Chapter 19 — Front Matter
Purpose and Overview
You have seen steam rising from a pot of boiling water, frost forming on a car window overnight, and ice melting on a warm afternoon. You have watched puddles disappear without boiling and noticed that a glass of water left open for a week holds less water than before. All of these observations involve the same thing: matter changing from one physical state to another.
In Chapter 18, you studied gases and the laws that govern how gas particles behave. Gases are the simplest state of matter to model mathematically because gas particles move freely, spread apart, and interact very little with one another. Liquids and solids are more complex. Their particles are close together, and the forces between particles matter enormously.
This chapter explores what happens inside liquids and solids at the particle level — and what happens when matter crosses the boundary from one state to another. You will study melting, freezing, vaporization, condensation, sublimation, and deposition. You will learn how to read heating curves and cooling curves. And you will connect these ideas to things you already know: why water boils at 100 °C, why some solids feel harder than others, why humidity affects how hot a day feels, and why dry ice seems to vanish into air.
The central idea of this chapter is simple but powerful: the state of matter that a substance occupies depends on the balance between the energy of its particles and the strength of the forces pulling those particles together.
Opening Question
Why does dry ice disappear into vapor without ever forming a puddle? Why does water boil at a lower temperature in Denver than in New York City? Why does a steam burn feel worse than a boiling-water burn at exactly the same temperature? By the end of this chapter, you will be able to explain all three using the same set of ideas: particle energy, intermolecular forces, and phase changes.
9Sections
6Phase Changes
3Worked Examples
30+Key Terms
Learning Objectives
- 1Compare the arrangement, motion, and energy of particles in solids, liquids, and gases.
- 2Distinguish between crystalline and amorphous solids and explain how their structures affect melting behavior.
- 3Explain how temperature relates to particle kinetic energy and why a distribution of energies matters.
- 4Connect the type and strength of intermolecular forces to a substance's melting point, boiling point, and state at room temperature.
- 5Describe and distinguish all six phase changes, classifying each as endothermic or exothermic.
- 6Explain why temperature remains constant during a phase change even when heat is being added or removed.
- 7Define vapor pressure and explain how it determines the boiling point of a liquid and how altitude affects boiling.
- 8Calculate the energy absorbed or released during melting, freezing, vaporization, or condensation using the heat of fusion and heat of vaporization.
- 9Read and interpret heating curves and cooling curves, identifying phases, phase changes, melting points, and boiling points.
- 10Interpret a phase diagram, identify the triple point and critical point, and use the diagram to predict phase behavior under given conditions.
- 11Connect phase change principles to real applications in weather, cooking, and materials science.
Sections 19.1 – 19.3
The Nature of Solids, Liquids, and Gases
Particle Arrangement · Motion · Intermolecular Forces · States at Room Temperature
Section 19.1
Comparing Solids, Liquids, and Gases
Matter on Earth exists in three common states: solid, liquid, and gas. These states are sometimes called the three phases of matter. A fourth state, plasma, exists at extremely high temperatures and is found in stars and lightning, but it is beyond the scope of this chapter.
The three common phases of matter differ in how particles are arranged, how much they move, and how strongly they attract one another.
Solids
In a solid, particles are packed tightly together and held in fixed positions by strong attractive forces. The particles vibrate in place but do not slide past one another or move freely through the material. Because particles are locked in place, solids have a definite shape and a definite volume.
Crystalline solids have particles arranged in a highly ordered, repeating geometric pattern called a crystal lattice. Table salt (NaCl), quartz (SiO₂), diamond (C), and sucrose are crystalline solids with sharp, well-defined melting points. Amorphous solids — glass, rubber, many plastics — have particles arranged irregularly and soften gradually over a range of temperatures rather than at one sharp point.
Liquids
In a liquid, particles are still close together but are not locked in fixed positions. They are free to slide past one another, which allows the liquid to flow and take the shape of whatever container holds it. A liquid has a definite volume but no definite shape. Because the particles are still close together, liquids are nearly incompressible.
Gases
In a gas, particles are spread far apart and move rapidly in all directions. The attractive forces between gas particles are so weak that particles behave almost independently. A gas has neither a definite shape nor a definite volume — it expands or compresses to fill the available space.
| Property | Solid | Liquid | Gas |
|---|
| Particle arrangement | Ordered, fixed positions | Close, disordered | Far apart, random |
| Particle movement | Vibrate in place | Slide past each other | Move rapidly in all directions |
| Shape | Definite | Takes container shape | Takes container shape |
| Volume | Definite | Definite | Expands to fill container |
| Compressibility | Very low | Very low | High |
| Intermolecular forces | Strong | Moderate | Very weak (negligible) |
- Explain the difference between a crystalline solid and an amorphous solid. Give one example of each.
- A liquid takes the shape of its container but does not expand to fill it. A gas expands to fill its container. Why do they behave differently? Explain in terms of particle spacing and intermolecular forces.
- Why are both solids and liquids nearly incompressible while gases are easily compressed?
- Water is a liquid at room temperature. Oxygen (O₂) is a gas at room temperature. What does this tell you about the relative strength of intermolecular forces in each substance?
Section 19.2
Particle Motion in Liquids and Solids
Key Idea — Temperature and Kinetic Energy
Temperature is a measure of the average kinetic energy of the particles in a substance. The higher the temperature, the faster the particles move on average. This holds true in all three phases of matter.
