Give water an energy budget. Watch its temperature and phase masses, catch a melting plateau, compare two starting states and inspect an open ice structure.
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Heat can change a sample’s temperature—or its phase mixture. During melting at fixed pressure, the energy can change how much is liquid while the thermometer barely changes.
A 0°C sample changes from 25% to 75% liquid by mass. Must the thermometer rise?
For the virtual 100 g sample, converting an additional 50 g requires 16.670 kJ while the coexistence temperature stays approximately 0°C.
Heat is energy transferred because of a temperature difference. Our controls prescribe a net heat amount. The energy receipt shows which part warms a phase and which part changes the mixture. Temperature and transferred heat have different units and roles.
At the modeled melting temperature, adding heat increases liquid mass while temperature stays at approximately 0°C. A half-melted sample is not halfway to the boiling temperature. Pin two different mixtures with the same thermometer reading.
A heat pulse may first warm cold ice to melting, then melt some of it. Any energy left after a boundary still counts. The receipt carries it into the next interval rather than discarding it or pretending the ice instantly became warm liquid.
Solid ice, liquid water and water vapor contain H₂O. Their arrangement and separation change. Boiling does not turn each water molecule into hydrogen gas and oxygen gas. A visible mist contains droplets; individual vapor molecules are not little white clouds.
The same heat shared by twice the single-phase mass produces half the temperature rise in this approximation. Phase-change budgets also scale with mass. Compare the curves per sample, then per kilogram.
Removing energy can freeze liquid or condense vapor. At an equilibrium plateau, the mixture changes without a temperature change. Zero net heat leaves the model’s temperature and phase masses fixed, even while the molecular illustration keeps moving.
Ordinary pure H₂O at nominal 101.325 kPa, retaining all mass with an ideal movable boundary. It represents uniform equilibrium states with pressure work only and negligible bulk kinetic/potential changes. It is not a rigid sealed vessel or an uncovered saucepan. At these conditions Q = ΔH, where H = U + pV; expansion work must not be subtracted a second time.
The model uses rounded transition temperatures 0°C and 100°C, with a domain of −20°C to 120°C. Ice cₚ = 2020, liquid cₚ = 4180 and vapor cₚ = 2040 J/(kg·K). Fusion enthalpy is 333400 J/kg and vaporization enthalpy 2256400 J/kg. These are not a full IAPWS equation-of-state implementation.
The ice capacity is a midpoint approximation to NIST’s −20, −10 and 0°C values. The vapor capacity comes from a 100–120°C enthalpy secant at 0.10 MPa. Vaporization enthalpy is tabulated at 100°C saturation and 101.42 kPa, slightly above the model’s nominal pressure. The constant-property choices and rounded transition points are explicitly approximate.
The enthalpy reference is ice at −20°C. For 100 g, interval boundaries are 4.040, 37.380, 79.180, 304.820 and 308.900 kJ. Within a single phase ΔT = Q/(mcₚ); within coexistence the converted mass is Q/L. A direct piecewise decoder preserves energy across any number of boundaries.
At approximately 0°C, temperature alone does not distinguish 25% from 75% liquid by mass. Changing between those mixtures for 100 g requires 16.670 kJ in the supplied model. At vaporization, the fraction is also by mass—not the fraction of vessel volume occupied by vapor.
Near melting, reference densities are 916.7 kg/m³ for ice and 999.84 kg/m³ for liquid. Equal 100 g masses occupy about 109.09 and 100.02 mL. Freezing therefore increases volume by about 9.07% in this comparison. These two densities are not used to predict thermal expansion throughout the whole heating path.
The representative molecule positions never calculate temperature, pressure or phase fractions. Their motion is an illustration, not molecular dynamics. Zero Celsius is not zero thermal motion. A paused drawing does not mean a physically motionless solid.
Water can evaporate into surrounding air below its boiling temperature. That open-system mass and heat transfer lies outside this closed, equilibrium enthalpy path. Real heat-transfer rates, gradients, nucleation, supercooling and dissolved substances require additional physics.
The lower and upper energy limits correspond to −20°C and 120°C in the supplied approximation. A requested pulse beyond a limit is split into accepted and unapplied energy. Unapplied energy is not transferred to the sample. Reversal uses accepted heat, not a request the model could not accept.
A phase change can absorb energy without a large temperature rise. The model explains that principle; predicting a real cold pack requires its material, quantity and heat-transfer conditions.
Its horizontal axis may be energy, energy per kilogram or time. Those choices make different claims. Our full energy axis preserves the much larger vaporization budget.
Floating ice and freezing expansion invite a density question. A separate equal-mass volume comparison keeps that question distinct from the phase-fraction bar.
Place them together on a tray indoors, away from strong drafts and direct sun. Label A and B. Note cup material and approximate ice sizes; similar-looking pieces are not necessarily equal mass.
Leave A without added liquid. Put about 100 mL room-temperature water in B. Predict how these different surroundings might affect melting and record a reason.
Use a spoon to place one ice piece in each cup at nearly the same time. Start the timer. Notice whether the piece in B floats; sketch it without pushing it under.
