. The key insight is understanding what "compressed" means - the liquid is under additional pressure beyond what's needed to keep it from vaporizing. Option B correctly captures this relationship. At any given temperature, compressed liquid must exist at higher pressure than the saturated liquid. This extra pressure "compresses" or "subcools" the liquid, keeping it away from its vaporization point. Option A reverses the temperature relationship. At the same pressure, saturated liquid exists at a lower temperature than compressed liquid would need to maintain its liquid state. Option C incorrectly suggests density differences - compressed liquid actually has slightly higher density due to the additional pressure compacting the molecules. Option D completely misunderstands the definitions: saturated liquid is precisely at its bubble point (the condition where vaporization begins), while compressed liquid is away from its bubble point. Remember this key pattern: "compressed" or "subcooled" liquid is always pushed away from the saturation curve by either lower temperature or higher pressure, while "saturated" liquid sits exactly on the phase boundary where liquid and vapor coexist.

Question 9

A pure substance exists as a wet vapor (liquid-vapor mixture) with a quality of 0.80.8. What does this quality value specifically indicate?

  1. 80%80\% of the total volume is occupied by vapor phase
  2. 80%80\% of the total mass is in the vapor phase (correct answer)
  3. 80%80\% of the total internal energy is due to vapor phase
  4. The mixture is 80%80\% of the way from liquid to vapor state
  5. 80%80\% of the molecules have sufficient energy to be vapor
Explanation: When you encounter questions about quality in thermodynamics, you're dealing with the fundamental definition of how we quantify two-phase mixtures. Quality (denoted as xx) is specifically defined as the mass fraction of vapor in a liquid-vapor mixture. A quality of 0.80.8 means that 80%80\% of the total mass is in the vapor phase, while the remaining 20%20\% is liquid. This is expressed mathematically as x=mvapormtotalx = \frac{m_{vapor}}{m_{total}}. So answer B correctly captures this mass-based definition. Let's examine why the other options miss the mark. Answer A confuses quality with volume fraction - while vapor does occupy more volume than liquid at the same conditions due to density differences, quality specifically refers to mass, not volume. Answer C incorrectly associates quality with internal energy distribution. Though vapor typically has higher specific internal energy than liquid, quality doesn't directly represent energy fractions. Answer D treats quality as a general "progress indicator" from liquid to vapor state, which sounds intuitive but misses the precise technical definition that quality is always about mass fraction. Remember that quality is always a mass-based property in thermodynamics. When you see quality values, think "mass fraction of vapor." This distinction is crucial because volume fractions, energy fractions, and mass fractions can all be different values for the same mixture due to the different properties of liquid and vapor phases.

Question 10

A substance undergoes sublimation at constant temperature and pressure. During this process, what happens to the average intermolecular distance?

  1. Decreases significantly as molecules pack more efficiently
  2. Remains approximately constant throughout the process
  3. Increases significantly as solid transforms to vapor (correct answer)
  4. First decreases then increases during the transition
  5. Changes are negligible compared to other molecular properties
Explanation: This question tests your understanding of phase transitions and molecular behavior, specifically what happens to molecular spacing when matter changes between different states. During sublimation, a solid transforms directly into a gas without passing through the liquid phase. This dramatic change in physical state fundamentally alters how molecules are arranged and spaced. In the solid phase, molecules are tightly packed in an ordered structure with minimal intermolecular distances. When sublimation occurs, these molecules gain enough energy to completely overcome intermolecular forces and enter the gas phase, where they become widely separated and move freely through space. The correct answer is C because the transition from solid to vapor involves a massive increase in intermolecular distance. Gas molecules occupy roughly 1000 times more volume than the same number of molecules in solid form, meaning the average distance between molecules increases dramatically. Option A contradicts basic phase behavior—molecules don't pack more efficiently when going from solid to gas. Option B is incorrect because constant intermolecular distance would mean no phase change occurred, since different phases are defined by different molecular arrangements and spacing. Option D suggests a two-step process, but sublimation is a direct solid-to-gas transition without intermediate decreasing distances. When studying phase transitions, remember that molecular spacing always increases when moving from more ordered phases (solid) to less ordered phases (gas). The key pattern is: solid < liquid < gas in terms of intermolecular distances and molecular freedom.

Question 11

Water vapor at 150°C150°C and 55 bar is compressed isothermally to 1010 bar. Given that the saturation pressure of water at 150°C150°C is approximately 4.764.76 bar, what is the final state?

