Heat is a form of energy—specifically, the energy that flows between two bodies because of differences in temperature. Therefore, the scientific definition of heat is different from, and more precise than, the everyday meaning. Physicists working in the area of thermodynamics study heat from a number of perspectives, including specific heat, or the amount of energy required to change the temperature of a substance, and calorimetry, the measurement of changes in heat as a result of physical or chemical changes. Thermodynamics helps us to understand such phenomena as the operation of engines and the gradual breakdown of complexity in physical systems—a phenomenon known as entropy.
HOW IT WORKS
Heat, Work, and Energy
Thermodynamics is the study of the relationships between heat, work, and energy. Work is the exertion of force over a given distance to displace or move an object, and is, thus, the product of force and distance exerted in the same direction. Energy, the ability to accomplish work, appears in numerous manifestations—including thermal energy, or the energy associated with heat.
Thermal and other types of energy, including electromagnetic, sound, chemical, and nuclear energy, can be described in terms of two extremes: kinetic energy, or the energy associated with movement, and potential energy, or the energy associated with position. If a spring is pulled back to its maximum point of tension, its potential energy is also at a maximum; once it is released and begins springing through the air to return to its original position, it begins gaining kinetic energy and losing potential energy.
All manifestations of energy appear in both kinetic and potential forms, somewhat like the way football teams are organized to play both offense or defense. Just as a football team takes an offensive role when it has the ball, and a defensive role when the other team has it, a physical system typically undergoes regular transformations between kinetic and potential energy, and may have more of one or the other, depending on what is taking place in the system.
What Heat Is and Is Not
Thermal energy is actually a form of kinetic energy generated by the movement of particles at the atomic or molecular level: the greater the movement of these particles, the greater the thermal energy. Heat is internal thermal energy that flows from one body of matter to another—or, more specifically, from a system at a higher temperature to one at a lower temperature. Thus, temperature, like heat, requires a scientific definition quite different from its common meaning: temperature measures the average molecular kinetic energy of a system, and governs the direction of internal energy flow between them.
Two systems at the same temperature are said to be in a state of thermal equilibrium. When this occurs, there is no exchange of heat. Though in common usage, "heat" is an expression of relative warmth or coldness, in physical terms, heat exists only in transfer between two systems. What people really mean by "heat" is the internal energy of a system—energy that is a property of that system rather than a property of transferred internal energy.
NO SUCH THING AS "COLD."
Though the term "cold" has plenty of meaning in the everyday world, in physics terminology, it does not. Cold and heat are analogous to darkness and light: again, darkness means something in our daily experience, but in physical terms, darkness is simply the absence of light. To speak of cold or darkness as entities unto themselves is rather like saying, after spending 20 dollars, "I have 20 non-dollars in my pocket."
If you grasp a snowball in your hand, of course, your hand gets cold. The human mind perceives this as a transfer of cold from the snowball, but, in fact, exactly the opposite happens: heat moves from your hand to the snow, and if enough heat enters the snowball, it will melt. At the same time, the departure of heat from your hand results in a loss of internal energy near the surface of your hand, which you experience as a sensation of coldness.
Transfers of Heat
In holding the snowball, heat passes from the surface of the hand by one means, conduction, then passes through the snowball by another means, convection. In fact, there are three methods heat is transferred: conduction, involving successive molecular collisions and the transfer of heat between two bodies in contact; convection, which requires the motion of fluid from one place to another; or radiation, which takes place through electromagnetic waves and requires no physical medium, such as water or air, for the transfer.
Solids, particularly metals, whose molecules are packed relatively close together, are the best materials for conduction. Molecules of liquid or nonmetallic solids vary in their ability to conduct heat, but gas is a poor conductor, because of the loose attractions between its molecules.
The qualities that make metallic solids good conductors of heat, as a matter of fact, also make them good conductors of electricity. In the conduction of heat, kinetic energy is passed from molecule to molecule, like a long line of people standing shoulder to shoulder, passing a secret. (And, just as the original phrasing of the secret becomes garbled, some kinetic energy is inevitably lost in the series of transfers.)
As for electrical conduction, which takes place in a field of electric potential, electrons are freed from their atoms; as a result, they are able to move along the line of molecules. Because plastic is much less conductive than metal, an electrician uses a screwdriver with a plastic handle; similarly, a metal cooking pan typically has a wooden or plastic handle.
Wherever fluids are involved—and in physics, "fluid" refers both to liquids and gases—convection is a common form of heat transfer. Convection involves the movement of heated material—whether it is air, water, or some other fluid.
Convection is of two types: natural convection and forced convection, in which a pump or other mechanism moves the heated fluid. When heated air rises, this is an example of natural convection. Hot air has a lower density than that of the cooler air in the atmosphere above it, and, therefore, is buoyant; as it rises, however, it loses energy and cools. This cooled air, now denser than the air around it, sinks again, creating a repeating cycle that generates wind.
Examples of forced convection include some types of ovens and even a refrigerator or air conditioner. These two machines both move warm air from an interior to an exterior place. Thus, the refrigerator pulls hot air from the compartment and expels it to the surrounding room, while an air conditioner pulls heat from a building and releases it to the outside.
But forced convection does not necessarily involve humanmade machines: the human heart is a pump, and blood carries excess heat generated by the body to the skin. The heat passes through the skin by means of conduction, and at the surface of the skin, it is removed from the body in a number of ways, primarily by the cooling evaporation of perspiration.
Outer space, of course, is cold, yet the Sun's rays warm the Earth, an apparent paradox. Because there is no atmosphere in space, convection is impossible. In fact, heat from the Sun is not dependant on any fluid medium for its transfer: it comes to Earth by means of radiation. This is a form of heat transfer significantly different from the other two, because it involves electromagnetic energy, instead of ordinary thermal energy generated by the action of molecules. Heat from the Sun comes through a relatively narrow area of the light spectrum, including infrared, visible light, and ultraviolet rays.
