Warning: In case you have arrived at this page directly from a search engine, it is not exactly what it might seem to be.  This is an illustration of bad (teaching)
explanations in science, not a critique of thermodynamics. This paper comes from a link in a paper about bad explanations at www.garlikov.com/science/sciteach.htm

The black print is from a Funk and Wagnall's Knowledge Center Web Page about Thermodynamics: http://www.fwkc.com/encyclopedia/low/articles/t/t025000863f.html

The red print contains my comments.

THERMODYNAMICS, field of physics that describes and correlates the physical properties of macroscopic systems of matter and energy as they are understood at any particular time and thought to be true.  "Laws" in physics are about our observations and constructs; we think they apply to nature, but they may not.  The principles of thermodynamics are of fundamental importance to all branches of science and engineering. 

A central concept of thermodynamics is that of the macroscopic system, defined as a geometrically isolable piece of matter in coexistence with an infinite, unperturbable environment. The state of a macroscopic system in equilibrium can be described in terms of such measurable properties as temperature, pressure, and volume, which are known as thermodynamic variables. These properties are constructs. They may or may not correlate with how we perceive their everyday language equivalents.  For example, "temperature" is not about how cold or warm something feels to human touch, but is about a construct that has to do with certain effects on thermometers, and the theoretical basis, as it is understood, for those effects. Many other variables (such as density, specific heat, compressibility, and the coefficient of thermal expansion) can be identified and correlated, to produce a more complete description of an object and its relationship to its environment. 

When a macroscopic system moves from one state of equilibrium to another, a thermodynamic process is said to take place. Some processes are reversible and others are irreversible. The laws of thermodynamics, discovered in the 19th century through painstaking experimentation, govern the nature of all thermodynamic processes and place limits on them.  "Govern" is the totally wrong word to use here, because it implies that nature has to act in accordance with these laws, whereas these laws (as all scientific "laws") are only descriptions of what is observed or of what, after much deliberation, seems to be observed, or of what seems to account for what is observed.  E.g., Newton's law that "an object in motion will maintain that motion unless acted upon by a force..." is not ever really observed.  What we do is posit forces in order to account for changes in motion.  Doing so, leads in various cases to our being able to predict and manipulate phenomena, but there is always the possibility that we are working within specific or 'local' conditions or perceptions which we mistakenly take to be universal, but outside of which our laws either do not apply, or do not give the whole picture.  E.g., Newton's Laws and "classical mechanics" do not describe phenomena that is at either great speeds/accelerations or tremendously small (quantum) scale.

Zeroth Law of Thermodynamics. 
The vocabulary of empirical sciences is often borrowed from daily language. Thus, although the term temperature appeals to common sense, its meaning suffers from the imprecision of nonmathematical language. A precise, though empirical, definition of temperature is provided by the so-called zeroth law of thermodynamics as explained below. 

When two systems are in equilibrium, they share a certain property. This property can be measured and a definite numerical value ascribed to it. A consequence of this fact is the zeroth law of thermodynamics, which states that when each of two systems is in equilibrium with a third, the first two systems must be in equilibrium with each other. This shared property of equilibrium is the temperature. This is either incomplete or it is not intuitively about temperature, for it is not clear that what we consider to be temperature is the only property that would meet this criteria. For example, if two objects have the same color as a third object, they will have the same color as each other.  Similarly mass, shape, size, etc. 

If any such system is placed in contact with an infinite environment that exists at some certain temperature, the system will eventually come into equilibrium with the environment-that is, reach the same temperature. (The so-called infinite environment is a mathematical abstraction called a thermal reservoir; in reality the environment need only be large relative to the system being studied.) 

