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Phase Changes of Matter (Phase Transitions)

A phase change or phase transition is a change between solid, liquid, gaseous, and sometimes plasma states of matter . The states of matter differ in the organization of particles and their energy. The main factors that cause phase changes are changes in temperature and pressure. At the phase transition, such as the boiling point between liquid and gas phases, the two states of matter have identical free energies and are equally likely to exist.

Here is a list of the key phase changes between solids, liquids, gases, and plasma. There are six phase changes between solids , liquids, and gases , and eight phase changes if you include plasma. There are additional phase changes if you explore condensed matter physics or metallurgy.

List of Phase Changes

Here is a list of the phase changes of matter.

• Melting (Solid → Liquid)
• Freezing (Liquid → Solid)
• Vaporization or Evaporation (Liquid → Gas)
• Condensation (Gas → Liquid)
• Deposition (Gas → Solid)
• Sublimation (Solid → Gas)
• Ionization (Gas → Plasma)
• Deionization or Recombination (Plasma → Gas)

Phase Changes for States of Matter

Another way to learn phase changes is to associate them with the starting state of matter:

• Solid : A solid can melt into liquid or sublimate into gas.
• Liquid : A liquid can freeze into a solid or vaporize into a gas.
• Gas : A gas can deposit into a solid, condense into a liquid, or ionize into plasma.
• Plasma : Plasma can deionize or recombine to form a gas. Remember, plasma is like a gas, except the particles are even further apart and they are ionized.

Examples of Phase Changes

• Melting : Solid ice melts into liquid water.
• Freezing : Freezing water changes it from a liquid into solid ice.
• Vaporization : An example of vaporization is the evaporation of rubbing alcohol from skin into the air.
• Condensation : A good example of condensation is dew formation from water vapor in air.
• Deposition : Hoarfrost is grayish-white frost that forms during clear, cold weather when water vapor deposits as ice. Another example is deposition of silver vapor onto glass to form a silver mirror.
• Sublimation : Dry ice undergoes sublimation to change from solid carbon dioxide directly into carbon dioxide gas. Another example is the transition from ice directly into water vapor on a cold, windy winter day.
• Ionization : When you turn on a plasma ball toy, the noble gases inside are ionized by an electric charge and become plasma. The aurora is another example of ionization.
• Deionization or Recombination : Lightning is an example of plasma. After a lightning strike, nitrogen ions eventually draw closer together and lose their charge to become N 2 gas.

Why Phase Changes Occur

Most phase changes occur because of a change in the energy of the system. Increasing temperature gives atoms and molecules more kinetic energy, helping them break bonds and move further apart. Similarly, decreasing temperature slows down particles and makes it easier for them to gain rigid structure. Increasing pressure forces particle together, while decreasing pressure lets them move away from each other. You can use a phase diagram to predict whether a substance will be a solid, liquid, or gas at a given combination of temperature and pressure. Matter must be ionized to become plasma. So, you can increase temperature to form ions, but decreasing pressure doesn’t automatically make plasma even if you go all the way to a vacuum.

• Blundell, Stephen J.; Katherine M. Blundell (2008). Concepts in Thermal Physics . Oxford University Press. ISBN 978-0-19-856770-7.
• IUPAC (1997). “Phase Transition”. Compendium of Chemical Terminology (2nd ed.) (the “Gold Book”). ISBN 0-9678550-9-8. doi: 10.1351/goldbook
• Jaeger, Gregg (1 May 1998). “The Ehrenfest Classification of Phase Transitions: Introduction and Evolution”.  Archive for History of Exact Sciences .  53  (1): 51–81. doi: 10.1007/s004070050021

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Key Questions

Heat decides the state of matter. Heat plays an important role in converting one state of matter to another. Adding heat or taking off heat brings in change of state/ phase.

Soild -----> Liquid This is called melting process ( solid needs to be heated, example melting of butter, solid on taking up heat changes to liquid)

Liquid ----> Solid This is called freezing process ( liquid needs to be cooled off, example freezing of water, liquid on losing heat to surroundings cools off to form soild.)

Liquid ---> Gas This is called evaporation process ( liquid needs to be heated, example evaporation of water, water on taking up heat from the flame or from the surroundings changes to gas)

Gas ---> Liquid This is called condensation process ( gas needs to be cooled, example condensation of steam on the mirror, when one throws steam on mirror or cold surface, steam or gas loses heat to the surface and changes to gas)

Soild -----> Gas This is called sublimation process ( solid needs to be heated, example leaving dry Ice on the surface, solid dry ice gains heat from surroundings and turns into gas.

Phase Changes

Phase changes and energy conservation.

During a phase transition, certain properties of the medium change, often discontinuously, as a result of some external condition.

Learning Objectives

Describe behavior of the medium during a phase transition

Key Takeaways

• The term is most commonly used to describe transitions between solid, liquid and gaseous states of matter and, in rare cases, plasma.
• Once water reaches the boiling point, extra energy is used to change the state of matter and increase the potential energy instead of the kinetic energy.
• Plots of pressure versus temperatures, an example of a phase diagram, provide considerable insight into thermal properties of substances.
• intermolecular : from one molecule to another; between molecules
• plasma : a state of matter consisting of partially ionized gas
• thermodynamic : Relating to the conversion of heat into other forms of energy.

A phase of a thermodynamic system and the states of matter have uniform physical properties. During a phase transition of a given medium certain properties of the medium change, often discontinuously, as a result of some external condition, such as temperature or pressure. For example, a liquid may become gas upon heating to the boiling point, resulting in an abrupt change in volume. The measurement of the external conditions at which the transformation occurs is termed the phase transition. The term is most commonly used to describe transitions between solid, liquid and gaseous states of matter and, in rare cases, plasma.

As an example, if you boil water, it never goes above 100 degrees Celsius. Only after it has completely evaporated will it get any hotter. This is because once water reaches the boiling point, extra energy is used to change the state of matter and increase the potential energy instead of the kinetic energy. The opposite happens when water freezes. To boil or melt one mole of a substance, a certain amount of energy is required. These amounts of energy are the molar heat of vaporization and molar heat of fusion. If that amount of energy is added to a mole of that substance at boiling or freezing point, all of it will melt or boil, but the temperature won’t change.

Temperature increases linearly with heat, until the melting point. But the heat added does not change the temperature; that heat energy is instead used to break intermolecular bonds and convert ice into water. At this point, there is a mixture of both ice and water. Once all ice has been melted, the temperature again rises linearly with heat added. At the boiling point, temperature no longer rises with heat added because the energy is once again being used to break intermolecular bonds. Once all water has been boiled to steam, the temperature will continue to rise linearly as heat is added.

Temperature vs. Heat : This graph shows the temperature of ice as heat is added.

The plots of pressure versus temperatures provide considerable insight into thermal properties of substances. There are well-defined regions on these graphs that correspond to various phases of matter, so PT graphs are called phase diagrams. Using the graph, if you know the pressure and temperature you can determine the phase of water. The solid lines—boundaries between phases—indicate temperatures and pressures at which the phases coexist (that is, they exist together in ratios, depending on pressure and temperature). For example, the boiling point of water is 100º C at 1.00 atm. As the pressure increases, the boiling temperature rises steadily to 374º C at a pressure of 218 atm. A pressure cooker (or even a covered pot) will cook food faster because the water can exist as a liquid at temperatures greater than 100º C without all boiling away. The curve ends at a point called the critical point, because at higher temperatures the liquid phase does not exist at any pressure. The critical temperature for oxygen is -118ºC, so oxygen cannot be liquefied above this temperature.

Phase Diagram of Water : In this typical phase diagram of water, the green lines mark the freezing point, and the blue line marks the boiling point, showing how they vary with pressure. The dotted line illustrates the anomalous behavior of water. Note that water changes states based on the pressure and temperature.