Every form of matter emits electromagnetic waves, though their presence may not be readily perceived. Thus, when a metal rod is heated, it experiences conduction, but part of its heat is radiated, manifested by its glow—visible light. Even when the heat in an object is not visible, however, it may be radiating electromagnetic energy, for instance, in the form of infrared light. And, of course, different types of matter radiate better than others: in general, the better an object is at receiving radiation, the better it is at emitting it.
The measurement of temperature by degrees in the Fahrenheit or Celsius scales is a part of everyday life, but measurements of heat are not as familiar to the average person. Because heat is a form of energy, and energy is the ability to perform work, heat is, therefore, measured by the same units as work.
The principal unit of work or energy in the metric system (known within the scientific community as SI, or the SI system) is the joule. Abbreviated "J," a joule is equal to 1 newton-meter (N · m). The newton is the SI unit of force, and since work is equal to force multiplied by distance, measures of work can also be separated into these components. For instance, the British measure of work brings together a unit of distance, the foot, and a unit of force, the pound. A foot-pound (ft · lb) is equal to 1.356 J, and 1 joule is equal to 0.7376 ft · lb.
In the British system, Btu, or British thermal unit, is another measure of energy used for machines such as air conditioners. One Btu is equal to 778 ft · lb or 1,054 J. The kilocalorie in addition to the joule, is an important SI measure of heat. The amount of energy required to change the temperature of 1 gram of water by 1°C is called a calorie, and a kilocalorie is equal to 1,000 calories. Somewhat confusing is the fact that the dietary Calorie (capital C), with which most people are familiar, is not the same as a calorie (lowercase C)—rather, a dietary Calorie is the equivalent of a kilocalorie.
Specific heat is the amount of heat that must be added to, or removed from, a unit of mass for a given substance to change its temperature by 1°C. Thus, a kilocalorie, because it measures the amount of heat necessary to effect that change precisely for a kilogram of water, is identical to the specific heat for that particular substance in that particular unit of mass.
The higher the specific heat, the more resistant the substance is to changes in temperature. Many metals, in fact, have a low specific heat, making them easy to heat up and cool down. This contributes to the tendency of metals to expand when heated (a phenomenon also discussed in the Thermal Expansion essay), and, thus, to their malleability.
MEASURING AND CALCULATING SPECIFIC HEAT.
The specific heat of any object is a function of its mass, its composition, and the desired change in temperature. The values of the initial and final temperature are not important—only the difference between them, which is the temperature change.
The components of specific heat are related to one another in the formula Q = mc δT. Here Q is the quantity of heat, measured in joules, which must be added. The mass of the object is designated by m, and the specific heat of the particular substance in question is represented with c. The Greek letter delta (δ) designates change, and δT stands for "change in temperature."
Specific heat is measured in units of J/kg · °C (joules per kilogram-degree Centigrade), though for the sake of convenience, this is usually rendered in terms of kilojoules (kJ), or 1,000 joules—that is, kJ/kg · °C. The specific heat of water is easily derived from the value of a kilo-calorie: it is 4.185, the same number of joules required to equal a kilocalorie.
The measurement of heat gain or loss as a result of physical or chemical change is called calorimetry (pronounced kal-IM-uh-tree). Like the word "calorie," the term is derived from a Latin root meaning "heat."
The foundations of calorimetry go back to the mid-nineteenth century, but the field owes much to scientists' work that took place over a period of about 75 years prior to that time. In 1780, French chemist Antoine Lavoisier (1743-1794) and French astronomer and mathematician Pierre Simon Laplace (1749-1827) had used a rudimentary ice calorimeter for measuring the heats in formations of compounds. Around the same time, Scottish chemist Joseph Black (1728-1799) became the first scientist to make a clear distinction between heat and temperature.
By the mid-1800s, a number of thinkers had come to the realization that—contrary to prevailing theories of the day—heat was a form of energy, not a type of material substance. Among these were American-British physicist Benjamin Thompson, Count Rumford (1753-1814) and English chemist James Joule (1818-1889)—for whom, of course, the joule is named.
Calorimetry as a scientific field of study actually had its beginnings with the work of French chemist Pierre-Eugene Marcelin Berthelot (1827-1907). During the mid-1860s, Berthelot became intrigued with the idea of measuring heat, and by 1880, he had constructed the first real calorimeter.
Essential to calorimetry is the calorimeter, which can be any device for accurately measuring the temperature of a substance before and after a change occurs. A calorimeter can be as simple as a styrofoam cup. Its quality as an insulator, which makes styrofoam ideal for holding in the warmth of coffee and protecting the hand from scalding as well, also makes styrofoam an excellent material for calorimetric testing. With a styrofoam calorimeter, the temperature of the substance inside the cup is measured, a reaction is allowed to take place, and afterward, the temperature is measured a second time.
The most common type of calorimeter used is the bomb calorimeter, designed to measure the heat of combustion. Typically, a bomb calorimeter consists of a large container filled with water, into which is placed a smaller container, the combustion crucible. The crucible is made of metal, having thick walls with an opening through which oxygen can be introduced. In addition, the combustion crucible is designed to be connected to a source of electricity.
In conducting a calorimetric test using a bomb calorimeter, the substance or object to be studied is placed inside the combustion crucible and ignited. The resulting reaction usually occurs so quickly that it resembles the explosion of a bomb—hence, the name "bomb calorimeter." Once the "bomb" goes off, the resulting transfer of heat creates a temperature change in the water, which can be readily gauged with a thermometer.
