If you take a can of cola from the refrigerator and leave it on the kitchen table, its temperature will rise—rapidly at first but then more slowly—until the temperature of the cola equals that of the room (the two are then in thermal equilibrium). In the same way, the temperature of a cup of hot coffee, left sitting on the table, will fall until it also reaches room temperature.
In generalizing this situation, we describe the cola or the coffee as a system (with temperature TS) and the relevant parts of the kitchen as the environment (with temperature TE) of that system. Our observation is that if TS is not equal to TE, then TS will change (TE can also change some) until the two temperatures are equal and thus thermal equilibrium is reached.
Such a change in temperature is due to a change in the thermal energy of the system because of a transfer of energy between the system and the system’s environment. (Recall that thermal energy is an internal energy that consists of the kinetic and potential energies associated with the random motions of the atoms, molecules, and other microscopic bodies within an object.) The transferred energy is called heat and is symbolized Q. Heat is positive when energy is transferred to a system’s thermal energy from its environment (we say that heat is absorbed by the system). Heat is negative when energy is transferred from a system’s thermal energy to its environment (we say that heat is released or lost by the system).
This transfer of energy is shown in Fig. 18-12. In the situation of Fig. 18-12a, in which TS > TE, energy is transferred from the system to the environment, so Q is negative. In Fig. 18-12b, in which TS = TE, there is no such transfer, Q is zero, and heat is neither released nor absorbed. In Fig. 18-12c, in which TS < TE, the transfer is to the system from the environment; so Q is positive.
We are led then to this definition of heat:
Fig. 18-12 If the temperature of a system exceeds that of its environment as in (a), heat Q is lost by the system to the environment until thermal equilibrium (b) is established. (c) If the temperature of the system is below that of the environment, heat is absorbed by the system until thermal equilibrium is established.
Heat is the energy transferred between a system and its environment because of a temperature difference that exists between them.
Recall that energy can also be transferred between a system and its environment as work W via a force acting on a system. Heat and work, unlike temperature, pressure, and volume, are not intrinsic properties of a system. They have meaning only as they describe the transfer of energy into or out of a system. Similarly, the phrase “a $600 transfer” has meaning if it describes the transfer to or from an account, not what is in the account, because the account holds money, not a transfer. Here, it is proper to say: “During the last 3 min, 15 J of heat was transferred to the system from its environment” or “During the last minute, 12 J of work was done on the system by its environment.” It is meaningless to say: “This system contains 450 J of heat” or “This system contains 385 J of work.”
Before scientists realized that heat is transferred energy, heat was measured in terms of its ability to raise the temperature of water. Thus, the calorie (cal) was defined as the amount of heat that would raise the temperature of 1 g of water from 14.5°C to 15.5°C. In the British system, the corresponding unit of heat was the British thermal unit (Btu), defined as the amount of heat that would raise the temperature of 1 lb of water from 63°F to 64°F.
In 1948, the scientific community decided that since heat (like work) is transferred energy, the SI unit for heat should be the one we use for energy—namely, the joule. The calorie is now defined to be 4.1868 J (exactly), with no reference to the heating of water. (The “calorie” used in nutrition, sometimes called the Calorie (Cal), is really a kilocalorie.) The relations among the various heat units are
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