Humidity, Evaporation, and Boiling

The amount of water vapor in air is a result of evaporation or boiling, until an equilibrium is reached.

Explain why water boils at 100 °C

• Relative humidity is the fraction of water vapor in a gas compared to the saturation value.
• Since the kinetic energy of a molecule is proportional to its temperature, evaporation proceeds more quickly at higher temperatures.
• Vapor pressure increases with temperature because molecular speeds are higher as temperature increases.
• Water boils at 100 °C because the vapor pressure exceeds atmospheric pressure at this temperature.
• equilibrium : The state of a body at rest or in uniform motion, the resultant of all forces on which is zero.
• vapor pressure : The pressure that a vapor exerts, or the partial pressure if it is mixed with other gases.
• humidity : The amount of water vapor in the air.

The term relative humidity refers to how much water vapor is in the air compared with the maximum possible. At its maximum, denoted as saturation, the relative humidity is 100%, and evaporation is inhibited. The amount of water vapor the air can hold depends on its temperature. For example, relative humidity rises in the evening, as air temperature declines, sometimes reaching the dew point. At the dew point temperature, relative humidity is 100%, and fog may result from the condensation of water droplets if they are small enough to stay in suspension. Conversely, if one wished to dry something, it is more effective to blow hot air over it rather than cold air, because, among other things, hot air can hold more water vapor.

Evaporation

The capacity of air to hold water vapor is based on vapor pressure of water. The liquid and solid phases are continuously giving off vapor because some of the molecules have high enough speeds to enter the gas phase, a process called evaporation; see (a). For the molecules to evaporate, they must be located near the surface, be moving in the proper direction, and have sufficient kinetic energy to overcome liquid-phase intermolecular forces. When only a small proportion of the molecules meet these criteria, the rate of evaporation is low. Since the kinetic energy of a molecule is proportional to its temperature, evaporation proceeds more quickly at higher temperatures.

If a lid is placed over the container, as in (b), evaporation continues, increasing the pressure, until sufficient vapor has built up for condensation to balance evaporation. Then equilibrium has been achieved, and the vapor pressure is equal to the partial pressure of water in the container. Vapor pressure increases with temperature because molecular speeds are higher as temperature increases.

As the faster-moving molecules escape, the remaining molecules have lower average kinetic energy, and the temperature of the liquid decreases. This phenomenon is also called evaporative cooling. This is why evaporating sweat cools the human body. Evaporation also tends to proceed more quickly with higher flow rates between the gaseous and liquid phase and in liquids with higher vapor pressure. For example, laundry on a clothes line will dry (by evaporation) more rapidly on a windy day than on a still day.

Application for Boiling Water

Why does water boil at 100ºC? The vapor pressure of water at 100ºC is 1.01×10 5 Pa, or 1.00 atm. Thus, it can evaporate without limit at this temperature and pressure. But why does it form bubbles when it boils? This is because water ordinarily contains significant amounts of dissolved air and other impurities, which are observed as small bubbles of air in a glass of water. If a bubble starts out at the bottom of the container at 20ºC, it contains water vapor (about 2.30%). The pressure inside the bubble is fixed at 1.00 atm (we ignore the slight pressure exerted by the water around it). As the temperature rises, the amount of air in the bubble stays the same, but the water vapor increases; the bubble expands to keep the pressure at 1.00 atm. At 100ºC, water vapor enters the bubble continuously since the partial pressure of water is equal to 1.00 atm in equilibrium. It cannot reach this pressure, however, since the bubble also contains air and total pressure is 1.00 atm. The bubble grows in size and thereby increases the buoyant force. The bubble breaks away and rises rapidly to the surface, resulting in boiling. (See. )

Close-up of the Boiling Process : (a) An air bubble in water starts out saturated with water vapor at 20ºC. (b) As the temperature rises, water vapor enters the bubble because its vapor pressure increases. The bubble expands to keep its pressure at 1.00 atm. (c) At 100ºC, water vapor enters the bubble continuously because water’s vapor pressure exceeds its partial pressure in the bubble, which must be less than 1.00 atm. The bubble grows and rises to the surface.

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Phase change.

A phase change is when matter changes to from one state (solid, liquid, gas, plasma) to another. (see figure 1). These changes occur when sufficient energy is supplied to the system (or a sufficient amount is lost), and also occur when the pressure on the system is changed. The temperatures and pressures under which these changes happen differ depending on the chemical and physical properties of the system. The energy associated with these transitions is called latent heat .

Water is a substance that has many interesting properties that influence its phase changes. Most people learn from an early age that water melts from ice to liquid at 0°C, and boils from a liquid to a gas at 100°C; but this isn't true in all circumstances. The pressure affects these transition points, so for water, the boiling point actually decreases as the pressure decreases. Water also has certain intermolecular forces which govern the temperatures at which these transitions occur. [2] This difference in boiling point is why the directions for cooking at high altitudes are sometimes slightly different (like boiling pasta longer).

The relatively large amount of energy needed to change the phase of water is one of the reasons why water is used to cool power plants . It's also part of why humans sweat in order to stay cool (through evaporation) and dogs pant. This high latent heat also makes water important for moderating the climate .

Boiling/condensing and freezing/melting are the most common pairs of phase changes experienced on Earth. However, there are other phase changes such as sublimation —which is going straight from a solid to a gas. Figure 1 also shows phase changes that are rare (on Earth, at least) known as plasma . However, figure 1 does not show what happens when gases or liquids get to sufficiently high pressures and temperatures that they can't be distinguished. This phase is called the supercritical fluid state (which is useful for some modern power plants).

Visit UC Davis' Chem wiki for more information on phase changes and other chemical phenomena.

For Further Reading

• Latent heat
• Boiling point
• Supercritical fluid
• Water cycle
• Or explore a random page
• ↑ Wikimedia Commons [Public Domain], Available: http://commons.wikimedia.org/wiki/File%3APhase_change_-_en.svg
• ↑ UC Davis Chem Wiki [Online] Available: http://chemwiki.ucdavis.edu/Physical_Chemistry/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Intermolecular_Forces/Van_der_Waals_Forces

13.5 Phase Changes

Up to now, we have considered the behavior of ideal gases. Real gases are like ideal gases at high temperatures. At lower temperatures, however, the interactions between the molecules and their volumes cannot be ignored. The molecules are very close (condensation occurs) and there is a dramatic decrease in volume, as seen in Figure 13.26 . The substance changes from a gas to a liquid. When a liquid is cooled to even lower temperatures, it becomes a solid. The volume never reaches zero because of the finite volume of the molecules.

High pressure may also cause a gas to change phase to a liquid. Carbon dioxide, for example, is a gas at room temperature and atmospheric pressure, but becomes a liquid under sufficiently high pressure. If the pressure is reduced, the temperature drops and the liquid carbon dioxide solidifies into a snow-like substance at the temperature – 78 º C – 78 º C size 12{ +- "78"°C} {} . Solid CO 2 CO 2 size 12{"CO" rSub { size 8{2} } } {} is called “dry ice.” Another example of a gas that can be in a liquid phase is liquid nitrogen ( LN 2 ) ( LN 2 ) size 12{ \( "LN" rSub { size 8{2} } \) } {} . LN 2 LN 2 size 12{"LN" rSub { size 8{2} } } {} is made by liquefaction of atmospheric air (through compression and cooling). It boils at 77 K ( – 196 º C ) ( – 196 º C ) size 12{ \( –"196"°C \) } {} at atmospheric pressure. LN 2 LN 2 size 12{"LN" rSub { size 8{2} } } {} is useful as a refrigerant and allows for the preservation of blood, sperm, and other biological materials. It is also used to reduce noise in electronic sensors and equipment, and to help cool down their current-carrying wires. In dermatology, LN 2 LN 2 size 12{"LN" rSub { size 8{2} } } {} is used to freeze and painlessly remove warts and other growths from the skin.