To study heat changes at temperatures higher than the boiling point of water (212°F or 100°C), physicists use substances with higher boiling points. For experiments involving extremely large temperature ranges, an aneroid (without liquid) calorimeter may be used. In this case, the lining of the combustion crucible must be of a metal, such as copper, with a high coefficient or factor of thermal conductivity.
The bomb calorimeter that Berthelot designed in 1880 measured the caloric value of fuels, and was applied to determining the thermal efficiency of a heat engine. A heat engine is a machine that absorbs heat at a high temperature, performs mechanical work, and as a result, gives off heat at a lower temperature.
The desire to create efficient heat engines spurred scientists to a greater understanding of thermodynamics, and this resulted in the laws of thermodynamics, discussed at the conclusion of this essay. Their efforts were intimately connected with one of the greatest heat engines ever created, a machine that literally powered the industrialized world during the nineteenth century: the steam engine.
HOW A STEAM ENGINE WORKS.
Like all heat engines (except reverse heat engines such as the refrigerator, discussed below), a steam engine pulls heat from a high-temperature reservoir to a low-temperature reservoir, and in the process, work is accomplished. The hot steam from the high-temperature reservoir makes possible the accomplishment of work, and when the energy is extracted from the steam, the steam condenses in the low-temperature reservoir, becoming relatively cool water.
A steam engine is an external-combustion engine, as opposed to the internal-combustion engine that took its place at the forefront of industrial technology at the beginning of the twentieth century. Unlike an internal-combustion engine, a steam engine burns its fuel outside the engine. That fuel may be simply firewood, which is used to heat water and create steam. The thermal energy of the steam is then used to power a piston moving inside a cylinder, thus, converting thermal energy to mechanical energy for purposes such as moving a train.
EVOLUTION OF STEAM POWER.
As with a number of advanced concepts in science and technology, the historical roots of the steam engine can be traced to the Greeks, who—just as they did with ideas such as the atom or the Sun-centered model of the universe—thought about it, but failed to develop it. The great inventor Hero of Alexandria (c. 65-125) actually created several steam-powered devices, but he perceived these as mere novelties, hardly worthy of scientific attention. Though Europeans adopted water power, as, for instance, in waterwheels, during the late ancient and medieval periods, further progress in steam power did not occur for some 1,500 years.
Following the work of French physicist Denis Papin (1647-1712), who invented the pressure cooker and conducted the first experiments with the use of steam to move a piston, English engineer Thomas Savery (c. 1650-1715) built the first steam engine. Savery had abandoned the use of the piston in his machine, but another English engineer, Thomas Newcomen (1663-1729), reintroduced the piston for his own steam-engine design.
Then in 1763, a young Scottish engineer named James Watt (1736-1819) was repairing a Newcomen engine and became convinced he could build a more efficient model. His steam engine, introduced in 1769, kept the heating and cooling processes separate, eliminating the need for the engine to pause in order to reheat. These and other innovations that followed—including the introduction of a high-pressure steam engine by English inventor Richard Trevithick (1771-1833)—transformed the world.
CARNOT PROVIDES THEORETICAL UNDERSTANDING.
The men who developed the steam engine were mostly practical-minded figures who wanted only to build a better machine; they were not particularly concerned with the theoretical explanation for its workings. Then in 1824, a French physicist and engineer by the name of Sadi Carnot (1796-1832) published his sole work, the highly influential Reflections on the Motive Power of Fire (1824), in which he discussed heat engines scientifically.
In Reflections, Carnot offered the first definition of work in terms of physics, describing it as "weight lifted through a height." Analyzing Watt's steam engine, he also conducted groundbreaking studies in the nascent science of thermodynamics. Every heat engine, he explained, has a theoretical limit of efficiency related to the temperature difference in the engine: the greater the difference between the lowest and highest temperature, the more efficient the engine.
Carnot's work influenced the development of more efficient steam engines, and also had an impact on the studies of other physicists investigating the relationship between work, heat, and energy. Among these was William Thomson, Lord Kelvin (1824-1907). In addition to coining the term "thermodynamics," Kelvin developed the Kelvin scale of absolute temperature and established the value of absolute zero, equal to −273.15°C or −459.67°F.
According to Carnot's theory, maximum effectiveness was achieved by a machine that could reach absolute zero. However, later developments in the understanding of thermodynamics, as discussed below, proved that both maximum efficiency and absolute zero are impossible to attain.
REVERSE HEAT ENGINES.
It is easy to understand that a steam engine is a heat engine: after all, it produces heat. But how is it that a refrigerator, an air conditioner, and other cooling machines are also heat engines? Moreover, given the fact that cold is the absence of heat and heat is energy, one might ask how a refrigerator or air conditioner can possibly use energy to produce cold, which is the same as the absence of energy. In fact, cooling machines simply reverse the usual process by which heat engines operate, and for this reason, they are called "reverse heat engines." Furthermore, they use energy to extract heat.
A steam engine takes heat from a high-temperature reservoir—the place where the water is turned into steam—and uses that energy to produce work. In the process, energy is lost and the heat moves to a low-temperature reservoir, where it condenses to form relatively cool water. A refrigerator, on the other hand, pulls heat from a low-temperature reservoir called the evaporator, into which flows heat from the refrigerated compartment—the place where food and other perishables are kept. The coolant from the evaporator take this heat to the condenser, a high-temperature reservoir at the back of the refrigerator, and in the process it becomes a gas. Heat is released into the surrounding air; this is why the back of a refrigerator is hot.
Instead of producing a work output, as a steam engine does, a refrigerator requires a work input—the energy supplied via the wall outlet. The principles of thermodynamics show that heat always flows from a high-temperature to a low-temperature reservoir, and reverse heat engines do not defy these laws. Rather, they require an external power source in order to effect the transfer of heat from a low-temperature reservoir, through the gases in the evaporator, to a high-temperature reservoir.
The Laws of Thermodynamics
THE FIRST LAW OF THERMODYNAMICS.