In a solid, particles vibrate back and forth around their fixed positions. As temperature increases, those vibrations become more energetic. If enough energy is added, the vibrations become so strong that particles begin to break free of their fixed positions, and the solid melts.
Not all particles in a sample have the same kinetic energy at a given temperature. Some are moving faster than average, some slower. This distribution of energies explains why a liquid can evaporate at temperatures below its boiling point — even at 20 °C, some water molecules at the surface have enough energy to escape into the gas phase. This is why puddles dry up even on cool, dry days.
Thermal Expansion
As temperature increases, particles in a solid or liquid move with more energy, and the average distance between particles increases slightly. This is called thermal expansion. It explains why metal bridges have expansion joints, why thermometers work, and why glass jars can crack when hot liquid is poured in suddenly.
Water is a famous exception. Liquid water is densest at 4 °C, not at 0 °C. Below 4 °C, water expands as it cools, which means ice is less dense than liquid water. This is why ice floats — a fact with enormous consequences for aquatic life and Earth's climate systems.
Why Ice Floats Matters
If ice sank instead of floating, lakes and ponds would freeze solid from the bottom up, killing aquatic life and dramatically altering Earth's climate. The anomalous expansion of water near its freezing point is one of the properties that makes liquid water uniquely suited to supporting life.
- A sample of liquid water sits at 25 °C. Do all the water molecules have the same kinetic energy? Explain why or why not, and connect your answer to the phenomenon of evaporation.
- Why do most solids expand when heated? What is happening at the particle level?
- Water is densest at 4 °C, not at 0 °C. What does this mean for ice versus liquid water? Why does this matter for life in frozen lakes?
Section 19.3
Intermolecular Forces and States of Matter
The state that a substance occupies at a given temperature depends not only on the kinetic energy of its particles but also on the intermolecular forces — the attractive forces between molecules. You studied these in depth in Chapter 12. Here is a brief review with emphasis on how force strength controls state and phase behavior.
| Force Type | Found In | Relative Strength |
|---|
| London dispersion forces | All molecules (polar and nonpolar) | Weakest; increases with molecular size |
| Dipole-dipole forces | Polar molecules | Moderate; stronger than dispersion alone |
| Hydrogen bonding | Molecules with H bonded to N, O, or F | Strongest common intermolecular force |
The stronger the intermolecular forces, the more energy is needed to pull particles apart — and therefore the higher the melting point and boiling point, and the more likely the substance is to be a liquid or solid at everyday temperatures.
| Substance | Formula | Dominant IMF | Boiling Point |
|---|
| Methane | CH₄ | Dispersion only | −161 °C |
| Hydrogen chloride | HCl | Dispersion + dipole-dipole | −85 °C |
| Water | H₂O | Dispersion + dipole-dipole + hydrogen bonding | 100 °C |
Pattern to Remember
Stronger intermolecular forces → higher boiling point → more likely to be liquid or solid at room temperature. Weaker intermolecular forces → lower boiling point → more likely to be a gas at room temperature.
- Nitrogen (N₂) boils at −196 °C. Water (H₂O) boils at 100 °C. Both are small molecules. Explain the enormous difference in boiling points using intermolecular forces.
- Ethanol (C₂H₅OH) has a boiling point of 78 °C. Dimethyl ether (CH₃OCH₃) has the same molecular formula (C₂H₆O) but boils at −24 °C. What type of intermolecular force does ethanol have that dimethyl ether lacks? (Hint: look at what is bonded to the oxygen in each molecule.)
- Predict whether each substance is most likely a solid, liquid, or gas at room temperature (25 °C), and explain your reasoning: (a) argon (Ar), a noble gas; (b) table salt (NaCl), an ionic solid with melting point 801 °C; (c) ethanol, boiling point 78 °C.
Sections 19.4 – 19.6
Phase Changes: Energy, Direction, and Examples
Melting · Freezing · Vaporization · Condensation · Sublimation · Deposition
Section 19.4
Melting and Freezing
Melting
Melting is the phase change from solid to liquid. During melting, particles in a solid gain enough energy to break free of their fixed lattice positions and begin to move past one another. The temperature at which this occurs at standard pressure (1 atm) is called the melting point of the substance.
| Substance | Melting Point |
|---|
| Ice (H₂O) | 0 °C |
| Table salt (NaCl) | 801 °C |
| Lead (Pb) | 328 °C |
| Gold (Au) | 1064 °C |
| Tungsten (W) | 3422 °C |
Critical Concept — Temperature During Melting
During melting, temperature does not rise. Energy being added is used to break apart attractive forces between particles — converting the solid to liquid. None of that energy is left over to increase particle speed, which is what a temperature rise would require. This is why the melting plateau on a heating curve is flat.
The amount of energy required to melt one gram of a substance at its melting point is the heat of fusion (also called enthalpy of fusion, ΔHfus). Values are given in joules per gram (J/g). The heat of fusion of water is 334 J/g.
Freezing
Freezing is the reverse of melting: the phase change from liquid to solid. Liquid particles lose enough energy that they slow down and are captured by the attractive forces of neighboring particles, locking into the ordered arrangement of a solid. For a pure substance, the freezing point equals the melting point. During freezing, energy is released — freezing is exothermic; melting is endothermic.