  1. Superheated vapor at higher pressure than initial state
  2. Saturated vapor at the dew point condition
  3. Liquid-vapor mixture with some condensation occurred (correct answer)
  4. Compressed liquid after complete condensation
  5. Saturated liquid at the bubble point condition
Explanation: When analyzing phase changes during compression, you need to compare the system pressure to the saturation pressure at the given temperature to determine what phase transitions occur. Initially, the water vapor is at 150°C150°C and 55 bar. Since the saturation pressure at 150°C150°C is 4.764.76 bar, and 5>4.765 > 4.76, the initial state is superheated vapor. During isothermal compression to 1010 bar, the temperature remains constant at 150°C150°C, but the pressure increases significantly above the saturation pressure. At 150°C150°C, once the pressure exceeds 4.764.76 bar during compression, condensation begins. Since the final pressure (1010 bar) is much higher than the saturation pressure, significant condensation occurs, but the system hasn't reached complete liquefaction. The result is a liquid-vapor mixture where both phases coexist at equilibrium. Option A is incorrect because while the final pressure is higher, the vapor has partially condensed, so it's no longer purely superheated vapor. Option B is wrong because the dew point condition occurs exactly at saturation pressure (4.764.76 bar), not at 1010 bar. Option D is incorrect because complete condensation to compressed liquid would require much higher pressures at this temperature. Remember this key principle: when pressure exceeds saturation pressure at constant temperature, condensation occurs. If the final pressure is moderately above saturation pressure, expect a liquid-vapor mixture rather than complete phase change. Always compare operating pressure to saturation pressure to predict phase behavior.

Question 12

A pure substance at 50°C50°C exists as a liquid-vapor mixture with 25%25\% vapor by mass. If the pressure is decreased while temperature remains constant, what is the most likely outcome?

  1. More liquid vaporizes, increasing the vapor mass fraction (correct answer)
  2. Some vapor condenses, decreasing the vapor mass fraction
  3. The mass fractions remain unchanged but total volume increases
  4. The mixture becomes entirely liquid due to pressure reduction
  5. The mixture becomes entirely vapor due to pressure reduction
Explanation: When you encounter a liquid-vapor mixture problem, you're dealing with phase equilibrium at the saturation temperature. At any given temperature, a pure substance can only exist as a mixture at one specific pressure—the saturation pressure. Understanding how pressure changes affect this equilibrium is crucial. At 50°C50°C, this substance exists at its saturation pressure with 25% vapor by mass. When pressure decreases while temperature stays constant, you're moving away from the saturation conditions. Since the temperature remains at 50°C50°C but pressure drops below the saturation pressure, the system responds by converting more liquid to vapor to maintain equilibrium. This increases the vapor mass fraction, making choice A correct. Choice B suggests the opposite—that vapor condenses when pressure decreases. This contradicts basic phase behavior; lower pressure actually favors the vapor phase, not the liquid phase. Choice C incorrectly assumes the mass fractions stay constant. While total volume would indeed increase, the key insight is that the mass distribution between phases must change to reach the new equilibrium state. Choice D represents a fundamental misunderstanding—reducing pressure on a liquid-vapor mixture promotes vaporization, not condensation to pure liquid. Remember this pattern: for a pure substance at constant temperature, decreasing pressure always shifts the equilibrium toward more vapor, while increasing pressure favors more liquid. This relationship is the foundation of many thermodynamic processes, including refrigeration cycles and steam power systems.

Question 13

Ice at −10°C-10°C and 11 atm is compressed isothermally to 500500 atm. Based on typical phase behavior, what is the most likely final phase?