There are three laws of thermodynamics, which provide parameters as to the operation of thermal systems in general, and heat engines in particular. The history behind the derivation of these laws is discussed in the essay on Thermodynamics; here, the laws themselves will be examined in brief form.
The physical law known as conservation of energy shows that within a system isolated from all outside factors, the total amount of energy remains the same, though transformations of energy from one form to another take place. The first law of thermodynamics states the same fact in a somewhat different manner.
According to the first law of thermodynamics, because the amount of energy in a system remains constant, it is impossible to perform work that results in an energy output greater than the energy input. Thus, it could be said that the conservation of energy law shows that "the glass is half full": energy is never lost. On the hand, the first law of thermodynamics shows that "the glass is half empty": no machine can ever produce more energy than was put into it. Hence, a perpetual motion machine is impossible, because in order to keep a machine running continually, there must be a continual input of energy.
THE SECOND LAW OF THERMODYNAMICS.
The second law of thermodynamics begins from the fact that the natural flow of heat is always from a high-temperature to a low-temperature reservoir. As a result, no engine can be constructed that simply takes heat from a source and performs an equivalent amount of work: some of the heat will always be lost. In other words, it is impossible to build a perfectly efficient engine.
In effect, the second law of thermodynamics compounds the "bad news" delivered by the first law with some even worse news: though it is true that energy is never lost, the energy available for work output will never be as great as the energy put into a system. Linked to the second law is the concept of entropy, the tendency of natural systems toward breakdown, and specifically, the tendency for the energy in a system to be dissipated. "Dissipated" in this context means that the high-and low-temperature reservoirs approach equal temperatures, and as this occurs, entropy increases.
THE THIRD LAW OF THERMODYNAMICS.
Entropy also plays a part in the third law of thermodynamics, which states that at the temperature of absolute zero, entropy also approaches zero. This might seem to counteract the "worse news" of the second law, but in fact, what the third law shows is that absolute zero is impossible to reach.
As stated earlier, Carnot's engine would achieve perfect efficiency if its lowest temperature were the same as absolute zero; but the second law of thermodynamics shows that a perfectly efficient machine is impossible. Relativity theory (which first appeared in 1905, the same year as the third law of thermodynamics) showed that matter can never exceed the speed of light. In the same way, the collective effect of the second and third laws is to prove that absolute zero—the temperature at which molecular motion in all forms of matter theoretically ceases—can never be reached.
WHERE TO LEARN MORE
Beiser, Arthur. Physics, 5th ed. Reading, MA: Addison-Wesley, 1991.
Bonnet, Robert L and Dan Keen. Science Fair Projects: Physics. Illustrated by Frances Zweifel. New York: Sterling, 1999.
Encyclopedia of Thermodynamics (Web site). <http://therion.minpet.unibas.ch/minpet/groups/thermodict/> (April 12, 2001).
Friedhoffer, Robert. Physics Lab in the Home. Illustrated by Joe Hosking. New York: Franklin Watts, 1997.
Manning, Mick and Brita Granström. Science School. New York: Kingfisher, 1998.
Macaulay, David. The New Way Things Work. Boston: Houghton Mifflin, 1998.
Moran, Jeffrey B. How Do We Know the Laws of Thermodynamics? New York: Rosen Publishing Group, 2001.
Santrey, Laurence. Heat. Illustrated by Lloyd Birmingham. Mahwah, NJ: Troll Associates, 1985.
Suplee, Curt. Everyday Science Explained. Washington, D.C.: National Geographic Society, 1996.
"Temperature and Thermodynamics" PhysLINK.com (Web site). <http://www.physlink.com/ae_thermo.cfm> (April 12, 2001).
The temperature, defined as 0K on the Kelvin scale, at which the motion of molecules in a solid virtually ceases. The third law of thermodynamics establishes the impossibility of actually reaching absolute zero.
BTU (BRITISH THERMAL UNIT):
A measure of energy or heat in the Britishsystem, often used in reference to the capacity of an air conditioner. A Btu is equal to 778 foot-pounds, or 1,054 joules.
A measure of heat or energy in the SI or metric system, equal to the heat that must be added to or removed from 1 gram of water to change its temperature by1°C. The dietary Calorie (capital C) with which most people are familiar is the same as the kilocalorie.
The measurement of heat gain or loss as a result of physical or chemical change.
The transfer of heat by successive molecular collisions. Conduction is the principal means of heat transfer in solids, particularly metals.
CONSERVATION OF ENERGY:
A law of physics stating that within a system isolated from all other outside factors, the total amount of energy remains the same, though transformations of energy from one form to another take place. The first law of thermodynamics is the same as the conservation of energy.
The transfer of heat through the motion of hot fluid from oneplace to another. In physics, a "fluid" can be either a gas or a liquid, and convection is the principal means of heat transfer, for instance, in air and water.
The ability to accomplishwork.
The tendency of natural systems toward breakdown, and specifically, the tendency for the energy in a system to be dissipated. Entropy is closely related to the second law of thermodynamics.
FIRST LAW OF THERMODYNAMICS:
A law stating that the amount of energy in a system remains constant, and therefore, it is impossible to perform work that results in an energy, output greater than the energy input. This is the same as the conservation of energy.
The principal unit of energy—and, thus, of heat—in the British or English system. The metric or SI unit is the joule. A foot-pound (ft · lb) is equal to 1.356 J.
Internal thermal energy that flows from one body of matter to another. Heat is transferred by three methods conduction, convection, and radiation.
A machine that absorbs heat at a high temperature, performs mechanical work, and, as a result, gives off heat at a lower temperature.
The principal unit of energy—and, thus, of heat—in the SI or metric system, corresponding to 1 newton-meter (N · m). A joule (J) is equal to 0.7376 foot-pounds.