Worked Example 19.1
Calculating energy to melt ice
- ProblemHow much energy is required to melt 45.0 g of ice at 0 °C? The heat of fusion of water is 334 J/g.
- Formula
q = m × ΔH_fus
- Solve
q = 45.0 g × 334 J/g
q = 15,030 J = 15.0 kJ
- NoteThis calculation covers only the melting step. Raising the temperature of the resulting liquid water requires an additional calculation using specific heat (Chapter 20).
Answer: 15,030 J (15.0 kJ) are needed to melt 45.0 g of ice at 0 °C.
- How much energy is needed to melt 80.0 g of ice at 0 °C? (ΔHfus of water = 334 J/g)
- How much energy is released when 250. g of liquid water at 0 °C freezes completely?
- The heat of fusion of ethanol is 104 J/g. How much energy is needed to melt 55.0 g of solid ethanol?
- A sample of ice absorbs 20,040 J while melting completely at 0 °C. What is the mass of the sample?
- Explain why freezing is described as an exothermic process. Where does the released energy go?
- A student says the melting point and freezing point of a pure substance are different temperatures. Is this correct? Explain.
Section 19.5
Vaporization and Condensation
Evaporation
Evaporation occurs at the surface of a liquid at temperatures below the boiling point. Even when a liquid is cool, some molecules near the surface have enough kinetic energy to overcome the attractive forces of their neighbors and escape into the gas phase. Evaporation is a slow, surface-level, endothermic process — energy is absorbed from the surroundings. This is why your skin feels cool after sweating.
| Factor | Effect on Evaporation Rate |
|---|
| Temperature | Higher temperature → more particles with sufficient energy → faster evaporation |
| Surface area | Larger surface → more particles near the liquid-gas interface → faster evaporation |
| Air movement (wind) | Carries away vapor → prevents re-entry → faster evaporation |
| Humidity | More water vapor in air → vapor re-enters liquid more → slower evaporation |
Vapor Pressure
In a closed container, evaporation and condensation reach a balance called vapor-liquid equilibrium. The pressure exerted by the vapor in equilibrium with its liquid at a given temperature is the vapor pressure of the liquid. Vapor pressure increases with temperature because more particles have enough energy to escape. Substances with weak intermolecular forces have high vapor pressures and are called volatile.
Boiling Point and Altitude
Boiling occurs when the vapor pressure of the liquid equals the external atmospheric pressure. At sea level (1 atm), water boils at 100 °C. At high altitude, atmospheric pressure is lower, so water boils at a lower temperature — about 95 °C at 2500 meters. This is why food takes longer to cook at high altitude: the water is boiling at a lower temperature. A pressure cooker raises pressure above 1 atm inside a sealed pot, raising the boiling point above 100 °C and cooking food faster.
Temperature During Boiling
Just as temperature does not rise during melting, temperature does not rise during boiling. All the energy added at the boiling point goes into converting liquid to vapor — breaking intermolecular forces — rather than increasing particle speed.
The amount of energy required to vaporize one gram of a liquid at its boiling point is the heat of vaporization (ΔHvap). For water, ΔHvap = 2260 J/g — about 6.8 times larger than the heat of fusion, because vaporization requires completely separating particles, while melting only loosens them from fixed positions.
Condensation
Condensation is the reverse of vaporization: the phase change from gas to liquid. It is exothermic — energy is released to the surroundings. The energy released during condensation equals the energy absorbed during vaporization. This is why steam burns are often more severe than boiling-water burns: condensing steam releases its heat of vaporization directly into your skin in addition to the thermal energy from the hot water.
Worked Example 19.2
Energy released during condensation
- ProblemHow much energy is released when 12.5 g of steam at 100 °C condenses into liquid water at 100 °C? ΔHvap of water = 2260 J/g.
- Formula
q = m × ΔH_vap
- Solve
q = 12.5 g × 2260 J/g
q = 28,250 J = 28.3 kJ
- Interpret28,250 J of energy are released to the surroundings — which is why steam at 100 °C burns more severely than liquid water at 100 °C.
Answer: 28,250 J (28.3 kJ) are released when 12.5 g of steam condenses.
- How much energy is required to vaporize 35.0 g of water at 100 °C? (ΔHvap = 2260 J/g)
- How much energy is released when 40.0 g of steam condenses to liquid water at 100 °C?
- Methanol has a heat of vaporization of 1100 J/g. How much energy is released when 18.0 g of methanol vapor condenses?
- Explain why a humid summer day feels hotter than a dry day at the same air temperature.
- Water boils at approximately 95 °C at the altitude of Denver, Colorado (about 1600 m). Why does pasta take longer to cook there than at sea level?
- Explain using vapor pressure concepts why a pressure cooker reduces cooking time.
- Why does sweating cool the human body? Name the phase change involved and classify it as endothermic or exothermic.
- How many grams of water can be vaporized by 100,000 J of energy? (Use ΔHvap = 2260 J/g)
Section 19.6
Sublimation and Deposition
Sublimation is the phase change directly from solid to gas, bypassing the liquid phase. It occurs when the vapor pressure of the solid equals or exceeds the external pressure without the substance ever reaching its melting point at that pressure. Sublimation is endothermic.
The most familiar example is dry ice — solid carbon dioxide (CO₂). At standard atmospheric pressure, solid CO₂ does not melt. Instead, it sublimes directly at −78.5 °C. Other examples include snow subliming on cold, dry, sunny days (this is why snowfields shrink even when temperatures stay below freezing), iodine crystals producing visible violet vapor at room temperature, and the freeze-drying of food (ice is removed from frozen food by sublimation under reduced pressure).