  1. The ice remains in the same solid phase structure
  2. The ice transforms to liquid water due to pressure
  3. The ice transforms to a different solid phase structure (correct answer)
  4. The ice transforms directly to vapor phase
  5. The ice becomes a solid-liquid mixture
Explanation: When you encounter phase transformation problems involving extreme pressure changes, you need to consider how pressure affects the stability of different crystal structures, not just the basic solid-liquid-gas transitions you learned in general chemistry. At −10°C-10°C and 11 atm, ice exists in its familiar hexagonal structure (Ice Ih). However, water has multiple solid phases with different crystal arrangements that become stable under different pressure conditions. When you compress ice to 500500 atm while maintaining constant temperature, you're moving into a pressure regime where a different crystal structure becomes thermodynamically favored. Looking at the answer choices: Option A incorrectly assumes the original ice structure remains stable under such high pressure - but different crystal packings have different densities and become favorable as pressure increases. Option B reflects the common misconception that pressure always melts ice, but at −10°C-10°C, even 500500 atm isn't enough to overcome the temperature effect keeping water in solid form. Option D suggests sublimation, which makes no physical sense since increasing pressure actually suppresses vapor formation. Option C correctly identifies that ice will transform to a different solid phase structure - likely Ice III or Ice V, which have different molecular arrangements that are more stable under high pressure conditions. Study tip: For thermodynamics exams, remember that water has at least 15 different solid phases. When you see extreme pressure changes with ice, think about polymorphic transitions between different crystal structures, not just the simple solid-liquid-gas diagram from introductory courses.

Question 14

A vessel contains both liquid water and water vapor in equilibrium at 150°C150°C. If heat is added at constant temperature, what happens to the pressure inside the vessel?

  1. Pressure increases proportionally to the heat added
  2. Pressure decreases as more liquid vaporizes
  3. Pressure remains constant until all liquid vaporizes (correct answer)
  4. Pressure first increases then decreases
  5. Pressure oscillates around the initial value
Explanation: When you encounter a problem involving liquid-vapor equilibrium at constant temperature, you're dealing with a saturated system where the vapor pressure depends only on temperature, not on the relative amounts of liquid and vapor present. At 150°C150°C, water has a specific saturation pressure (approximately 4.76 bar). When you add heat to this two-phase system at constant temperature, the energy goes entirely into converting liquid water to vapor through the latent heat of vaporization. Since temperature remains constant, the pressure must also remain constant at the saturation pressure for that temperature. This continues until all liquid has been converted to vapor. Option A is incorrect because it assumes the system behaves like a single-phase gas where added heat would increase temperature and pressure. However, in a two-phase system, temperature and pressure are locked together by the saturation relationship. Option B gets the physics backward. While more liquid does vaporize when heat is added, this doesn't cause pressure to decrease. The pressure stays constant at the saturation value. Option D suggests some complex pressure variation, but there's no physical basis for pressure to first increase then decrease in this scenario. The saturation pressure remains constant throughout the phase change process. Remember this key principle: in any liquid-vapor equilibrium at constant temperature, pressure remains fixed at the saturation pressure regardless of how much heat you add, until you reach a single-phase boundary (all liquid vaporized or temperature changes).

Question 15

A liquid-vapor mixture of refrigerant has 30%30\% liquid by mass at −10°C-10°C. If the mixture is heated at constant pressure until the quality becomes 0.90.9, what change occurred?

  1. 60%60\% of the original liquid mass vaporized
  2. 20%20\% of the total mixture mass vaporized
  3. 67%67\% of the original liquid mass vaporized (correct answer)
  4. 90%90\% of the original mixture mass vaporized
  5. 10%10\% of the original vapor mass condensed
Explanation: When you encounter problems involving liquid-vapor mixtures and quality changes, you're working with mass conservation during phase transitions. The key insight is tracking what happens to the original liquid and vapor components separately. Let's define the problem systematically. Initially, you have 30% liquid by mass, which means the quality (vapor fraction) is 0.7. After heating, the quality becomes 0.9. Let's say you start with 100 kg total mass for easy calculation. Initially: 30 kg liquid + 70 kg vapor Finally: 10 kg liquid + 90 kg vapor (since quality = 0.9) The change in liquid mass is 30 kg - 10 kg = 20 kg of liquid vaporized. To find what percentage of the original liquid this represents: 20 kg vaporized30 kg original liquid=23=67%\frac{20 \text{ kg vaporized}}{30 \text{ kg original liquid}} = \frac{2}{3} = 67\% This confirms answer C is correct. Now for the wrong answers: A) calculates 2030=0.6=60%\frac{20}{30} = 0.6 = 60\% but makes an error in the arithmetic. B) calculates what fraction of the total mixture mass vaporized (20100=20%\frac{20}{100} = 20\%), which answers a different question entirely. D) misinterprets quality as the percentage of total mass that vaporized, confusing the final state with the change that occurred. Study tip: In quality problems, always distinguish between the final state and the change that occurred. Quality tells you the current vapor fraction, but to find how much vaporized, you need to compare initial and final liquid masses specifically.