Established by William Thomson, Lord Kelvin (1824-1907), the Kelvin scale measures temperature in relation to absolute zero, or 0K.(Units in the Kelvin system, known as Kelvins, do not include the word or symbol for degree.) The Kelvin and Celsius scales are directly related; hence, Celsius temperatures can be converted to Kelvins by adding 273.15.
A measure of heat or energy in the SI or metric system, equal to the heat that must be added to or removed from 1 kilogram of water to change its temperature by 1°C. As its name suggests, a kilocalorie is 1,000 calories. The dietary Calorie (capital C) with which most people are familiar, is the same as the kilocalorie.
The energy that an object possesses by virtue of its motion.
The energy that an object possesses due to its position.
The transfer of heat by means of electromagnetic waves, which require no physical medium (for example, water or air) for the transfer. Earth receives the Sun's heat by means of radiation.
SECOND LAW OF THERMODYNAMICS:
A law of thermodynamics stating that no engine can be constructed that simply takes heat from a source and performs an equivalent amount of work. Some of the heat will always be lost, and, therefore, it is impossible to build a perfectly efficient engine. This is a result of the fact that the natural flow of heat is always from a high-temperature reservoir to a low-temperature reservoir—a fact expressed in the concept of entropy. The second law is sometimes referred to as "the law of entropy."
The amount of heat that must be added to, or removed from, a unit of mass of a given substance to change its temperature by 1°C. A kilocalorie is the specific heat of 1 gram of water.
In physics, the term "system" usually refers to any set of physical interactions isolated from the rest of the universe. Anything outside of the system, including all factors and forces irrelevant to a discussion of that system, is known as the environment.
The direction of internal energy flow between two systems when heat is being transferred. Temperature measures the average molecular kinetic energy in transit between those systems.
Heat energy, a form of kinetic energy produced by the movement of atomic or molecular particles. The greater the movement of the separticles, the greater the thermal energy.
The statethat exists when two systems have the same temperature. As a result, there is no exchange of heat between them.
The study of the relationships between heat, work, and energy.
THIRD LAW OF THERMODYNAMICS:
A law of thermodynamics which states that at the temperature of absolute zero, entropy also approaches zero. Zero entropy would contradict the second law of thermodynamics, meaning that absolute zerois, therefore, impossible to reach.
The exertion of force over a given distance to displace or move an object. Work is, thus, the product of force and distance exerted in the same direction.
Heat, within the science of physics, is defined as the transfer of thermal energy from one part of a material to another part of a material, or from one body to another body.
Heat exchange reflects and drives changes in energy state between two objects—or more generally systems—in thermal contact due to a difference in temperature. Heat flows from a system at higher temperature to one at lower temperature until both systems are at the same temperature. Systems at the same temperature are said to be in thermal equilibrium.
The term heat is sometimes used, incorrectly, to refer to a form of energy that a system contains. Heat is a form of energy-in-transit; it is not energy-in-resi-dence. The energy contained in a system (exclusive of energy depending on external factors) is called internal energy and, unlike heat, is a property of a system like the volume or mass.
The concept of heat has always been an important consideration in scientific thought as well as with everyday life. Some things are hot, while others are warm and still others are cold. When hot or warm objects are placed in contact with cold ones, the warmer object cools down as the colder one heats up. Apparently, something causes an object to be warm and that something flows from warm objects to cold ones. This something is called heat.
The controversy over whether heat is a material substance or is due to the kinetic motion of particles continued into the early twentieth century. English physicist and mathematician Sir Isaac Newton (1642–1727), Dutch mathematician Christiaan Huygens (1629–1695), English scientist Robert Hooke (1635–1703), English chemist and physicist Henry Cavendish (1731–1810), English chemist Sir Humphry Davy (1778–1829), and American physicist Benjamin Thompson, Count Rumford (1753–1814) supported the kinetic interpretation. Dutch chemist and physician Hermann Boerhaave (1668–1738), French chemist Antoine-Laurent Lavoisier (1743–1794) and French mathematician Pierre-Simon Laplace (1749–1827) favored the material interpretation.
Scottish chemist and physicist Joseph Black (1728–1799) and Swedish physicist Johann Wilcke (1732–1779) developed the concept of specific heat capacity, defined as the amount of heat that raises the temperature of a substance by one degree. The empirical definition of heat is based on this concept: the amount of heat involved in a process is equal to the product of the specific heat of the substance multiplied by the change in temperature that occurs when heat is added to or taken away from the substance. The unit given to the amount of heat is the calorie, a metric unit, defined as the amount of heat that raises the temperature of one gram of water by one degree of Celsius. It is also represented by the British thermal unit (Btu), an English unit. One Btu is defined as the amount of heat that raises the temperature of one pound of water by one degree Fahrenheit. Specific heat is the same for any substance, whether it is defined by the calorie or the Btu. In addition, scientists often represent heat with respect to the joule, a unit used to represent all types of energy.
English physicist James Joule (1818-1889) showed that, unlike material substances, heat can be created and destroyed by changing it into work. One calorie of heat is equivalent to about 4.2 joules of work. Subsequently, the kinetic theory proposed by German scientists August K. Kronig (1822-1879) and Rudolf Clausius (1822-1888) explained heat as the result of translational, rotational, and vibrational motions of molecules: the faster they move, the higher the temperature. The work of German-American Albert Einstein (1879-1955) finally settled the argument in favor of the kinetic interpretation of heat.
The first law of thermodynamics states that the internal energy of a system can change only if energy flows into or out of the system. This flow, or energy-in-transit, appears as heat or as work (or a combination), and the change in internal energy is equal to the total of heat and work appearing during the change. After the change, however, the system contains neither heat nor work; it contains internal energy.