Deposition is the reverse of sublimation: the direct phase change from gas to solid, bypassing the liquid phase. Deposition is exothermic. Frost forms by deposition — water vapor deposits directly as ice crystals on cold surfaces when the surface temperature is below 0 °C. Frost is not frozen dew. Dew forms by condensation (gas to liquid) and may then freeze; frost forms in a single step directly from vapor to solid.
| Phase Change | Direction | Energy Change | Example |
|---|
| Melting | Solid → Liquid | Endothermic | Ice melting in a glass |
| Freezing | Liquid → Solid | Exothermic | Liquid water freezing overnight |
| Vaporization | Liquid → Gas | Endothermic | Water boiling; puddle evaporating |
| Condensation | Gas → Liquid | Exothermic | Dew forming; steam condensing |
| Sublimation | Solid → Gas | Endothermic | Dry ice disappearing; snow shrinking |
| Deposition | Gas → Solid | Exothermic | Frost forming on a car window |
Common Misconception — Frost vs. Dew
Frost is not frozen dew. Frost forms by deposition — water vapor transitions directly from gas to solid ice. Dew forms by condensation — water vapor becomes liquid water. If the surface later drops below 0 °C after dew has formed, that dew can freeze, but frost itself involves no liquid step.
- Name the phase change taking place in each description: (a) liquid ethanol evaporating from an open flask; (b) water droplets forming on a cold glass; (c) dry ice shrinking on a table; (d) molten iron hardening in a mold; (e) snow crystals forming in clouds; (f) iodine crystals producing violet vapor.
- Classify each phase change in problem 1 as endothermic or exothermic.
- Why does dry ice not produce a liquid puddle as it warms at atmospheric pressure? What would need to change for liquid CO₂ to form?
- Explain freeze-drying using the concept of sublimation. Why does this preserve food better than heat-drying?
- A student says: "Frost is just frozen dew." Correct this misconception using the names of the phase changes involved.
Sections 19.7 – 19.9
Heating Curves, Phase Diagrams, and Real-World Applications
Heating and Cooling Curves · Phase Diagrams · Triple Point · Critical Point · Applications
Section 19.7
Heating Curves and Cooling Curves
A heating curve is a graph that shows how the temperature of a substance changes as energy is steadily added to it. The horizontal axis shows energy added (or time, if heat is added at a constant rate). The vertical axis shows temperature. Heating curves have five distinct regions for a substance that has distinct melting and boiling points.
Five Regions of a Water Heating Curve
Region 1 — Heating ice below 0 °C: Temperature rises as ice warms. Slope depends on specific heat of ice (2.09 J/g·°C).
Region 2 — Melting plateau at 0 °C: Temperature constant. Energy breaks the crystal lattice. All of it goes to the phase change.
Region 3 — Heating liquid water (0 °C to 100 °C): Temperature rises. Slope depends on specific heat of liquid water (4.18 J/g·°C).
Region 4 — Boiling plateau at 100 °C: Temperature constant. Energy breaks all intermolecular forces between liquid molecules. Much longer than the melting plateau because ΔHvap ≫ ΔHfus.
Region 5 — Heating steam above 100 °C: Temperature rises. Slope depends on specific heat of steam (1.87 J/g·°C).
The slope of each rising region reflects the specific heat of the substance in that phase. A gentler slope means more energy is needed per degree — the substance has a higher specific heat. The flat regions (plateaus) occur at the melting and boiling points and represent phase changes at constant temperature.
A cooling curve is the reverse process: a graph showing how temperature changes as energy is steadily removed. It has the same five regions in reverse. The flat regions on a cooling curve represent condensation and freezing, both of which release energy without a temperature change.
Worked Example 19.3
Interpreting a heating curve for an unknown substance
- ProblemA student heats a pure substance starting at −20 °C and ending at 120 °C. The heating curve shows a flat region at 25 °C and another flat region at 85 °C. What conclusions can the student draw?
- ThinkFlat regions on a heating curve are phase changes at constant temperature. The first flat region is the melting point; the second is the boiling point.
- Conclude
Melting point of substance: 25 °C
Boiling point of substance: 85 °C
State at room temperature (~22 °C): solid, since 22 °C < 25 °C melting point
The melting point is 25 °C, the boiling point is 85 °C, and at room temperature the substance is a solid.
Use the data table for questions 1–10. A researcher heats a 10.0 g sample of a pure substance, recording the following:
| Region | Temperature Range | Observation |
|---|
| 1 | −30 °C to 20 °C | Temperature rises steadily |
| 2 | 20 °C | Temperature constant for several minutes |
| 3 | 20 °C to 110 °C | Temperature rises steadily |
| 4 | 110 °C | Temperature constant for a long time |
| 5 | 110 °C to 150 °C | Temperature rises steadily |
- What phase is the substance in during Region 1?
- What phase change is occurring during Region 2?
- What is the melting point of this substance?
- What phase is the substance in during Region 3?
- What phase change is occurring during Region 4?
- What is the boiling point of this substance?
- What phase is the substance in during Region 5?
- Why does temperature remain constant in Regions 2 and 4 even though heat is being added?