Question 16

A container holds saturated water vapor at 200°C200°C. The container volume is slowly decreased while maintaining constant temperature. What sequence of events occurs?

  1. Vapor → superheated vapor → supercritical fluid
  2. Saturated vapor → liquid-vapor mixture → saturated liquid → compressed liquid (correct answer)
  3. Saturated vapor → compressed vapor → liquid vapor mixture → compressed liquid
  4. Vapor → direct transformation to solid phase
  5. Vapor remains saturated throughout with increasing pressure only
Explanation: When you encounter phase transition problems involving constant temperature and changing volume, you're dealing with the pressure-volume relationship along isotherms on a phase diagram. At 200°C200°C, water exists as saturated vapor at its saturation pressure. As you decrease the container volume while maintaining constant temperature, you're increasing the pressure of the system. This follows a predictable sequence: The saturated vapor first reaches a point where condensation begins, creating a liquid-vapor mixture where both phases coexist. As volume continues decreasing, more vapor condenses until you have only saturated liquid. Further volume reduction increases pressure beyond saturation, creating compressed liquid (also called subcooled liquid). Option A is incorrect because superheated vapor occurs when temperature exceeds saturation temperature at a given pressure, not when pressure increases at constant temperature. A supercritical fluid only exists above the critical point (374°C374°C and 22.1 MPa22.1 \text{ MPa} for water). Option C incorrectly suggests "compressed vapor" as an intermediate state. At constant temperature below the critical point, increasing pressure on saturated vapor immediately causes condensation—there's no compressed vapor phase. Option D is wrong because at 200°C200°C, water cannot freeze under any pressure conditions. Solid formation would require much lower temperatures. Remember this key principle: At constant temperature below the critical point, increasing pressure on a saturated vapor always leads to condensation, not superheating. Focus on understanding the saturation curve and how pressure changes affect phase equilibrium.

Question 17

A pure substance at its critical point undergoes a small decrease in temperature at constant pressure. What phase behavior is observed?

  1. The substance immediately separates into distinct liquid and vapor phases
  2. The substance gradually transitions from vapor-like to liquid-like properties (correct answer)
  3. The substance remains in a single phase with unchanged properties
  4. The substance oscillates between liquid and vapor phases
  5. The substance transforms directly to solid phase
Explanation: When analyzing phase behavior near the critical point, you need to understand that this represents a unique thermodynamic state where the distinction between liquid and vapor phases disappears. At the critical point, the densities of liquid and vapor become identical, eliminating the phase boundary. Starting from the critical point and decreasing temperature at constant pressure, the substance doesn't undergo an abrupt phase transition. Instead, it experiences a continuous, gradual change in properties. The substance smoothly transitions from having vapor-like characteristics (lower density, higher compressibility) to more liquid-like properties (higher density, lower compressibility) without any discontinuous jump or phase separation. This makes option B correct. Option A is wrong because there's no sudden separation into distinct phases—that would occur if you crossed a phase boundary away from the critical region, not when starting from the critical point itself. Option C incorrectly suggests no property changes occur, but density, compressibility, and other intensive properties do change continuously as you move away from the critical point. Option D describes an impossible oscillating behavior that has no physical basis in thermodynamics. The key insight is that the critical point is special precisely because it's where the phase boundary terminates. Moving away from it along certain paths (like constant pressure cooling) results in smooth, continuous property changes rather than the sharp transitions you'd see elsewhere on the phase diagram. Study tip: Remember that the critical point eliminates phase boundaries—think "smooth transitions" whenever you encounter critical point problems.

Question 18

A substance has a negative slope on its solid-liquid phase boundary in a pressure-temperature diagram. When this solid is compressed at constant temperature, what happens?