Units of heat are units of energy. One classical unit, the calorie, is defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius. A more precise definition recognizes that this energy depends slightly on the temperature of the water, so the interval was specified as 14.5 to 15.5°C (58.1-59.9°F). The dietary calorie (capital C) is a kilocalorie (1,000 calories). The energy available from the metabolism of a given amount of food is commonly given in calories.
In the International System of Units (SI or extended metric system) the joule is the unit of energy. Although based on mechanical rather than thermal considerations, the joule is now the preferred energy unit for both mechanical and thermal applications. The joule is about 1/4 of a calorie and, now, formally defines the calorie. One calorie is by definition exactly 4.184 joules, although the practical difference between this definition and the original one is negligible.
The specific heat capacity, or specific heat, is the heat required to raise the temperature of one gram of substance one degree Celsius. The specific heats of a few substances in joules per gram per degree Celsius are listed in Table 1.
For example, to raise the temperature of equal amounts of all four of these substances, the water
|Table 1. (Thomson Gale.)|
|Substance||Specific heat (J/g°C) at 25°C|
would require considerably more heat than the others (over 9 times as much as the iron, for example, because 4.18 divided by 0.45 is 9.3). Alternatively, if one added the same amount of heat to equal amounts of all four of these substances, the temperature of the water would rise least. In short, it is more difficult to change the temperature of water than most other substances. This is one of the main reasons coastal climates usually have smaller seasonal temperature variations than inland climates. Because of its relatively high specific heat, water is a good thermal moderator.
Heat is the energy that flows between two objects because of a difference in temperature. Heat always flows from a body at a higher temperature to one at a lower temperature.
Scientists use the term heat differently than do nonscientists. The average person may think of heat as the amount of energy contained in a body. The correct term for that property, however, is thermal energy.
Thermal energy and temperature
According to the kinetic theory of matter, all matter is composed of particles that are constantly in motion. Temperature is a measure of the motion of those particles. The more rapidly particles are in motion, the higher the temperature; the less rapidly they are moving, the lower the temperature.
In theory, it would be possible to reduce the motion of the particles in an object to zero. In that case, the object would contain no thermal energy. The temperature at which all particle motion ends is called absolute zero. Scientists have come within a few millionths of a degree of absolute zero but have never actually reached that point.
Another way to think of heat is as a transfer of thermal energy from one place to another. This process occurs in one of three ways: conduction, convection, and radiation.
Conduction. Conduction is the process of heat transfer. Rapidly moving molecules in a hot material collide with slower moving molecules in a cool material. The fast-moving molecules slow down and the slow moving molecules increase their speed. Conduction occurs, then, when two bodies of different temperatures are in contact with each other.
Convection. Convection is the process by which large masses of a fluid (a liquid or gas) move, carrying thermal energy. When water in a container is heated, for example, it expands. Cooler water around it pushes the lighter water upward. As the warm water rises, it begins to cool and starts to move downward in the liquid again. Eventually, a circular motion is produced within the liquid, forcing heat to be transferred throughout the liquid.
Radiation. Finally, thermal energy can be transferred by radiation. Hot bodies emit electromagnetic radiation that corresponds to their temperature. This radiation passes through space until it comes into contact with a body with less thermal energy. The cooler body then absorbs this radiation and becomes warmer.
Transfer of thermal energy. The transfer of thermal energy from one place to another occurs in one of three ways: conduction, convection, and radiation. In conduction, rapidly moving molecules in a hot material collide with slower moving molecules in a cool material. The fast-moving molecules slow down and the slow-moving molecules increase their speed. Conduction occurs, then, when two bodies of different temperatures are in contact with each other.
Convection is the process by which large masses of a fluid (a liquid or gas) move, carrying thermal energy. When water in a container is heated, for example, it expands. Cooler water around it pushes the lighter water upward. As the warm water rises, it begins to cool and starts to move downward in the liquid again. Eventually, a circular motion is produced within the liquid, forcing heat to be transferred throughout the liquid.
Finally, thermal energy can be transferred by radiation. Hot bodies emit electromagnetic radiation that corresponds to their temperature. This radiation passes through space until it comes into contact with a body with less thermal energy. The cooler body then absorbs this radiation and becomes warmer.
Since heat is a form of energy, the units used to measure heat are the same as those used to measure energy. In the metric system, one of the earliest units used to measure heat was the calorie. The calorie is defined as the amount of heat energy needed to raise the temperature of one gram of water one degree Celsius. To be precise, the temperature change is specified as an increase from 14.5°C to 15.5°C.
In the International System of Units (the SI system), the unit of energy is the joule. A calorie is defined as 4.184 joules.
Materials differ from each other with regard to how easily they can be warmed. One could add a joule of heat to a gram of water, a gram of iron, a gram of mercury, and a gram of ethyl alcohol and notice very different results. The temperature of the mercury would rise the most, and the temperature of the water would rise the least.
The specific heat capacity (or just specific heat) of a material is defined as the amount of heat required to raise the temperature of one gram of the material one degree Celsius. It takes 4.18 joules to raise the temperature of 1 gram of water 1 degree Celsius (at a temperature of 25°C). In comparison, it takes only 0.14 joule to raise the temperature of the same amount of mercury by one degree Celsius and 0.45 joule to raise the temperature of the same amount of iron by one degree Celsius. It takes 2.46 joules to raise the temperature of the same amount of ethyl alcohol by one degree Celsius.
[See also Energy; Temperature; Thermodynamics ]
heat, nonmechanical energy in transit, associated with differences in temperature between a system and its surroundings or between parts of the same system.
Measures of Heat
Temperature is a measure of the average translational kinetic energy of the molecules of a system. Heat is commonly expressed in either of two units: the calorie, an older metric unit, and the British thermal unit (Btu), an English unit commonly used in the United States. Scientists express heat in terms of the joule, a unit used for all forms of energy.