PV Diagrams

We can examine aspects of the behavior of a substance by plotting a graph of pressure versus volume, called a PV diagram . When the substance behaves like an ideal gas, the ideal gas law describes the relationship between its pressure and volume. That is,

Now, assuming the number of molecules and the temperature are fixed,

For example, the volume of the gas will decrease as the pressure increases. If you plot the relationship PV = constant PV = constant size 12{ size 11{ ital "PV"="constant"}} {} on a PV PV size 12{ ital "PV"} {} diagram, you find a hyperbola. Figure 13.27 shows a graph of pressure versus volume. The hyperbolas represent ideal-gas behavior at various fixed temperatures, and are called isotherms . At lower temperatures, the curves begin to look less like hyperbolas—the gas is not behaving ideally and may even contain liquid. There is a critical point —that is, a critical temperature —above which liquid cannot exist. At sufficiently high pressure above the critical point, the gas will have the density of a liquid but will not condense. Carbon dioxide, for example, cannot be liquefied at a temperature above 31 . 0 º C 31 . 0 º C size 12{"31" "." 0°C} {} . Critical pressure is the minimum pressure needed for liquid to exist at the critical temperature. Table 13.3 lists representative critical temperatures and pressures.

Phase Diagrams

The plots of pressure versus temperatures provide considerable insight into thermal properties of substances. There are well-defined regions on these graphs that correspond to various phases of matter, so PT PT size 12{ ital "PT"} {} graphs are called phase diagrams . Figure 13.28 shows the phase diagram for water. Using the graph, if you know the pressure and temperature you can determine the phase of water. The solid lines—boundaries between phases—indicate temperatures and pressures at which the phases coexist (that is, they exist together in ratios, depending on pressure and temperature). For example, the boiling point of water is 100 º C 100 º C size 12{"100"°C} {} at 1.00 atm. As the pressure increases, the boiling temperature rises steadily to 374 º C 374 º C size 12{"374"°C} {} at a pressure of 218 atm. A pressure cooker (or even a covered pot) will cook food faster because the water can exist as a liquid at temperatures greater than 100 º C 100 º C size 12{"100"°C} {} without all boiling away. The curve ends at a point called the critical point , because at higher temperatures the liquid phase does not exist at any pressure. The critical point occurs at the critical temperature, as you can see for water from Table 13.3 . The critical temperature for oxygen is – 118 º C – 118 º C size 12{ +- "118"°C} {} , so oxygen cannot be liquefied above this temperature.

Similarly, the curve between the solid and liquid regions in Figure 13.28 gives the melting temperature at various pressures. For example, the melting point is 0 º C 0 º C size 12{0°C} {} at 1.00 atm, as expected. Note that, at a fixed temperature, you can change the phase from solid (ice) to liquid (water) by increasing the pressure. Ice melts from pressure in the hands of a snowball maker. From the phase diagram, we can also say that the melting temperature of ice falls with increased pressure. When a car is driven over snow, the increased pressure from the tires melts the snowflakes; afterwards the water refreezes and forms an ice layer.

At sufficiently low pressures there is no liquid phase, but the substance can exist as either gas or solid. For water, there is no liquid phase at pressures below 0.00600 atm. The phase change from solid to gas is called sublimation . It accounts for large losses of snow pack that never make it into a river, the routine automatic defrosting of a freezer, and the freeze-drying process applied to many foods. Carbon dioxide, on the other hand, sublimates at standard atmospheric pressure of 1 atm. (The solid form of CO 2 CO 2 size 12{"CO" rSub { size 8{2} } } {} is known as dry ice because it does not melt. Instead, it moves directly from the solid to the gas state.)

All three curves on the phase diagram meet at a single point, the triple point , where all three phases exist in equilibrium. For water, the triple point occurs at 273.16 K ( 0 . 01 º C ) ( 0 . 01 º C ) size 12{ \( 0 "." "01"°C \) } {} , and is a more accurate calibration temperature than the melting point of water at 1.00 atm, or 273.15 K ( 0 . 0 º C ) ( 0 . 0 º C ) size 12{ \( 0 "." 0°C \) } {} . See Table 13.4 for the triple point values of other substances.

Equilibrium

Liquid and gas phases are in equilibrium at the boiling temperature. (See Figure 13.29 .) If a substance is in a closed container at the boiling point, then the liquid is boiling and the gas is condensing at the same rate without net change in their relative amount. Molecules in the liquid escape as a gas at the same rate at which gas molecules stick to the liquid, or form droplets and become part of the liquid phase. The combination of temperature and pressure has to be “just right”; if the temperature and pressure are increased, equilibrium is maintained by the same increase of boiling and condensation rates.

One example of equilibrium between liquid and gas is that of water and steam at 100 º C 100 º C size 12{"100"°C} {} and 1.00 atm. This temperature is the boiling point at that pressure, so they should exist in equilibrium. Why does an open pot of water at 100 º C 100 º C size 12{"100"°C} {} boil completely away? The gas surrounding an open pot is not pure water: it is mixed with air. If pure water and steam are in a closed container at 100 º C 100 º C size 12{"100"°C} {} and 1.00 atm, they would coexist—but with air over the pot, there are fewer water molecules to condense, and water boils. What about water at 20 . 0 º C 20 . 0 º C size 12{"20" "." 0°C} {} and 1.00 atm? This temperature and pressure correspond to the liquid region, yet an open glass of water at this temperature will completely evaporate. Again, the gas around it is air and not pure water vapor, so that the reduced evaporation rate is greater than the condensation rate of water from dry air. If the glass is sealed, then the liquid phase remains. We call the gas phase a vapor when it exists, as it does for water at 20 . 0 º C 20 . 0 º C size 12{"20" "." 0°C} {} , at a temperature below the boiling temperature.

Check Your Understanding

Explain why a cup of water (or soda) with ice cubes stays at 0 º C 0 º C size 12{0°C} {} , even on a hot summer day.

The ice and liquid water are in thermal equilibrium, so that the temperature stays at the freezing temperature as long as ice remains in the liquid. (Once all of the ice melts, the water temperature will start to rise.)

Vapor Pressure, Partial Pressure, and Dalton’s Law

Vapor pressure is defined as the pressure at which a gas coexists with its solid or liquid phase. Vapor pressure is created by faster molecules that break away from the liquid or solid and enter the gas phase. The vapor pressure of a substance depends on both the substance and its temperature—an increase in temperature increases the vapor pressure.

Partial pressure is defined as the pressure a gas would create if it occupied the total volume available. In a mixture of gases, the total pressure is the sum of partial pressures of the component gases , assuming ideal gas behavior and no chemical reactions between the components. This law is known as Dalton’s law of partial pressures , after the English scientist John Dalton (1766–1844), who proposed it. Dalton’s law is based on kinetic theory, where each gas creates its pressure by molecular collisions, independent of other gases present. It is consistent with the fact that pressures add according to Pascal’s Principle . Thus water evaporates and ice sublimates when their vapor pressures exceed the partial pressure of water vapor in the surrounding mixture of gases. If their vapor pressures are less than the partial pressure of water vapor in the surrounding gas, liquid droplets or ice crystals (frost) form.

Is energy transfer involved in a phase change? If so, will energy have to be supplied to change phase from solid to liquid and liquid to gas? What about gas to liquid and liquid to solid? Why do they spray the orange trees with water in Florida when the temperatures are near or just below freezing?

Yes, energy transfer is involved in a phase change. We know that atoms and molecules in solids and liquids are bound to each other because we know that force is required to separate them. So in a phase change from solid to liquid and liquid to gas, a force must be exerted, perhaps by collision, to separate atoms and molecules. Force exerted through a distance is work, and energy is needed to do work to go from solid to liquid and liquid to gas. This is intuitively consistent with the need for energy to melt ice or boil water. The converse is also true. Going from gas to liquid or liquid to solid involves atoms and molecules pushing together, doing work and releasing energy.