  1. The solid always remains solid regardless of pressure increase
  2. The solid may melt if sufficient pressure is applied (correct answer)
  3. The solid will sublimate directly to vapor phase
  4. The solid will transform to a different crystal structure only
  5. The solid will decompose into constituent elements
Explanation: When you encounter phase diagrams in thermodynamics, pay close attention to the slopes of phase boundaries—they reveal crucial information about how substances behave under different conditions. Most substances have a positive slope on their solid-liquid boundary, but a few important ones like water have a negative slope. A negative slope on the solid-liquid phase boundary means that as pressure increases at constant temperature, you move from the solid region into the liquid region of the phase diagram. This is exactly what happens when you compress the solid—if you apply sufficient pressure, the solid will melt. This counterintuitive behavior occurs because the liquid phase is actually denser than the solid phase for these substances, so increased pressure favors the denser liquid state. Looking at the wrong answers: Choice A incorrectly assumes the solid always remains solid, ignoring the phase boundary crossing that occurs with sufficient pressure. Choice C confuses this with sublimation, which would involve crossing the solid-vapor boundary, not the solid-liquid boundary we're dealing with here. Choice D focuses on crystal structure changes within the solid phase, which misses the point entirely—we're talking about a complete phase transition from solid to liquid. Remember this key pattern: negative slope on the solid-liquid boundary means "pressure can cause melting." Water is the classic example—think of ice skating, where pressure from the blade helps create a thin layer of liquid water. Always trace your path on the phase diagram when pressure or temperature changes.

Question 19

Steam at 120°C120°C and 11 atm is cooled at constant pressure until it reaches 80°C80°C. What is the correct description of this process?

  1. Superheated vapor → saturated vapor → liquid-vapor mixture → compressed liquid (correct answer)
  2. Superheated vapor → liquid-vapor mixture → compressed liquid only
  3. Saturated vapor → liquid-vapor mixture → compressed liquid only
  4. Superheated vapor → compressed liquid directly without two-phase region
  5. Superheated vapor → saturated vapor → compressed liquid without mixture
Explanation: When analyzing phase changes at constant pressure, you need to determine the initial and final states, then trace the path between them using saturation properties. At 1 atm, water's saturation temperature is 100°C. Since the steam starts at 120°C and 1 atm, it's above the saturation temperature, making it superheated vapor. As cooling begins, the steam remains superheated until it reaches 100°C, where it becomes saturated vapor. Continued cooling at constant pressure forces condensation in the two-phase region (liquid-vapor mixture) as temperature stays constant at 100°C. Once fully condensed, further cooling below 100°C creates compressed liquid (also called subcooled liquid) until reaching the final state of 80°C. Answer A correctly captures this complete sequence: superheated vapor → saturated vapor → liquid-vapor mixture → compressed liquid. Answer B skips the saturated vapor state, jumping directly from superheated to two-phase, which ignores the critical transition point at saturation temperature. Answer C incorrectly assumes the process begins with saturated vapor, missing that 120°C at 1 atm is actually superheated conditions. Answer D suggests the steam bypasses the two-phase region entirely, which violates thermodynamic principles—any vapor cooling at constant pressure below its saturation temperature must pass through the liquid-vapor region. Study tip: Always check if your initial state is above, at, or below saturation conditions for the given pressure. This determines whether you start with superheated vapor, saturated vapor, or liquid, which is crucial for mapping the correct phase change sequence.

Question 20

At the triple point of a pure substance, which condition must be satisfied?

  1. The substance exists only in the vapor phase due to low pressure
  2. All three phases have identical densities and cannot be distinguished
  3. The rates of phase transitions between all three phases are equal
  4. All three phases coexist in thermodynamic equilibrium simultaneously (correct answer)
  5. The substance rapidly cycles between all three phases in sequence
Explanation: When you encounter questions about triple points, you're dealing with one of thermodynamics' most specific equilibrium conditions. The triple point represents a unique state where temperature and pressure are precisely defined for a given substance. The correct answer is D because the triple point is fundamentally defined as the unique combination of temperature and pressure where solid, liquid, and vapor phases coexist in thermodynamic equilibrium. At this point, all three phases are present simultaneously and stable - none is converting to another on a net basis. This equilibrium means the chemical potential of all three phases is identical, making phase transitions thermodynamically neutral. Let's examine why the other options miss the mark. Option A incorrectly suggests only vapor exists due to low pressure. In reality, triple points typically occur at relatively low pressures, but all three phases must be present by definition. Option B contains a fundamental misconception - the phases maintain their distinct densities and physical properties at the triple point; they don't become indistinguishable. Option C confuses kinetic rates with thermodynamic equilibrium. While individual molecules continuously transition between phases, the net rates are zero (equal forward and reverse transitions), not that all transition rates are equal to each other. Remember this key distinction: the triple point isn't about phase transition rates or phase properties becoming identical - it's about the thermodynamic stability of three distinct phases existing together. When you see "triple point" on exams, immediately think "three phases in equilibrium simultaneously."