- The plateau in Region 4 lasts much longer than Region 2. What does this tell you about the heats of fusion and vaporization?
- Describe in words what a cooling curve for this substance would look like as it cools from 150 °C back to −30 °C.
Use: specific heat of ice = 2.09 J/(g·°C); specific heat of liquid water = 4.18 J/(g·°C); ΔHfus = 334 J/g; ΔHvap = 2260 J/g.
- A 50.0 g sample of ice begins at −10 °C. How much energy is needed to bring it from −10 °C to 0 °C?
- How much additional energy is needed to melt the 50.0 g ice sample completely at 0 °C?
- How much energy is needed to heat the resulting 50.0 g of liquid water from 0 °C to 100 °C?
- How much energy is needed to vaporize all 50.0 g of water at 100 °C?
- What is the total energy required to convert 50.0 g of ice at −10 °C to steam at 100 °C? Which step required the most energy? Explain why in terms of what is happening at the particle level.
Section 19.8
Phase Diagrams
A phase diagram is a graph that shows which phase of a substance is stable under different combinations of temperature and pressure. Temperature appears on the horizontal axis; pressure on the vertical axis. The diagram is divided into three regions — solid, liquid, gas — separated by boundary lines.
| Boundary Line | Separates | Meaning |
|---|
| Solid-liquid line (melting curve) | Solid and liquid regions | Conditions where solid and liquid coexist; the melting/freezing point at each pressure |
| Liquid-gas line (vapor pressure curve) | Liquid and gas regions | Conditions where liquid and gas coexist; the boiling point at each pressure |
| Solid-gas line (sublimation curve) | Solid and gas regions | Conditions where solid and gas coexist; sublimation at each pressure |
Two Special Points
Triple point — The exact temperature and pressure where all three phases coexist simultaneously in equilibrium. Every pure substance has exactly one triple point. For water: 0.01 °C and 611.7 Pa (about 0.006 atm).
Critical point — The temperature and pressure above which the liquid-gas boundary disappears. Above the critical temperature, a substance cannot be liquefied by pressure alone. The substance becomes a supercritical fluid with properties of both liquid and gas. For water: 374 °C and 218 atm.
To use a phase diagram, locate the temperature-pressure point of interest and read which region it falls in. For example, carbon dioxide at 1 atm: moving horizontally from low temperature, the 1 atm line crosses directly from the solid region into the gas region, never entering the liquid region. This confirms that dry ice sublimes at atmospheric pressure. Liquid CO₂ only exists above about 5.11 atm (the triple point pressure of CO₂) — which is why CO₂ fire extinguishers must be pressurized to contain liquid CO₂.
Water's Unusual Solid-Liquid Line
For most substances, the solid-liquid line on the phase diagram slopes to the right — higher pressure raises the melting point. Water is an exception: its solid-liquid line slopes slightly to the left, meaning higher pressure lowers the melting point of ice. This is connected to ice being less dense than liquid water. It means that ice under very high pressure will melt even below 0 °C.
Use this information for questions 1–8: Substance X has a triple point at 5 °C and 0.012 atm, a normal melting point of 5.0 °C (at 1 atm), a normal boiling point of 65 °C (at 1 atm), and a critical point at 200 °C and 75 atm.
- What phase is Substance X at 20 °C and 1 atm?
- What phase is Substance X at −10 °C and 1 atm?
- What phase is Substance X at 80 °C and 1 atm?
- What is happening to Substance X at 5 °C and 0.012 atm?
- At pressures below 0.012 atm, what phase change will solid Substance X undergo when heated?
- Can Substance X exist as a liquid above 200 °C regardless of pressure? Explain.
- A sample of Substance X is at 30 °C and 1 atm. The pressure is slowly decreased to 0.005 atm. Predict what will happen.
- Explain the difference between the triple point and the critical point.
Section 19.9
Phase Changes in Weather, Cooking, and Materials
Weather and Climate
The water cycle is driven almost entirely by phase changes. Evaporation of ocean, lake, and river water adds water vapor to the atmosphere, absorbing roughly 2260 J per gram — energy stored as latent heat in the vapor. When water vapor rises, cools, and condenses around dust particles into cloud droplets, that stored energy is released back into the atmosphere. This release of latent heat fuels thunderstorms and hurricanes.
Frost forms by deposition when the surface temperature drops below 0 °C and air moisture is sufficient. Snow can sublime directly on cold, sunny, dry days — losing mass without ever becoming liquid. High humidity slows evaporation of sweat, making humid heat feel more oppressive than dry heat at the same temperature.
Cooking
Boiling pasta at high altitude takes longer because lower atmospheric pressure reduces the boiling point — water boils at a lower temperature, so it heats food more slowly. A pressure cooker does the opposite: sealed pressure above 1 atm raises the boiling point, often to 120 °C or more, cooking food much faster. Baking involves evaporation of surface moisture (allowing temperatures above 100 °C for browning), and freeze-drying food uses sublimation under low pressure to remove moisture without heat damage, preserving flavor and nutrition.
Materials Science
Casting metals involves melting (solid → liquid) and controlled solidification (liquid → solid) — the cooling rate determines crystal structure and material properties. Glass is made by melting silica and allowing it to solidify as an amorphous solid. Welding joins metals by locally melting and resolidifying them. Steel heat treatment (annealing, quenching, tempering) manipulates the iron-carbon crystal structure through controlled phase transitions, dramatically changing hardness and toughness. Vapor deposition coats surfaces by vaporizing a target material and allowing it to deposit on a cooler surface — used to make anti-reflection lens coatings, mirror reflective layers, and semiconductor devices.