PhET Explorations

States of matter—basics.

Heat, cool, and compress atoms and molecules and watch as they change between solid, liquid, and gas phases.

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December 15, 2021

Understanding phase change materials for thermal energy storage

by American Institute of Physics

As the world searches for practical ways to decarbonize our activities and mitigate associated climate change, approaches to alternative energy are hampered by the intermittent nature of energy sources, such as solar and wind. One possible solution to help boost reliability and adoption of such renewable energy sources is improved energy storage capabilities.

In the Journal of Applied Physics , researchers from Lawrence Berkeley National Laboratory, Georgia Institute of Technology, and the University of California, Berkeley, describe advances in understanding the fundamental physics of phase change materials used for energy storage.

Phase change materials absorb thermal energy as they melt, holding that energy until the material is again solidified. Better understanding the liquid state physics of this type of thermal storage may help accelerate technology development for the energy sector.

"Modeling the physics of gases and solids is easier than liquids," said co-author Ravi Prasher. "Gases are free moving, and solids merely vibrate, but liquids behave more like a solid when melting and more like a vapor as they heat up."

This behavior makes it difficult to model and predict storage-system behavior during the phase change critical to its function.

To best capitalize on phase change phenomena of materials for thermal storage, material parameters, including molecular motion and entropy, must be mathematically described, so behavior and theoretical limits can be predicted. The researchers describe a step toward this predictive power by discussing past literature and new developments in the field of liquid state physics.

"The amount of energy that gets stored during phase change depends on the entropy of melting," said Prasher. "Once you know how to predict the entropy change, you know how to design materials that will cater to specific needs."

Developing high-performance thermal energy storage material is important, as heat energy dominates energy use in buildings and manufacturing. Thermal storage is also safer than many other forms of energy storage, since it does not have the capability to release stored energy rapidly and destructively in the case of a malfunction.

Finally, thermal storage holds promise for functioning at large scales and over long durations, and individualized and/or novel materials can be manufactured to suite specific needs. More sophisticated models are needed to further aid in the ability to screen and create materials for optimal thermal energy storage applications.

Journal information: Journal of Applied Physics

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• Science Clarified
• Real-Life Chemistry Vol 1
• Properties of Matter

Properties of Matter - Real-life applications

Types of solids.

Particles of solids resist attempts to compress them, or push them together, and because of their close proximity, solid particles are fixed in an orderly and definite pattern. As a result, a solid usually has a definite volume and shape.

A crystalline solid is a type of solid in which the constituent parts are arranged in a simple, definite geometric pattern that is repeated in all directions. But not all crystalline solids are the same. Table salt is an example of an ionic solid: a form of crystalline solid that contains ions. When mixed with a solvent such as water, ions from the salt move freely throughout the solution, making it possible to conduct an electric current.

Regular table sugar (sucrose) is a molecular solid, or one in which the molecules have a neutral electric charge—that is, there are no ions present. Therefore, a solution of water and sugar would not conduct electricity. Finally, there are crystalline solids known as atomic solids, in which atoms of one element bond to one another. Examples include diamonds (made of pure carbon), silicon, and all metals.

Other solids are said to be amorphous, meaning that they possess no definite shape. Amorphous solids—an example of which is clay—either possess very tiny crystals, or consist of several varieties of crystal mixed randomly. Still other solids, among them glass, do not contain crystals.

Freezing and Melting

Vibrations and freezing..

Because of their slow movement in relation to one another, solid particles exert strong attractions; yet as slowly as they move, solid particles do move—as is the case with all forms of matter at the atomic level. Whereas the particles in a liquid or gas move fast enough to be in relative motion with regard to one another, however, solid particles merely vibrate from a fixed position.

As noted earlier, the motion and attraction of particles in matter has a direct effect on thermal energy, and thus on heat and temperature. The cooler the solid, the slower and weaker the vibrations, and the closer the particles are to one another. Thus, most types of matter contract when freezing, and their density increases. Absolute zero, or 0K on the Kelvin scale of temperature—equal to −459.67°F (−273°C)—is the point at which vibration virtually ceases.

Note that the vibration virtually stops, but does not totally stop. In fact, as established in the third law of thermodynamics, absolute zero is impossible to achieve: thus, the relative motion of molecules never ceases. The lowest temperature actually achieved, at a Finnish nuclear laboratory in 1993, is 2.8 · 10 −10 K, or 0.00000000028K—still above absolute zero.

UNUSUAL CHARACTERISTICS OF SOLID AND LIQUID WATER.

The behavior of water when frozen is interesting and exceptional. Above 39.2°F (4°C) water, like most substances, expands when heated. In other words, the molecules begin moving further apart as expected, because—in this temperature range, at least—water behaves like other substances, becoming "less solid" as the temperature increases.

Between 32°F (0°C) and 39.2°F (4°C), however, water actually contracts. In this temperature range, it is very "cold" (that is, it has relatively little heat), but it is not frozen. The density of water reaches its maximum—in other words, water molecules are as closely packed as they can be—at 39.2°F; below that point, the density starts to decrease again. This is highly unusual: in most substances, the density continues to increase with lowered temperatures, whereas water is actually most dense slightly above the freezing point.

Below the freezing point, then, water expands, and therefore when water in pipes freezes, it may increase in volume to the point where it bursts the pipe. This is also the reason why ice floats on water: its weight is less than that of the water it has displaced, and thus it is buoyant. Additionally, the buoyant qualities of ice atop very cold water helps explain the behavior of lake water in winter; although the top of a lake may freeze, the entire lake rarely freezes solid—even in the coldest of inhabited regions.

Instead of freezing from the bottom up, as it would if ice were less buoyant than the water, the lake freezes from the top down—an important thing to remember when ice-fishing! Furthermore, water in general (and ice in particular) is a poor conductor of heat, and thus little of the heat from the water below it escapes. Therefore, the lake does not freeze completely—only a layer at the top—and this helps preserve animal and plant life in the body of water.

When heated, particles begin to vibrate more and more, and therefore move further apart. If a solid is heated enough, it loses its rigid structure and becomes a liquid. The temperature at which a solid turns into a liquid is called the melting point, and melting points are different for different substances. The melting point of a substance, incidentally, is the same as its freezing point: the difference is a matter of orientation—that is, whether the process is one of a solid melting to become a liquid, or of a liquid freezing to become a solid.

The energy required to melt 1 mole of a solid substance is called the molar heat of fusion. It can be calculated by the formula Q = smδT, where Q is energy, s is specific heat capacity, m is mass, and δT means change in temperature. (In the symbolic language often employed by scientists, the Greek letter δ, or delta, stands for "change in.") Specific heat capacity is measured in units of J/g · °C (joules per gram-degree Celsius), and energy in joules or kilojoules (kJ)—that is, 1,000 joules.

In melting, all the thermal energy in a solid is used in breaking up the arrangement of crystals, called a lattice. This is why water melted from ice does not feel any warmer than the ice did: the thermal energy has been expended, and there is none left over for heating the water. Once all the ice is melted, however, the absorbed energy from the particles—now moving at much greater speeds than when the ice was in a solid state—causes the temperature to rise.

For the most part, solids composed of particles with a higher average atomic mass require more energy—and hence higher temperatures—to induce the vibrations necessary for melting. Helium, with an average atomic mass of 4.003 amu, melts or freezes at an incredibly low temperature: −457.6°F (−272°C), or close to absolute zero. Water, for which, as noted earlier, the average atomic mass is the sum of the masses for its two hydrogen atoms and one oxygen atom, has an average molecular mass of 18.016 amu. Ice melts (or water freezes) at much higher temperatures than helium: 32°F (0°C). Copper, with an average atomic mass of 63.55 amu, melts at much, much higher temperatures than water: 1,985°F (1,085°C).

