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.
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.
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.
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.
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.
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.
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 = mcδT. 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.
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.
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.
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.
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.
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 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.
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.
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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.
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.
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.
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.
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.
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.