- A large hurricane releases an enormous amount of energy. Where does this energy come from? Name the phase change involved and classify it as endothermic or exothermic.
- Explain why food cooks faster in a pressure cooker than in an open pot of boiling water.
- A ski slope loses snow to sublimation on a cold, clear, dry winter day even though temperatures stay well below 0 °C. Explain this using phase change concepts.
- Vapor deposition creates thin films of aluminum on glass mirrors. Name the two phase changes involved in aluminum (first in the deposition chamber, then on the cool surface) and classify each as endothermic or exothermic.
- A welder joins two steel plates. Name the phase changes that occur in the metal at the weld zone and explain whether each is endothermic or exothermic.
- Why does salting roads in winter help prevent ice from forming? (Salt lowers the freezing point of water below 0 °C — a property called freezing point depression.) Explain in terms of what phase change is being prevented.
Chapter 19 — Back Matter
Chapter Summary
- Sec 19.1Matter exists in three common phases — solid, liquid, gas — differing in particle arrangement, motion, and force strength. Crystalline solids have ordered lattices and sharp melting points; amorphous solids soften gradually.
- Sec 19.2Temperature measures average particle kinetic energy. A distribution of energies explains evaporation below the boiling point. Most substances show thermal expansion; water is denser at 4 °C than at 0 °C, making ice less dense than liquid water.
- Sec 19.3Intermolecular forces — dispersion, dipole-dipole, and hydrogen bonding — determine melting and boiling points. Stronger forces mean higher transition temperatures and a greater tendency to be liquid or solid at room temperature.
- Sec 19.4Melting (solid → liquid) is endothermic; freezing (liquid → solid) is exothermic. Temperature stays constant during both. The heat of fusion (ΔHfus) gives the energy per gram. For water, ΔHfus = 334 J/g.
- Sec 19.5Evaporation is surface vaporization below the boiling point. Vapor pressure increases with temperature. Boiling occurs when vapor pressure equals external pressure, explaining altitude effects. The heat of vaporization (ΔHvap) for water is 2260 J/g. Condensation is exothermic and releases the same amount of energy.
- Sec 19.6Sublimation (solid → gas) and deposition (gas → solid) bypass the liquid phase. Dry ice sublimes at atmospheric pressure. Frost forms by deposition, not by freezing of dew. All six phase changes can be classified as endothermic or exothermic.
- Sec 19.7Heating curves have five regions: three rising segments (warming each phase) and two flat plateaus (melting and boiling). Cooling curves are the reverse. Plateau length reflects the energy of the phase change; slope reflects specific heat of each phase.
- Sec 19.8Phase diagrams map stable phases across temperature-pressure combinations. Boundary lines mark coexistence conditions. The triple point is where all three phases coexist; the critical point marks where the liquid-gas distinction disappears.
- Sec 19.9Phase changes drive Earth's water cycle and weather, underlie cooking and food preservation, and are manipulated deliberately in metal casting, welding, heat treatment, and vapor deposition.
Key Terms
amorphous solid
A solid with a disordered particle arrangement that softens gradually over a range of temperatures rather than melting at one sharp point.
boiling
Vaporization throughout a liquid when its vapor pressure equals external atmospheric pressure.
boiling point
The temperature at which a liquid's vapor pressure equals external pressure. At 1 atm, this is the normal boiling point.
condensation
The phase change from gas to liquid. An exothermic process that releases energy equal to the heat of vaporization.
critical point
The temperature and pressure above which liquid and gas phases are indistinguishable. Above this point, the substance is a supercritical fluid.
crystalline solid
A solid in which particles are arranged in a highly ordered, repeating pattern (crystal lattice). Has a sharp, well-defined melting point.
deposition
The phase change directly from gas to solid, bypassing the liquid phase. An exothermic process. Example: frost formation.
dew point
The temperature at which air becomes saturated with water vapor and condensation begins.
evaporation
Vaporization at the surface of a liquid at temperatures below the boiling point. An endothermic process.
freezing
The phase change from liquid to solid. An exothermic process. For a pure substance, the freezing point equals the melting point.
heat of fusion (ΔHfus)
The energy required to melt one gram of a substance at its melting point without changing temperature. Water: 334 J/g.
heat of vaporization (ΔHvap)
The energy required to vaporize one gram of a substance at its boiling point without changing temperature. Water: 2260 J/g.
heating curve
A graph of temperature versus energy added. Features rising regions (warming) and flat plateaus (phase changes at constant temperature).
intermolecular forces
Attractive forces between molecules — dispersion, dipole-dipole, and hydrogen bonding — that determine state and phase transition temperatures.
latent heat
Energy absorbed or released during a phase change that does not produce a temperature change.
melting
The phase change from solid to liquid. An endothermic process occurring at the melting point.
melting point
The temperature at which a solid becomes a liquid at a given pressure. For a pure substance, equals the freezing point.
phase
A distinct state of matter — solid, liquid, or gas — with uniform composition and properties throughout.
phase change
A transition between phases of matter. Occurs at a specific temperature and pressure; involves absorption or release of energy without a temperature change.
phase diagram
A graph showing which phase of a substance is stable across a range of temperature and pressure conditions.
sublimation
The phase change directly from solid to gas, bypassing the liquid phase. An endothermic process. Example: dry ice.
supercritical fluid
A substance above its critical temperature and pressure, having properties of both liquid and gas simultaneously.
thermal expansion
The tendency of matter to increase in volume as temperature increases, due to greater average particle separation.
triple point
The specific temperature and pressure at which all three phases of a substance coexist simultaneously in equilibrium.
vapor pressure
The pressure exerted by a vapor in equilibrium with its liquid at a given temperature. Increases with temperature.
vaporization
The phase change from liquid to gas, including both evaporation and boiling. An endothermic process.
volatile
Describes a substance with a high vapor pressure at room temperature that evaporates readily.