The particles of a liquid, as compared to those of a solid, have more energy, more motion, and—generally speaking—less attraction to one another. The attraction, however, is still fairly strong: thus, liquid particles are in close enough proximity that the liquid resists attempts at compression.

On the other hand, their arrangement is loose enough that the particles tend to move around one another rather than simply vibrate in place the way solid particles do. A liquid is therefore not definite in shape. Due to the fact that the particles in a liquid are farther apart than those of a solid, liquids tend to be less dense than solids. The liquid phase of a substance thus tends to be larger in volume than its equivalent in solid form. Again, however, water is exceptional in this regard: liquid water actually takes up less space than an equal mass of frozen water.

When a liquid experiences an increase in temperature, its particles take on energy and begin to move faster and faster. They collide with one another, and at some point the particles nearest the surface of the liquid acquire enough energy to break away from their neighbors. It is at this point that the liquid becomes a gas or vapor.

As heating continues, particles throughout the liquid begin to gain energy and move faster, but they do not immediately transform into gas. The reason is that the pressure of the liquid, combined with the pressure of the atmosphere above the liquid, tends to keep particles in place. Those particles below the surface, therefore, remain where they are until they acquire enough energy to rise to the surface.

The heated particle moves upward, leaving behind it a hollow space—a bubble. A bubble is not an empty space: it contains smaller trapped particles, but its small mass, relative to that of the liquid it disperses, makes it buoyant. Therefore, a bubble floats to the top, releasing its trapped particles as gas or vapor. At that point, the liquid is said to be boiling.

THE EFFECT OF ATMOSPHERIC PRESSURE.

The particles thus have to overcome atmospheric pressure as they rise, which means that the boiling point for any liquid depends in part on the pressure of the surrounding air. Normal atmospheric pressure (1 atm) is equal to 14 lb/in 2 (1.013 × 10 5 Pa), and is measured at sea level. The greater the altitude, the less the air pressure, because molecules of air—since air is a gas, and therefore its particles are fast-moving and non-attractive—respond less to Earth's gravitational pull. This is why airplanes require pressurized cabins to maintain an adequate oxygen supply; but even at altitudes much lower than the flight path of an airplane, differences in air pressure are noticeable.

It is for this reason that cooking instructions often vary with altitude. Atop Mt. Everest, Earth's highest peak at about 29,000 ft (8,839 m) above sea level, the pressure is approximately one-third of normal atmospheric pressure. Water boils at a much lower temperature on Everest than it does elsewhere: 158°F (70°C), as opposed to 212°F (100°C) at sea level. Of course, no one lives on the top of Mt. Everest—but people do live in Denver, Colorado, where the altitude is 5,577 ft (1,700 m) and the boiling point of water is 203°F (95°C).

Given the lower boiling point, one might assume that food would cook faster in Denver than in New York, Los Angeles, or in any city close to sea level. In fact, the opposite is true: because heated particles escape the water so much faster at high altitudes, they do not have time to acquire the energy needed to raise the temperature of the water. It is for this reason that a recipe may include a statement such as "at altitudes above XX feet, add XX minutes to cooking time."

If lowered atmospheric pressure means a lowered boiling point, what happens in outer space, where there is no atmospheric pressure? Liquids boil at very, very low temperatures. This is one of the reasons why astronauts have to wear pressurized suits: if they did not, their blood would boil—even though space itself is incredibly cold.

LIQUID TO GAS AND BACK AGAIN.

Note that the process of changing a liquid to a gas is similar to that which occurs when a solid changes to a liquid: particles gain heat and therefore energy, begin to move faster, break free from one another, and pass a certain threshold into a new phase of matter. And just as the freezing and melting point for a given substance are the same temperature—the only difference being one of orientation—the boiling point of a liquid transforming into a gas is the same as the condensation point for a gas turning into a liquid.

The behavior of water in boiling and condensation makes possible distillation, one of the principal methods for purifying sea water in various parts of the world. First the water is boiled, then it is allowed to cool and condense, thus forming water again. In the process, the water separates from the salt, leaving it behind in the form of brine. A similar separation takes place when salt water freezes: because salt, like most crystalline solids, has a much lower freezing point than water, very little of it remains joined to the water in ice. Instead, the salt takes the form of a briny slush.

A liquid that is vaporized, or any substance that exists normally as a gas, is quite different in physical terms from a solid or a liquid. This is illustrated by the much higher energy component in the molar heat of vaporization, or the amount of energy required to turn 1 mole of a liquid into a gas.

Consider, for instance, what happens to water when it experiences phase changes. Assuming that heat is added at a uniform rate, when ice reaches its melting point, there is only a relatively small period of time when the H 2 O is composed of both ice and liquid. But when the liquid reaches its boiling point, the water is present both as a liquid and a vapor for a much longer period of time. In fact, it takes almost seven times as much energy to turn liquid water into pure steam than it does to turn ice into purely liquid water. Thus, the molar heat of fusion for water is 6.02 kJ/mol, while the molar heat of vaporization is 40.6 kJ/mol.

Although liquid particles exert a moderate attraction toward one another, particles in a gas (particularly a substance that normally exists as a gas at ordinary temperatures on Earth) exert little to no attraction. They are thus free to move, and to move quickly. The overall shape and arrangement of gas is therefore random and indefinite—and, more importantly, the motion of gas particles provides much greater kinetic energy than is present in any other major form of matter on Earth.

The constant, fast, and random motion of gas particles means that they are regularly colliding and thereby transferring kinetic energy back and forth without any net loss of energy. These collisions also have the overall effect of producing uniform pressure in a gas. At the same time, the characteristics and behavior of gas particles indicate that they tend not to remain in an open container. Therefore, in order to have any pressure on a gas—other than normal atmospheric pressure—it is necessary to keep it in a closed container.

The Phase Diagram

The vaporization of water is an example of a change of phase—the transition from one phase of matter to another. The properties of any substance, and the points at which it changes phase, are plotted on what is known as a phase diagram. The phase diagram typically shows temperature along the x-axis, and pressure along the y-axis.

For simple substances, such as water and carbon dioxide (CO 2 ), the solid form of the substance appears at a relatively low temperature and at pressures anywhere from zero upward. The line between solids and liquids, indicating the temperature at which a solid becomes a liquid at any pressure above a certain level, is called the fusion curve. Though it appears to be a more or less vertical line, it is indeed curved, indicating that at high pressures, a solid well below the normal freezing point may be melted to create a liquid.

Liquids occupy the area of the phase diagram corresponding to relatively high temperatures and high pressures. Gases or vapors, on the other hand, can exist at very low temperatures, but only if the pressure is also low. Above the melting point for the substance, gases exist at higher pressures and higher temperatures. Thus, the line between liquids and gases often looks almost like a 45° angle. But it is not a straight line, as its name, the vaporization curve, implies. The curve of vaporization demonstrates that at relatively high temperatures and high pressures, a substance is more likely to be a gas than a liquid.

THE CRITICAL POINT.

There are several other interesting phenomena mapped on a phase diagram. One is the critical point, found at a place of very high temperature and pressure along the vaporization curve. At the critical point, high temperatures prevent a liquid from remaining a liquid, no matter how high the pressure.

At the same time, the pressure causes gas beyond that point to become increasingly more dense, but due to the high temperatures, it does not condense into a liquid. Beyond the critical point, the substance cannot exist in anything other than the gaseous state. The temperature component of the critical point for water is 705.2°F (374°C)—at 218 atm, or 218 times ordinary atmospheric pressure. For helium, however, critical temperature is just a few degrees above absolute zero. This is, in part, why helium is rarely seen in forms other than a gas.

THE SUBLIMATION CURVE.