As heat is added to a substance in the solid state, the molecules of the substance gain kinetic energy and the temperature of the substance rises. The amount of heat needed to raise a unit of mass of the substance one degree of temperature is called the specific heat of the substance. Because of the way in which the calorie and the Btu are defined, the specific heat of any substance is the same in either system of measurement. For example, the specific heat of water is 1 calorie per gram per degree Celsius; i.e., 1 calorie of heat is needed to raise the temperature of 1 gram of water by 1 degree Celsius; it is also 1 Btu per pound per degree Fahrenheit.
Heat of Fusion
When a solid reaches a certain temperature, it changes to a liquid. This temperature is a particular property of the substance and is called its melting point. While the solid-liquid transition is taking place, there is no change in temperature. All of the heat being added is being converted to the internal potential energy associated with the liquid state. The amount of heat needed to convert one unit of mass of a substance from a solid to liquid is called the heat of fusion, or latent heat of fusion, of the substance. Like specific heat, latent heat is also a property of the particular substance. The latent heat of fusion for the ice-to-water transition is 80 calories per gram.
Heat of Vaporization
After a substance is completely changed from a solid to a liquid, further addition of heat again causes the temperature to rise until it reaches the boiling point, the particular temperature at which the given substance changes from a liquid to a gas. During the liquid-gas transition, the temperature remains constant until the change is completed. The heat of vaporization, or latent heat of vaporization, is the heat that must be added to convert one unit of mass of the substance from a liquid to a gas.
Transfer of Heat
Heat may be transferred from one substance to another by three means—conduction, convection, and radiation. Conduction involves the transfer of energy from one molecule to adjacent molecules without the substance as a whole moving. Convection involves the movement of warmer parts of a substance away from the source of heat and takes place only in fluids, i.e., liquids and gases. Radiation is the transfer of heat energy in the form of electromagnetic radiation, principally in the infrared radiation portion of the spectrum.
Study and Analysis of Heat
The study of heat and its relationship to useful work is called thermodynamics and involves macroscopic quantities such as pressure, temperature, and volume without regard for the molecular basis of these quantities. Low-temperature physics is concerned with phenomena that occur at extremely low temperatures. The analysis of heat on the basis of the structure of matter is considered in the kinetic-molecular theory of gases and provides an explanation for the various gas laws. The gas laws in turn serve to define an absolute temperature scale based on theoretical considerations (see Kelvin temperature scale).
See M. C. Mott-Smith, Heat and Its Workings (1933, repr. 1962); R. Becker, Theory of Heat (tr. 1967).
See also 90. COLD ; 162. FIRE
- the ability of light and heat and other forms of radiant energy to cause chemical changes, as hormonal changes in birds causing them to migrate or brood. —actinic , adj.
- imperviousness to radiant heat or infrared radiation. Also called athermancy .
- the science of measuring heat. —calorimeter , n. —calorimetric , adj.
- Rare. one who believed the caloric theory, that heat is a material substance. —caloristic , adj.
- the process of generating heat by means of an electric current.
- the branch of geology that measures temperatures deep below the surface of the earth; geologic thermometry.
- the production or generation of heat. —pyrogenetic , adj.
- the chemical process of decomposition under the effect of heat. —pyrolitic , adj.
- a type of pyrometer that measures temperature optically or photometrically.
- a moderate warmth; lukewarmness. —tepid , adj.
- Medicine. the study of heat as a medical remedy or therapy. Also called thermotherapy .
- the science or study of the emission of electrons from substances at high temperatures. —thermionic , adj.
- the branch of chemistry that studies the relationship of heat to chemical changes, including the production of energy. —thermochemist , n. —thermochemical , adj.
- the branch of physics that studies the relationship of heat and mechanical energy and the conversion, in various materials, of one into the other. —thermodynamicist , n. —thermodynamic, thermodynamical , adj.
- thermogenesis, thermogeny
- the production of heat, especially in an animal body by physiological processes. —thermogenic, thermogenous , adj.
- 1. Engineering, a method of measuring surf ace temperatures by using luminescent materials.
- 2. a printing or photocopying process using infrared rays and heat.
- 3. a process of photography using far-infrared radiation; thermal photography. —thermographer , n. —thermographic , adj.
- the study of the movement of heat. —thermokinematic , adj.
- Archaic. the science and study of heat. Also called thermotics .
- Atomic Physics. any luminescence appearing in materials upon application of heat, caused by electron movement which increases as the temperature rises. —thermoluminescent , adj.
- Physiology. the dispersion of heat from the body. —thermolytic , adj.
- the branch of physics that deals with the measurement of temperature. —thermometric , adj.
- an abnormal fear of heat.
- a device for giving an approximation of the temperature change of a substance by noting the accompanying change in its volume. —thermoscopic , adj.
- the science or study of the equilibrium of heat.
- 1. Biology. the movement of an organism toward or away from a source of heat.
- 2. Physiology. the regulation of body temperature by various physiological processes. —thermotactic, thermotaxic , adj.