Formula Review
Energy for Melting or Freezing
q = m × ΔH_fus
q = energy (J) · m = mass (g) · ΔHfus = heat of fusion (J/g)
Melting: q is positive (endothermic) · Freezing: q is negative (exothermic)
Energy for Vaporization or Condensation
q = m × ΔH_vap
q = energy (J) · m = mass (g) · ΔHvap = heat of vaporization (J/g)
Vaporization: q is positive (endothermic) · Condensation: q is negative (exothermic)
| Quantity for Water | Value |
|---|
| Melting/freezing point (at 1 atm) | 0 °C |
| Boiling/condensation point (at 1 atm) | 100 °C |
| Heat of fusion (ΔHfus) | 334 J/g |
| Heat of vaporization (ΔHvap) | 2260 J/g |
| Specific heat of ice | 2.09 J/(g·°C) |
| Specific heat of liquid water | 4.18 J/(g·°C) |
| Specific heat of steam | 1.87 J/(g·°C) |
Worked Example Review
Example A
Energy to melt a sample
- ProblemHow much energy is needed to melt 80.0 g of ice at 0 °C?
- Solve
q = 80.0 g × 334 J/g = 26,720 J = 26.7 kJ
Answer: 26.7 kJ must be absorbed.
Example B
Energy released during freezing
- ProblemHow much energy is released when 250. g of liquid water at 0 °C freezes completely?
- Solve
q = 250. g × 334 J/g = 83,500 J = 83.5 kJ
Answer: 83.5 kJ is released to surroundings.
Example C
Energy to vaporize water
- ProblemHow much energy is required to vaporize 35.0 g of liquid water at 100 °C?
- Solve
q = 35.0 g × 2260 J/g = 79,100 J = 79.1 kJ
Answer: 79.1 kJ must be absorbed.
Example D
Reading a heating curve plateau
- ProblemA heating curve shows a flat region between minute 5 and minute 13, both at 58 °C. What does this represent and what is the melting point?
- ReasonA flat region means temperature is not changing while heat is added — this is a phase change. Since the substance is being heated and a flat region appears, this is the melting plateau.
- ConcludeThe region represents melting. The melting point is 58 °C.
Answer: The substance is melting; its melting point is 58 °C.
Concept Check
- How do the arrangements of particles in a solid, liquid, and gas differ from one another?
- What does temperature measure at the particle level?
- Why does temperature stay constant during a phase change even when heat is being added?
- Explain the difference between evaporation and boiling.
- What is vapor pressure, and how does it change with temperature?
- How does altitude affect the boiling point of water? Explain using vapor pressure.
- What is sublimation? Give one real-world example.
- What is the difference between frost and dew in terms of the phase changes involved?
- On a heating curve, what do the flat regions represent? What do the sloped regions represent?
- What does the triple point on a phase diagram represent?
Additional Practice
- Liquid ethanol evaporates from an open flask left overnight.
- Water droplets appear on a cold glass of water.
- Dry ice on a lab bench slowly shrinks and disappears.
- Molten iron poured into a mold hardens into a solid shape.
- Water vapor high in the atmosphere forms snowflake crystals.
- Iodine crystals produce violet vapor without forming a liquid.
- A puddle freezes overnight as temperature falls below 0 °C.
- Liquid nitrogen spills from a container and immediately turns to vapor.
- Water in a pot is heated until it boils vigorously.
- Frost appears on a car windshield on a cold, clear night.
- How much energy is needed to melt 120. g of ice at 0 °C?
- How much energy is released when 75.0 g of liquid water freezes at 0 °C?
- How much energy is required to vaporize 22.0 g of water at 100 °C?
- How much energy is released when 40.0 g of steam condenses to liquid water at 100 °C?
- Benzene has a heat of vaporization of 394 J/g. How much energy is needed to vaporize 30.0 g of liquid benzene?
- A student adds 10.0 g of ice to a beverage. How many joules does the ice absorb as it melts at 0 °C?
- A chemical process requires 500. g of water vapor to condense to liquid. How many kilojoules are released?
- How many grams of ice can be melted by 50,000 J of energy?
Error Analysis
Error 1 — Formula misuse
q = 40.0 g + 334 J/g = 374 J
What is the error in this calculation for melting 40.0 g of ice? What is the correct setup and answer?
Error 2 — Wrong value used
Energy to vaporize 50.0 g of water: q = 50.0 g × 334 J/g = 16,700 J
What is the error? What constant should be used for vaporization, and what is the correct answer?
Error 3 — Temperature during phase change
"When steam condenses to water, the water's temperature rises because energy is being released."