Another interesting phenomenon is the sublimation curve, or the line between solid and gas. At certain very low temperatures and pressures, a substance may experience sublimation, meaning that a gas turns into a solid, or a solid into a gas, without passing through a liquid stage.

A well-known example of sublimation occurs when "dry ice," made of carbon dioxide, vaporizes at temperatures above (−78.5°C). Carbon dioxide is exceptional, however, in that it experiences sublimation at relatively high pressures that occur in everyday life: for most substances, the sublimation point transpires at such a low pressure point that it is seldom witnessed outside of a laboratory.

THE TRIPLE POINT.

The phenomenon known as the triple point shows how an ordinary substance such as water or carbon dioxide can actually be a liquid, solid, and vapor—all at once. Most people associate water as a gas or vapor (that is, steam) with very high temperatures. Yet, at a level far below normal atmospheric pressure, water can be a vapor at temperatures as low as −4°F (−20 °C). (All of the pressure values in the discussion of water at or near the triple point are far below atmospheric norms: the pressure at which water turns into a vapor at −4°F, for instance, is about 0.001 atm.)

Just as water can exist as a vapor at low temperatures and low pressures, it is also possible for water at temperatures below freezing to remain liquid. Under enough pressure, ice melts and is thereby transformed from a solid to a liquid, at temperatures below its normal freezing point. On the other hand, if the pressure of ice falls below a very low threshold, it will sublimate.

The phase diagram of water shows a line between the solid and liquid states that is almost, but not quite, exactly perpendicular to the x-axis. But in fact, it is a true fusion curve: it slopes slightly upward to the left, indicating that solid ice turns into water with an increase of pressure. Below a certain level of pressure is the vaporization curve, and where the fusion curve intersects the vaporization curve, there is a place called the triple point. Just below freezing, in conditions equivalent to about 0.007 atm, water is a solid, liquid, and vapor all at once.

Other States of Matter

Principal among states of matter other than solid, liquid, and gas is plasma, which is similar to gas. (The term "plasma," when referring to the state of matter, has nothing to do with the word as it is often used, in reference to blood plasma.) As with gas, plasma particles collide at high speeds—but in plasma the speeds are even greater, and the kinetic energy levels even higher.

The speed and energy of these collisions is directly related to the underlying property that distinguishes plasma from gas. So violent are the collisions between plasma particles that electrons are knocked away from their atoms. As a result, plasma does not have the atomic structure typical of a gas; rather, it is composed of positive ions and electrons. Plasma particles are thus electrically charged, and therefore greatly influenced by electric and magnetic fields.

Formed at very high temperatures, plasma is found in stars. The reaction between plasma and atomic particles in the upper atmosphere is responsible for the aurora borealis, or "northern lights." Though found on Earth only in very small quantities, plasma—ubiquitous in other parts of the universe—may be the most plentiful of all the states of matter.

QUASI-STATES.

Among the quasi-states of matter discussed by scientists are several terms describing the structure in which particles are joined, rather than the attraction and relative movement of those particles. Thus "crystalline," "amorphous," and "glassy" are all terms to describe what may be individual states of matter; so too is "colloidal."

A colloid is a structure intermediate in size between a molecule and a visible particle, and it has a tendency to be dispersed in another medium—the way smoke, for instance, is dispersed in air. Brownian motion describes the behavior of most colloidal particles. When one sees dust floating in a ray of sunshine through a window, the light reflects off colloids in the dust, which are driven back and forth by motion in the air otherwise imperceptible to the human senses.

DARK MATTER.

The number of states or phases of matter is clearly not fixed, and it is quite possible that more will be discovered in outer space, if not on Earth. One intriguing candidate is called dark matter, so described because it neither reflects nor emits light, and is therefore invisible. In fact, luminous or visible matter may very well make up only a small fraction of the mass in the universe, with the rest being taken up by dark matter.

If dark matter is invisible, how do astronomers and physicists know it exists? By analyzing the gravitational force exerted on visible objects in such cases where there appears to be no visible object to account for that force. An example is the center of our galaxy, the Milky Way. It appears to be nothing more than a dark "halo," but in order to cause the entire galaxy to revolve around it—in the same way that planets revolve around the Sun, though on a vastly larger scale—it must contain a staggering quantity of invisible mass.

The Bose-Einstein Condensate

Physicists at the Joint Institute of Laboratory Astrophysics in Boulder, Colorado, in 1995 revealed a highly interesting aspect of atomic behavior at temperatures approaching absolute zero. Some 70 years before, Einstein had predicted that, at extremely low temperatures, atoms would fuse to form one large "superatom." This hypothesized structure was dubbed the Bose-Einstein Condensate (BEC) after Einstein and Satyendranath Bose (1894-1974), an Indian physicist whose statistical methods contributed to the development of quantum theory.

Cooling about 2,000 atoms of the element rubidium to a temperature just 170 billionths of a degree Celsius above absolute zero, the physicists succeeded in creating an atom 100 micrometers across—still incredibly small, but vast in comparison to an ordinary atom. The superatom, which lasted for about 15 seconds, cooled down all the way to just 20 billionths of a degree above absolute zero. The Colorado physicists won the Nobel Prize in physics in 1997 for their work.

In 1999, researchers in a lab at Harvard University also created a superatom of BEC, and used it to slow light to just 38 MPH (61.2 km/h)—about 0.02% of its ordinary speed. Dubbed a "new" form of matter, the BEC may lead to a greater understanding of quantum mechanics, and may aid in the design of smaller, more powerful computer chips.

Some Unusual Phase Transitions

At places throughout this essay, references have been made variously to "phases" and "states" of matter. This is not intended to confuse, but rather to emphasize a particular point. Solids, liquids, and gases are referred to as "phases" because many (though far from all) substances on Earth regularly move from one phase to another.

There is absolutely nothing incorrect in referring to "states of matter." But "phases of matter" is used in the present context as a means of emphasizing the fact that substances, at the appropriate temperature and pressure, can be solid, liquid, or gas. The phases of matter, in fact, can be likened to the phases of a person's life: infancy, babyhood, childhood, adolescence, adulthood, old age. The transition between these stages is indefinite, yet it is easy enough to say when a person is at a certain stage.

LIQUID CRYSTALS.

A liquid crystal is a substance that, over a specific range of temperature, displays properties both of a liquid and a solid. Below this temperature range, it is unquestionably a solid, and above this range it is just as certainly a liquid. In between, however, liquid crystals exhibit a strange solid-liquid behavior: like a liquid, their particles flow, but like a solid, their molecules maintain specific crystalline arrangements.

The cholesteric class of liquid crystals is so named because the spiral patterns of light through the crystal are similar to those which appear in cholesterols. Depending on the physical properties of a cholesteric liquid crystal, only certain colors may be reflected. The response of liquid crystals to light makes them useful in liquid crystal displays (LCDs) found on laptop computer screens, camcorder views, and in other applications.

LIQUEFACTION OF GASES.

One interesting and useful application of phase change is the liquefaction of gases, or the change of gas into liquid by the reduction in its molecular energy levels. Liquefied natural gas (LNG) and liquefied petroleum gas (LPG), the latter a mixture of by-products obtained from petroleum and natural gas, are among the examples of liquefied gas in daily use. In both cases, the volume of the liquefied gas is far less than it would be if the gas were in a vaporized state, thus enabling ease and economy of transport.

Liquefied gases are used as heating fuel for motor homes, boats, and homes or cabins in remote areas. Other applications of liquefied gases include liquefied oxygen and hydrogen in rocket engines; liquefied oxygen and petroleum used in welding; and a combination of liquefied oxygen and nitrogen used in aqualung devices. The properties of liquefied gases figure heavily in the science of producing and studying low-temperature environments. In addition, liquefied helium is used in studying the behavior of matter at temperatures close to absolute zero.

COAL GASIFICATION.