Temperatures are measured with devices called thermometers. A thermometer contains a substance with conveniently identifiable and reproducible states, such as the normal boiling and freezing points of pure water. If a graduated scale is marked between two such states, the temperature of any system can be determined by having that system brought into thermal contact with the thermometer, provided that the system is large relative to the thermometer. This is an operational definition of "temperature" -- the thermometer is not really "measuring" anything; it is registering an effect of being in proximity of the object being examined for its causal effects on the thermometer.  Because different thermometer temperatures feel to human touch as being warm, hot, cold, tepid, or whatever, we consider thermometers to measure what we are feeling, but they do not, because they cannot feel what we are feeling.  What we feel when we touch something warm or hot is a human response or perception or result of being in proximity with an object.

First Law of Thermodynamics. 
The first law of thermodynamics gives a precise definition of heat, another commonly used concept. 

When an object is brought into contact with a relatively colder object, a process takes place that brings about an equalization of temperatures of the two objects. To explain this phenomenon, 18th-century scientists hypothesized that a substance more abundant at higher temperature flowed toward the region at a lower temperature. This hypothetical substance, called "caloric," was thought to be a fluid capable of moving through material media. The first law of thermodynamics instead identifies caloric, or heat, as a form of energy. It can be converted into mechanical work, and it can be stored, but is not a material substance. Heat, measured originally in terms of a unit called the calorie, and work and energy, measured in ergs, were shown by experiment to be totally equivalent. One calorie is equivalent to 4.186 ´ 107 ergs, or 4.186 joules. 

The first law, then, is a law of energy conservation. It states that, because energy cannot be created or destroyed --setting aside the later ramifications of the equivalence of mass and energy-- the amount of heat transferred into a system plus the amount of work done on the system must result in a corresponding increase of internal energy in the system. Heat and work are mechanisms by which systems exchange energy with one another. 

In any machine some amount of energy is converted into work; therefore, no machine can exist in which no energy is converted into work. Such a hypothetical machine (in which no energy is required for performing work) is termed a "perpetual-motion machine of the first kind." Since the input energy must now take heat into account (and in a broader sense chemical, electrical, nuclear, and other forms of energy as well), the law of energy conservation rules out the possibility of such a machine ever being invented. The first law is sometimes given in a contorted form as a statement that precludes the existence of perpetual-motion machines of the first kind. 

Second Law of Thermodynamics. 
The second law of thermodynamics gives a precise definition of a property called entropy. Entropy can be thought of as a measure of how close a system is to equilibrium; it can also be thought of as a measure of the disorder in the system. Notice, however, that these are not precise and don't even seem to be definitions, not clear ones at least.  The law states that the entropy-that is, the disorder-of an isolated system can never decrease. Thus, when an isolated system achieves a configuration of maximum entropy, it can no longer undergo change: It has reached equilibrium. Nature, then, seems to "prefer" disorder or chaos. Yet living organisms are things which are ordered that have developed from disorder or chaos. And if the universe undergoes expansions and collapses, are not the collapses a turn or 'return' to some sort of order?  Does not gravity, acting upon things of different density, act to separate those things into "ordered" layers? Of course, if you put cream into coffee, the cream molecules disperse among the coffee molecules; but if you put olive oil into coffee and stir it up, the oil molecules will separate from the coffee molecules.  Why is the cream phenomenon somehow a model of a universal law while the olive oil phenomenon not?  It can be shown that the second law stipulates that, in the absence of work, heat cannot be transferred from a region at a lower temperature to one at a higher temperature. No, it simply stipulates that this has never been observed to happen.

The second law poses an additional condition on thermodynamic processes. It is not enough to conserve energy and thus obey the first law. A machine that would deliver work while violating the second law is called a "perpetual-motion machine of the second kind," since, for example, energy could then be continually drawn from a cold environment to do work in a hot environment at no cost. The second law of thermodynamics is sometimes given as a statement that precludes perpetual-motion machines of the second kind. Which just means that the law cannot be true and there exist (the possibility of) a perpetual motion macine of the second kind.  The law is incompatible with the existence of such a machine, but the existence of such a machine would then invalidate the law.  When scientific laws are broken, there is no penalty; the breakage merely invalidates the law. In jurisprudence "laws" are meant to govern and supercede behavior, but in science, observed  phenomena engender, determne, govern and supercede laws.

