You start with water vapor at very high temperature and gradually increase pressure while keeping the temperature constant above the critical temperature. Which of the following correctly describes what happens?
AThe vapor crosses the liquid-vapor boundary and becomes liquid
BThe vapor undergoes a sharp phase transition to supercritical fluid
CThe vapor continuously densifies into a supercritical fluid with no distinct phase transition
DThe vapor cannot be compressed above the critical temperature
Above the critical temperature, the liquid-vapor boundary no longer exists — it ends at the critical point. Increasing pressure continuously densifies the fluid without any sharp transition or latent heat. This 'going around' the critical point is a key feature of supercritical fluids: you can move between gas-like and liquid-like densities continuously. Option A is wrong because the liquid-vapor boundary only exists below the critical temperature.
Question 2 Multiple Choice
Why does increasing pressure cause ice to melt, a behavior opposite to almost all other solids?
AWater has an unusually high latent heat of fusion, so pressure supplies the needed energy
BIce is less dense than liquid water, so the solid-liquid boundary has a negative slope
CThe triple point of water is below atmospheric pressure, forcing melting at high pressure
DIce has stronger hydrogen bonds than liquid water, which are broken by pressure
The Clausius-Clapeyron equation dP/dT = L / (TΔv) governs the slope of phase boundaries. For the solid-liquid transition, Δv = v_liquid − v_solid. For most substances, solids are denser than liquids (Δv > 0), giving a positive slope. But for water, ice is less dense than liquid water (Δv < 0), so the slope is negative — increasing pressure pushes the boundary leftward (toward lower temperature), meaning ice melts under increased pressure at constant temperature. This negative slope is what makes ice skating and pressure-induced melting possible.
Question 3 True / False
The critical point and the triple point of a substance occur at the same pressure and temperature.
TTrue
FFalse
Answer: False
These are entirely different features of a phase diagram. The triple point is the unique P-T coordinate where all three phases (solid, liquid, vapor) coexist in equilibrium — for water it is at 273.16 K and 0.006 atm. The critical point marks where the liquid-vapor boundary terminates; above it, the substance is a supercritical fluid. For water the critical point is at 647 K and 218 atm. They cannot coincide because the triple point is at the intersection of three phase boundaries, while the critical point is the endpoint of one.
Question 4 True / False
At pressures below the triple point of a substance, heating a solid will cause it to sublimate directly to vapor without passing through a liquid phase.
TTrue
FFalse
Answer: True
The triple point is the minimum pressure at which the liquid phase can exist. Below this pressure, the liquid-vapor boundary does not exist — the phase diagram goes directly from solid to vapor. Heating a solid at such low pressures causes sublimation (solid → vapor), bypassing the liquid phase entirely. This is exactly how freeze-drying works: food is frozen and placed in a vacuum below water's triple-point pressure (0.006 atm), so the ice sublimes rather than melting.
Question 5 Short Answer
Why is water's phase diagram considered anomalous compared to most other substances, and what physical property of water causes this anomaly?
Think about your answer, then reveal below.
Model answer: For most substances, the solid is denser than the liquid, so the solid-liquid boundary has a positive slope — increasing pressure favors the denser solid phase. Water is anomalous because ice is less dense than liquid water (due to hydrogen bonding creating an open crystalline lattice). This means Δv for melting is negative, flipping the Clausius-Clapeyron slope to negative: the solid-liquid line tilts to the left. Practical consequence: ice melts under pressure rather than solidifying.
The anomaly has significant physical consequences beyond phase diagrams — it's why ice floats (insulating aquatic life beneath frozen surfaces), why glaciers flow, and why pipes burst when water freezes. The underlying cause is hydrogen bonding: when water freezes, molecules lock into a hexagonal lattice that is actually more open (lower density) than the disordered arrangement in liquid water.