Coal gasification, as one might discern from the name, is the conversion of coal to gas. Developed before World War II, it fell out of favor after the war, due to the lower cost of oil and natural gas. However, increasingly stringent environmental regulations imposed by the federal government on industry during the 1970s, combined with a growing concern for the environment on the part of the populace as a whole, led to a resurgence of interest in coal gasification.

Though widely used as a fuel in power plants, coal, when burned by ordinary means, generates enormous air pollution. Coal gasification, on the other hand, makes it possible to burn "clean" coal. Gasification involves a number of chemical reactions, some exothermic or heat-releasing, and some endothermic or heat-absorbing. At one point, carbon monoxide is released in an exothermic reaction, then mixed with hydrogen released from the coal to create a second exothermic reaction. The energy discharged in these first two reactions is used to initiate a third, endothermic, reaction.

The finished product of coal gasification is a mixture containing carbon monoxide, methane, hydrogen, and other substances, and this—rather than ordinary coal—is burned as a fuel. The composition of the gases varies according to the process used. Products range from coal synthesis gas and medium-Btu gas (both composed of carbon monoxide and hydrogen, though combined in different forms) to substitute natural gas, which consists primarily of methane.

Not only does coal gasification produce a clean-burning product, but it does so without the high costs associated with flue-gas desulfurization systems. The latter, often called "scrubbers," were originally recommended by the federal government to industry, but companies discovered that coal gasification could produce the same results for much less money. In addition, the waste products from coal gasification can be used for other purposes. At the Cool Water Integrated Gasification Combined Cycle Plant, established in Barstow, California, in 1984, sulfur obtained from the reduction of sulfur dioxide is sold off for about $100 a ton.

The Chemical Dimension to Changes of Phase

Throughout much of this essay, we have discussed changes of phase primarily in physical terms; yet clearly these changes play a significant role in chemistry. Furthermore, coal gasification serves to illustrate the impact chemical processes can have on changes of state.

Much earlier, figures were given for the melting points of copper, water, and helium, and these were compared with the average atomic mass of each. Those figures, again, are:

Average Atomic Mass and Melting Points of a Sample Gas, Liquid, and Solid

• Helium: 4.003 amu; −457.6°F (−210°C)
• Water: 40.304 amu 32°F (0°C)
• Copper: 63.55 amu; 1,985°F (1,085°C).

Something seems a bit strange about those comparisons: specifically, the differences in melting point appear to be much more dramatic than the differences in average atomic mass. Clearly, another factor is at work—a factor that relates to the difference in the attractions between molecules in each. Although the differences between solids, liquids, and gases are generally physical, the one described here—a difference between substances—is clearly chemical in nature.

To discuss this in the detail it deserves would require a lengthy digression on the chemical dimensions of intermolecular attraction. Nonetheless, it is possible here to offer at least a cursory answer to the question raised by these striking differences in response to temperature.

Dipoles, Electron Seas, and London Dispersion

Water molecules are polar, meaning that one area of a water molecule is positively charged, while another area has a negative charge. Thus the positive side of one molecule is drawn to the negative side of another, and vice versa, which gives water a much stronger intermolecular bond than, for instance, oil, in which the positive and negative charges are evenly distributed throughout the molecule.

Yet the intermolecular attraction between the dipoles (as they are called) in water is not nearly as strong as the bond that holds together a metal. Particles in copper or other metals "float" in a tightly packed "sea" of highly mobile electrons, which provide a bond that is powerful, yet lacking in a firm directional orientation. Thus metals are both strong and highly malleable (that is, they can be hammered very flat without breaking.)

Water, of course, appears most often as a liquid, and copper as a solid, precisely because water has a very high boiling point (the point at which it becomes a vapor) and copper has a very high melting point. But consider helium, which has the lowest freezing point of any element: just above absolute zero. Even then, a pressure equal to 25 times that of normal atmospheric pressure is required to push it past the freezing point.

Helium and other Group 8 or Group 18 elements, as well as non-polar molecules such as oils, are bonded by what is called London dispersion forces. The latter, as its name suggests, tends to keep molecules dispersed, and induces instantaneous dipoles when most of the electrons happen to be on one side of an atom. Of course, this happens only for an infinitesimal fraction of time, but it serves to create a weak attraction. Only at very low temperatures do London dispersion forces become strong enough to result in the formation of a solid.

WHERE TO LEARN MORE

Biel, Timothy L. Atom: Building Blocks of Matter. San Diego, CA: Lucent Books, 1990.

Feynman, Richard. Six Easy Pieces: Essentials of Physics Explained by Its Most Brilliant Teacher. New introduction by Paul Davies. Cambridge, MA: Perseus Books, 1995.

"High School Chemistry Table of Contents—Solids and Liquids" Homeworkhelp.com (Web site). <http://www.homeworkhelp.com/homeworkhelp/freemember/text/chem/high/topic09.htm> (April 10, 2001).

"Matter: Solids, Liquids, Gases." Studyweb (Web site). <http://www.studyweb.com/links/4880.html> (April 10, 2001).

"The Molecular Circus" (Web site). <http://www.cpo.com/Weblabs/circus.htm> (April 10, 2001).

Paul, Richard. A Handbook to the Universe: Explorations of Matter, Energy, Space, and Time for Beginning Scientific Thinkers. Chicago: Chicago Review Press, 1993.

"Phases of Matter" (Web site). <http://pc65.frontier.osrhe.edu/hs/science/pphase.htm> (April 10, 2001).

Royston, Angela. Solids, Liquids, and Gasses. Chicago: Heinemann Library, 2001.

Wheeler, Jill C. The Stuff Life's Made Of: A Book About Matter. Minneapolis, MN: Abdo & Daughters Publishing, 1996.

Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.

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7.4: Phase Changes

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Learning Objectives

• Determine the heat associated with a phase change.

Matter can exist in one of several different states, including a gas, liquid, or solid state. The amount of energy in molecules of matter determines the state of matter .

• A gas is a state of matter in which atoms or molecules have enough energy to move freely. The molecules come into contact with one another only when they randomly collide.
• A liquid is a state of matter in which atoms or molecules are constantly in contact but have enough energy to keep changing positions relative to one another.
• A solid is a state of matter in which atoms or molecules do not have enough energy to move. They are constantly in contact and in fixed positions relative to one another.

The following are the changes of state:

• If heat is added to a substance, such as in melting, vaporization, and sublimation, the process is endothermic . In this instance, heat is increasing the speed of the molecules causing them move faster (examples: solid to liquid; liquid to gas; solid to gas).
• If heat is removed from a substance, such as in freezing and condensation, then the process is exothermic . In this instance, heat is decreasing the speed of the molecules causing them move slower (examples: liquid to solid; gas to liquid). These changes release heat to the surroundings.
• The amount of heat needed to change a sample from solid to liquid would be the same to reverse from liquid to solid. The only difference is the direction of heat transfer.

Example \(\PageIndex{1}\)

Label each of the following processes as endothermic or exothermic.

• water boiling
• ice forming on a pond
• endothermic - you must put a pan of water on the stove and give it heat in order to get water to boil. Because you are adding heat/energy, the reaction is endothermic.
• exothermic - think of ice forming in your freezer instead. You put water into the freezer, which takes heat out of the water, to get it to freeze. Because heat is being pulled out of the water, it is exothermic. Heat is leaving.

Exercise \(\PageIndex{1}\)

• water vapor condensing
• gold melting

a. exothermic

b. endothermic

A phase change is a physical process in which a substance goes from one phase to another. Usually the change occurs when adding or removing heat at a particular temperature, known as the melting point or the boiling point of the substance. The melting point is the temperature at which the substance goes from a solid to a liquid (or from a liquid to a solid). The boiling point is the temperature at which a substance goes from a liquid to a gas (or from a gas to a liquid). The nature of the phase change depends on the direction of the heat transfer. Heat going into a substance changes it from a solid to a liquid or a liquid to a gas. Removing heat from a substance changes a gas to a liquid or a liquid to a solid.