What is wrong with this statement? Describe correctly what happens to temperature and energy during condensation.
Error 4 — Heating curve interpretation
"The flat region on the heating curve means no heat is being added to the sample."
Correct this misconception. What is actually happening in a flat region? Where is the added energy going?
Error 5 — Frost vs. dew
"Frost is just dew that has frozen solid."
Use phase change terminology to correct this statement. How does frost actually form?
Error 6 — Dry ice behavior
"Dry ice melts into liquid CO₂ at atmospheric pressure before turning into gas."
Is this correct? What actually happens to dry ice at atmospheric pressure? What would have to change for liquid CO₂ to form?
Chapter 19 Checkpoint
Chapter 19 Checkpoint
Complete without assistance — assesses all key skills and concepts from Chapter 19
Section 1 — Vocabulary and Concepts
- Explain the difference between a crystalline solid and an amorphous solid. Give one example of each and predict how each behaves as it is heated toward its melting temperature.
- Define vapor pressure and explain in two sentences how it determines the boiling point of a liquid.
- What is the triple point of a substance? What makes it unique compared to other points on a phase diagram?
Section 2 — Phase Change Identification
- Name the phase change: water vapor deposits as ice crystals on a cold metal surface.
- Name the phase change: solid iodine produces a violet vapor above the crystal surface.
- Name the phase change and classify as endothermic or exothermic: liquid water becomes steam in a boiling pot.
- Name the phase change and classify as endothermic or exothermic: molten lava solidifies into rock.
Section 3 — Calculations
- How much energy is required to melt 65.0 g of ice at 0 °C? (ΔHfus = 334 J/g)
- How much energy is released when 30.0 g of steam at 100 °C condenses to liquid water? (ΔHvap = 2260 J/g)
- A sample absorbs 83,500 J while vaporizing completely at its boiling point. The heat of vaporization is 835 J/g. What is the mass of the sample?
Section 4 — Heating Curve
- A heating curve for a pure substance shows: temperature rising from −20 °C to 45 °C, then staying at 45 °C, then rising to 130 °C, then staying at 130 °C, then rising above 130 °C. What are the melting and boiling points of this substance? What phase is it in at 80 °C?
- The boiling plateau in that same heating curve is much longer than the melting plateau. What does this reveal about the relative sizes of ΔHfus and ΔHvap for this substance?
Section 5 — Application
- A hiker cooking pasta at 3000 m altitude notices the water boils at about 90 °C and the pasta takes much longer to cook than at home. Explain using vapor pressure and boiling point concepts.
- Explain why steam burns at 100 °C are typically more severe than liquid water burns at 100 °C. Name the phase change involved and calculate how much energy 5.0 g of steam releases when it condenses.
Mixed Review
These questions draw on content from Chapters 12, 18, and 19. They are designed to keep earlier learning connected to new topics.
- (Chapter 12 review) Water has a higher boiling point than hydrogen sulfide (H₂S), even though H₂S is a larger molecule. Explain using intermolecular forces.
- (Chapter 18 review) A sealed container of gas is cooled at constant volume. According to the ideal gas law, what happens to the pressure?
- (Chapter 19) Explain why the boiling point of water decreases at higher altitude.
- (Chapters 18 and 19) A gas is compressed until it liquefies. What phase change has occurred? Is this change endothermic or exothermic?
- (Chapter 12 review) Pentane (C₅H₁₂) is a nonpolar hydrocarbon. What type of intermolecular force holds liquid pentane together? Predict whether pentane has a higher or lower boiling point than water, and explain why.
- (Chapter 19) A scientist wants to determine if an unknown sample is a pure substance or a mixture. She heats it carefully and watches the temperature. If it is pure, what should the heating curve show at the melting and boiling points?
- (Chapter 18 review) If temperature (in Kelvin) doubles at constant pressure, what happens to gas volume?
- (Chapter 19) A cooling curve for a substance shows a very short condensation plateau and a much longer freezing plateau. What can you conclude about the heats of vaporization and fusion of this substance?
Looking Ahead
Chapter 20: Thermochemistry
In Chapter 19, you studied the energy absorbed and released during phase changes. You calculated the energy needed to melt or vaporize a substance using the heat of fusion and the heat of vaporization. You saw that temperature stays constant during a phase change because all the energy goes into changing the state rather than changing particle speed.
In Chapter 20, you will extend your understanding of energy to the energy involved when temperature does change — the energy absorbed or released as substances heat up or cool down. You will study specific heat, which explains why some substances heat up faster than others, and calorimetry, the laboratory technique for measuring heat flow. You will also study enthalpy change in chemical reactions — the energy absorbed or released when chemical bonds break and form — and how to read energy diagrams that map the energy landscape of a reaction. Chapter 20 brings temperature, heat, and chemistry into a complete picture of thermochemistry.
MGU Library Connection
- Chapter 12: Molecular Geometry and Intermolecular Forces — prerequisite for Section 19.3
- Chapter 18: Gases and Gas Laws — prerequisite for vapor pressure and boiling point concepts
- Chapter 20: Thermochemistry — extends energy concepts from this chapter to specific heat and enthalpy
- Appendix B: Measurement and Laboratory Reference — temperature scales and measurement equipment
- MGU Dictionary entries: phase, phase change, sublimation, vapor pressure, triple point, enthalpy of fusion, enthalpy of vaporization