- the property or quality by which matter permits the passage of heat. —transcalent , adj.
heat / hēt/ • n. 1. the quality of being hot; high temperature: it is sensitive to both heat and cold. ∎ hot weather conditions: the oppressive heat was making both men sweat. ∎ a source or level of heat for cooking: remove from the heat and beat in the butter. ∎ a spicy quality in food that produces a burning sensation in the mouth: chili peppers add taste and heat to food. ∎ Physics heat seen as a form of energy arising from the random motion of the molecules of bodies, which may be transferred by conduction, convection, or radiation. ∎ technical the amount of heat that is needed to cause a specific process or is evolved in such a process: the heat of formation. ∎ technical a single operation of heating something, esp. metal in a furnace. 2. intensity of feeling, esp. of anger or excitement: words few men would dare use to another, even in the heat of anger. ∎ (the heat) inf. intensive and unwelcome pressure or criticism, esp. from the authorities: a flurry of legal proceedings turned up the heat in the dispute. 3. a preliminary round in a race or contest: the 200-meter heats. • v. make or become hot or warm: [tr.] the room faces north and is difficult to heat | [intr.] the pipes expand as they heat up. ∎ [intr.] (heat up) (of a person) become excited or impassioned. ∎ [intr.] (heat up) become more intense and exciting: the action really begins to heat up. ∎ [tr.] archaic inflame; excite: this discourse had heated them. PHRASES: in the heat of the moment while temporarily angry, excited, or engrossed, and without stopping for thought. in heat (of a female mammal) in the receptive period of the sexual cycle; in estrus.
Heat is the transfer of energy that results from the difference in temperature between a system and its surroundings. At a molecular level, heat is the transfer of energy that makes use of or stimulates disorderly molecular motion in the surroundings. For instance, when a hydrocarbon fuel burns, the energy released in the reaction stimulates the surrounding atoms and molecules into more vigorous random motion, and we refer to this escape of energy as heat. Heat is not stored: Heat is energy in transit.
The measurement of quantities of energy transferred as heat is called calorimetry. Such a measurement is commonly made by observing the rise in temperature caused by the process being studied and interpreting that rise in terms of the heat produced. Calorimetry is used to measure the changes in internal energy and enthalpy that accompany chemical reactions. The field of study is called thermochemistry, and it is used to assess the efficacy of fuels, the energy flow in chemical plants, and the strengths of chemical bonds. Measurements of the heat produced or absorbed by chemical reactions are central to thermodynamics, and to assessments of whether or not a particular reaction will tend to occur.
In thermodynamics, the quantity of energy transferred as heat as a result of a chemical reaction is identified with the change in the internal energy of the system if the transfer takes place without change in the system's volume, and with the change in enthalpy of the system if the transfer takes place at constant pressure. The energy or enthalpy change accompanying a chemical reaction that is inaccessible to measurement may be determined by using Hess's law, which states that the enthalpy change accompanying a chemical reaction can be regarded as the sum of the enthalpy changes of the reactions into which the overall reaction may be divided. Hess's law is no more than a special application of the first law of thermodynamics.
The source of heat as a fuel burns is the energy released when the bonds characteristic of the reactants are replaced by the bonds characteristic of the products. Energy is released when hydrocarbons burn because of the great strengths of the oxygen–hydrogen and oxygen–carbon bonds that are formed in the products (water and carbon dioxide), replacing the relatively weak carbon–hydrogen and carbon–carbon bonds of the fuel. Ultimately, the energy of burning fuel is the energy released as the electrons and atomic nuclei settle into more favorable arrangements (just as nucleons do in the much more exothermic processes accompanying nuclear rearrangements).
Although the term "heat energy" is commonly encountered in casual conversation, strictly speaking there is no such entity. The term is commonly used in place of the more precise term "energy of thermal motion," where thermal motion is random molecular motion, as in the motion of molecules in a gas. Nor is heat stored: Only energy is stored, and heat is one of the modes by which it may be increased or extracted.
see also Chemistry and Energy; Energy; Explosions; Temperature; Thermochemistry; Thermodynamics.
Smith, Crosbie (1998). The Science of Energy: A Cultural History of Energy Physics in Victorian Britain. Chicago: University of Chicago Press.
Heat exchange reflects and drives changes in energy state between two objects—or more generally systems—in thermal contact due to a difference in temperature . Heat flows from a system at higher temperature to one at lower temperature until both systems are at the same temperature. Systems at the same temperature are said to be in thermal equilibrium.
The term "heat" is sometimes used, incorrectly, to refer to a form of energy that a system contains. Heat is a form of energy-in-transit; it is not energy-in-residence. The energy contained in a system (exclusive of energy depending on external factors) is called internal energy and, unlike heat, is a property of a system like the volume or mass .
The first law of thermodynamics states that the internal energy of a system can change only if "energy" flows into or out of the system. This flow, or energy-in-transit, appears as heat or as work (or a combination), and the change in internal energy is equal to the total of heat and work appearing during the change. After the change, however, the system contains neither heat nor work; it contains internal energy.
Units of heat are units of energy. One classical unit, the calorie , was defined as the amount of energy required to raise the temperature of one gram of water one degree Celsius. A more precise definition recognizes that this energy depends slightly on the temperature of the water, so the interval was specified as 14.5–15.5°C (58.1–59.9°F). The dietary Calorie (capital C) is a kilocalorie (1000
|SUBSTANCE||SPECIFIC HEAT (J/g oC) at 25 oC|
calories). The energy available from the metabolism of a given amount of food is commonly given in Calories.
In the International System of Units (SI or extended metric system ) the joule is the unit of energy. Although based on mechanical rather than thermal considerations, the joule is now the preferred energy unit for both mechanical and thermal applications. The joule is about 1/4 of a calorie and now formally defines the calorie. One calorie is by definition exactly 4.184 joules, although the practical difference between this definition and the original one is negligible.
The specific heat capacity , or specific heat, is the heat required to raise the temperature of one gram of substance one degree Celsius. The specific heats of a few substances in joules per gram per degree Celsius are listed above.
For example, to raise the temperature of equal amounts of all four of these substances, the water would require considerably more heat than the others (over 9 times as much as the iron , for example, because 4.18 divided by 0.45 is 9.3). Or if you added the same amount of heat to equal amounts of all four of these substances, the temperature of the water would rise least. In short, it is more difficult to change the temperature of water than most other substances. This is one of the main reasons coastal climates usually have smaller seasonal temperature variations than inland climates. Because of its relatively high specific heat, water is a good thermal moderator.