Two key points are worth emphasizing. First, at a substance’s melting point or boiling point, two phases can exist simultaneously. Take water (H 2 O) as an example. On the Celsius scale, H 2 O has a melting point of 0°C and a boiling point of 100°C. At 0°C, both the solid and liquid phases of H 2 O can coexist. However, if heat is added, some of the solid H 2 O will melt and turn into liquid H 2 O. If heat is removed, the opposite happens: some of the liquid H 2 O turns into solid H 2 O. A similar process can occur at 100°C: adding heat increases the amount of gaseous H 2 O, while removing heat increases the amount of liquid H 2 O (Figure \(\PageIndex{1}\)).

Water is a good substance to use as an example because many people are already familiar with it. Other substances have melting points and boiling points as well.

Second, as shown in Figure \(\PageIndex{1}\), the temperature of a substance does not change as the substance goes from one phase to another . In other words, phase changes are isothermal (isothermal means “constant temperature”). Again, consider H 2 O as an example. Solid water (ice) can exist at 0°C. If heat is added to ice at 0°C, some of the solid changes phase to make liquid, which is also at 0°C. Remember, the solid and liquid phases of H 2 O can coexist at 0°C. Only after all of the solid has melted into liquid does the addition of heat change the temperature of the substance.

For each phase change of a substance, there is a characteristic quantity of heat needed to perform the phase change per gram (or per mole) of material. The heat of fusion (Δ H fus ) is the amount of heat per gram (or per mole) required for a phase change that occurs at the melting point. The heat of vaporization (Δ H vap ) is the amount of heat per gram (or per mole) required for a phase change that occurs at the boiling point. If you know the total number of grams or moles of material, you can use the Δ H fus or the Δ H vap to determine the total heat being transferred for melting or solidification using these expressions:

\[\text{heat} = n \times ΔH_{fus} \label{Eq1a} \]

wher e \(n\) is th e number of moles and \(ΔH_{fus}\) is expressed in energy/mole or

\[\text{heat} = m \times ΔH_{fus} \label{Eq1b} \]

where \(m\) is the mass in grams and \(ΔH_{fus}\) is expressed in energy/gram.

For the boiling or condensation, use these expressions:

\[\text{heat} = n \times ΔH_{vap} \label{Eq2a} \]

wher e \(n\) is the number of moles) and \(ΔH_{vap}\) is expressed in energy/mole or

\[\text{heat} = m \times ΔH_{vap} \label{Eq2b} \]

wh ere \(m\) i s the mass in grams and \(ΔH_{vap}\) is expressed in energy/gram.

Remember that a phase change depends on the direction of the heat transfer. If heat transfers in, solids become liquids, and liquids become solids at the melting and boiling points, respectively. If heat transfers out, liquids solidify, and gases condense into liquids. At these points, there are no changes in temperature as reflected in the above equations.

Example \(\PageIndex{2}\)

How much heat is necessary to melt 55.8 g of ice (solid H 2 O) at 0°C? The heat of fusion of H 2 O is 79.9 cal/g.

We can use the relationship between heat and the heat of fusion (Equation \(\PageIndex{1}\)) to determine how many cal of heat are needed to melt this ice:

\[ \begin{align*} \ce{heat} &= \ce{m \times ΔH_{fus}} \\[4pt] \mathrm{heat} &= \mathrm{(55.8\: \cancel{g})\left(\dfrac{79.9\: cal}{\cancel{g}}\right)=4,460\: cal} \end{align*} \nonumber \]

Exercise \(\PageIndex{2}\)

How much heat is necessary to vaporize 685 g of H 2 O at 100°C? The heat of vaporization of H 2 O is 540 cal/g.

\[ \begin{align*} \ce{heat} &= \ce{m \times ΔH_{vap}} \\[4pt] \mathrm{heat} &= \mathrm{(685\: \cancel{g})\left(\dfrac{540\: cal}{\cancel{g}}\right)=370,000\: cal} \end{align*} \nonumber \]

Table \(\PageIndex{1}\) lists the heats of fusion and vaporization for some common substances. Note the units on these quantities; when you use these values in problem solving, make sure that the other variables in your calculation are expressed in units consistent with the units in the specific heats or the heats of fusion and vaporization.

Sublimation

There is also a phase change where a solid goes directly to a gas:

\[\text{solid} \rightarrow \text{gas} \label{Eq3} \]

This phase change is called sublimation . Each substance has a characteristic heat of sublimation associated with this process. For example, the heat of sublimation (Δ H sub ) of H 2 O is 620 cal/g.

We encounter sublimation in several ways. You may already be familiar with dry ice, which is simply solid carbon dioxide (CO 2 ). At −78.5°C (−109°F), solid carbon dioxide sublimes, changing directly from the solid phase to the gas phase:

\[\mathrm{CO_2(s) \xrightarrow{-78.5^\circ C} CO_2(g)} \label{Eq4} \]

Solid carbon dioxide is called dry ice because it does not pass through the liquid phase. Instead, it does directly to the gas phase. (Carbon dioxide can exist as liquid but only under high pressure.) Dry ice has many practical uses, including the long-term preservation of medical samples.

Even at temperatures below 0°C, solid H 2 O will slowly sublime. For example, a thin layer of snow or frost on the ground may slowly disappear as the solid H 2 O sublimes, even though the outside temperature may be below the freezing point of water. Similarly, ice cubes in a freezer may get smaller over time. Although frozen, the solid water slowly sublimes, redepositing on the colder cooling elements of the freezer, which necessitates periodic defrosting (frost-free freezers minimize this redeposition). Lowering the temperature in a freezer will reduce the need to defrost as often.

Under similar circumstances, water will also sublime from frozen foods (e.g., meats or vegetables), giving them an unattractive, mottled appearance called freezer burn. It is not really a “burn,” and the food has not necessarily gone bad, although it looks unappetizing. Freezer burn can be minimized by lowering a freezer’s temperature and by wrapping foods tightly so water does not have any space to sublime into.

Key Takeaway

• There is an energy change associated with any phase change.

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Condensed Matter > Superconductivity

Title: the phase diagram of high-tc cuprates.

Abstract: We investigate the detailed structure of the T vs. doping phase diagram of hole doped High-Tc superconducting cuprates, from the point of view of a recently proposed comprehensive theory for these materials. We obtain, in particular, analytic expressions for the transition lines delimiting the Neel phase, TN (x), the Spin-Glass phase, Tg(x) and the CDW Charge Ordered phase, Tco(x). These, results along with the previously derived expressions for the transition lines delimiting the Superconducting phase, Tc (x), the Pseudogap (and Strange Metal) phases, T*(x) and the Fermi Liquid phase, Tfl(x) are in excellent agreement with the experimental data and form an unprecedentedly accurate description of cuprates. Our study reveals, in particular, the complementary role, which is played by the mechanisms responsible for the onset of superconducting (SC) and Spin-Glass (SG) phases in these materials. The absence of SG phases in electron doped High-Tc superconducting cuprates, strongly suggests a different SC mechanism should operate in those materials.

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An Essay on Juvenility, Phase Change, and Heteroblasty in Seed Plants

• PMID: 10572025
• DOI: 10.1086/314215

Phase change (the change from nonreproductive to reproductive status) and heteroblasty (ontogenetic changes in vegetative metamers) are two determinants of longitudinal asymmetry in plants. These concepts are critically important to understanding the regulation of plant development as well as morphological evolution and life-history variation. Since Goebel, the two have been conflated. This article questions how phase change and heteroblasty are delimited and explores some of the problems that arise in the explicit or implicit link between them, given that several lines of evidence indicate that they are distinct and independent facets of plant development. It is suggested that problems are perpetuated through use of the terms "juvenile" and "adult" to describe both phenomena.