Every two minutes for about 20 minutes, or until no solid is visible, record substantial ice, a smaller piece or no visible solid. A sketch is not a measurement of mass fraction. Retain incomplete melting observations too.
If using a suitable thermometer, optionally measure B after gentle consistent stirring, keeping the probe off the wall and visible ice. Record placement and instrument resolution. No particular temperature reading is required.
Water, air and the cup can supply heat; size, contact and starting temperatures also matter. Compare the histories without promising a universal winner or completion time.
On paper, draw the virtual 100 g half-melted sample. The model needs 8.335 kJ to make it 75% liquid at about 0°C. Mark that as calculated, beside your observed table.
Empty cups into a sink and wipe the tray. Explain why unknown room heat, initial ice temperature, dissolved substances and gradients prevent these observations from directly determining fusion enthalpy.
Where can the energy that melts ice come from?
Observe ordinary ice in open cups at room temperature. No boiling, sealed/heated vessel or pressure device. Do not drink the activity water. These uncontrolled observations do not measure latent heat or certify an ideal plateau.
Phase mixture and enthalpy, with temperature unchanged Converting 50 g in the supplied model needs 0.050 × 333400 = 16670 J.
40°C and 30°C With no phase boundary crossed, ΔT = Q/(mcₚ).
About 88 g ice and 12 g liquid at 0°C 1.010 kJ warms the ice; 3.990 kJ melts approximately 11.97 g.
Each keeps its H₂O identity The phase change concerns organization and spacing rather than chemical decomposition.
100 g ice ≈109 mL; 100 g liquid ≈100 mL Ordinary ice’s open structure gives a lower density than near-melting liquid water.
Same temperature and phase masses, with thermal motion still represented Macroscopic equilibrium can coexist with microscopic motion.
The supplied vaporization enthalpy is much larger than fusion enthalpy For 100 g, the two latent budgets are 225.64 and 33.34 kJ at their stated reference conditions.
No, evaporation into air below boiling is outside this closed path An open evaporating cloth requires mass/heat-transfer physics absent from the retained equilibrium sample.
COD9015208, revision 292002; Fortes et al. (2004), J. Chem. Phys. 120, 11376–11379, DOI 10.1063/1.1765099. Powder neutron-diffraction refinement at 100 K. Oxygen symmetry expanded; half-occupied H sites omitted, not assigned unique orientations. Source CIF is preserved.
COD / Fortes et al. · Ice structural dataCOD data are placed under CC0. Author and source attribution are retained; this does not license the original journal article artwork.
Crystallography Open Database · Data policyUnmodified owner-linked original JPEG, USGS source marked Public Domain, approximately 2012. No instrument, optical method or scale is inferred.
USGS · Ice bubbles and crystalsIAPWS R10-06(2009), release p. 3 and formulation verification tables. The constant-property lesson does not implement this complete equation of state.
IAPWS · Ice Ih formulationAllan H. Harvey, Properties of Ice and Supercooled Water (2019), visible ordinary-H₂O tables at 101.325 kPa. Ice cₚ varies over the range; 2020 J/(kg·K) is the chosen midpoint approximation.
Harvey / NIST · Ice and supercooled waterIAPWS-95 revised release. This model uses narrow constant-property approximations rather than the full Helmholtz formulation.
IAPWS · Ordinary water formulationFirst-party FAQs. Normal boiling is near 99.974°C on ITS-90; familiar 100°C is rounded here.
IAPWS · Water and steam FAQsMIT thermodynamics textbook, chapter 3 §§3.2–3.2.2. Equations are independently implemented; source artwork is not redistributed.
MIT · Thermodynamics and Climate ChangeNISTIR 5078 Table 1, 100°C saturation row: 101.42 kPa and 2256.4 kJ/kg. Reference calculations, not a measured household heating trace.
Harvey / NIST · Saturation by temperatureNISTIR 5078 Table 3, 0.10 MPa, 100–120°C enthalpy difference 2675.8→2716.6 kJ/kg. This pressure is 100 kPa, not exactly 101.325 kPa.
Harvey / NIST · Compressed water and superheated steamCipcigan et al., Scientific Reports 8, 1718 (2018), original QDO-water simulations. The lesson illustration does not reproduce this trajectory model.
Cipcigan et al. · Water structure and hydrogen bondingFirst-party chemistry teaching resource; no ACS animation or artwork copied.
American Chemical Society · Changing state: meltingOfficial water-cycle terminology. The molecular illustration is not a photograph of steam.
USGS · Water-cycle glossaryOfficial vapor-pressure explanation. Open-cup evaporation is distinct from the lesson’s closed equilibrium path.
USGS · Vapor pressure and waterIndependent subject review is pending.
Read the sources and model assumptions
Interactive explorationOpen a house wall. Swap its added layer, make it thicker, and reverse the temperatures to see why insulation helps in winter and summer.
Interactive explorationFollow water above and below the surface. Pave part of the land, change evaporation, and watch the water budget rearrange.
Interactive explorationLook through a cold room and follow its refrigerant loop. Count the heat removed and the work added, then try putting both sides in the same room.