Thermodynamic Cycles. 
All important thermodynamic relations used in engineering are derived from the first and second laws of thermodynamics. One useful way of discussing thermodynamic processes is in terms of cycles-processes that return a system to its original state after a number of stages, thus restoring the original values for all the relevant thermodynamic variables. In a complete cycle the internal energy of a system depends solely on these variables and cannot change. Thus, the total net heat transferred to the system must equal the total net work delivered from the system. 

An ideal cycle would be performed by a perfectly efficient heat engine-that is, all the heat would be converted to mechanical work. The 19th-century French scientist Nicolas Léonard Sadi Carnot, who conceived a thermodynamic cycle that is the basic cycle of all heat engines, showed that such an ideal engine cannot exist. Any heat engine must expend some fraction of its heat input as exhaust. The second law of thermodynamics places an upper limit on the efficiency of engines; that upper limit is less than 100 percent. The limiting case is now known as a Carnot cycle. 

Third Law of Thermodynamics. 
The second law suggests the existence of an absolute temperature scale that includes an absolute zero of temperature. The third law of thermodynamics states that absolute zero cannot be attained by any procedure in a finite number of steps. Absolute zero can be approached arbitrarily closely, but it can never be reached. 

Microscopic Basis of Thermodynamics. 
The recognition that all matter is made up of molecules provided a microscopic foundation for thermodynamics. A thermodynamic system consisting of a pure substance can be described as a collection of like molecules, each with its individual motion describable in terms of such mechanical variables as velocity and momentum. At least in principle, it should therefore be possible to derive the collective properties of the system by solving equations of motion for the molecules. In this sense, thermodynamics could be regarded as a mere application of the laws of mechanics to the microscopic system. 

Objects of ordinary size-that is, ordinary on the human scale-contain immense numbers (on the order of 1024) of molecules. Assuming the molecules to be spherical, each would need three variables to describe its position and three more to describe its velocity. Describing a macroscopic system in this way would be a task that even the largest modern computer could not manage. A complete solution of these equations, furthermore, would tell us where each molecule is and what it is doing at every moment. Such a vast quantity of information would be too detailed to be useful and too transient to be important. 

Statistical methods were devised therefore to obtain averages of the mechanical variables of the molecules in a system and to provide the gross features of the system. These gross features turn out to be, precisely, the macroscopic thermodynamic variables. The statistical treatment of molecular mechanics is called statistical mechanics, and it anchors thermodynamics to mechanics. 

Viewed from the statistical perspective, temperature represents a measure of the average kinetic energy of the molecules of a system. Increases in temperature reflect increases in the vigor of molecular motion. When two systems are in contact, energy is transferred between molecules as a result of collisions. The transfer will continue until uniformity is achieved, in a statistical sense, which corresponds to thermal equilibrium. The kinetic energy of the molecules also corresponds to heat and-together with the potential energy arising from interaction between molecules-makes up the internal energy of a system. 

The conservation of energy, a well-known law of mechanics , translates readily to the first law of thermodynamics, and the concept of entropy translates into the extent of disorder on the molecular scale. By assuming that all combinations of molecular motion are equally likely, thermodynamics shows that the more disordered the state of an isolated system, the more combinations can be found that could give rise to that state, and hence the more frequently it will occur. The probability of the more disordered state occurring overwhelms the probability of the occurrence of all other states. This probability provides a statistical basis for definitions of both equilibrium state and entropy. But, if it is just a matter of probability, this should mean that somewhere in the world at some time, someone should have a chance of noticing the cream go back into the center of the cup of coffee spontaneously, no? 

Finally, temperature can be reduced by taking energy out of a system, that is, by reducing the vigor of molecular motion. Absolute zero corresponds to the state of a system in which all its constituents are at rest. This is, however, a notion from classical physics. In terms of quantum mechanics, residual molecular motion will exist even at absolute zero. An analysis of the statistical basis of the third law goes beyond the scope of the present discussion. 


Reset June 24, 2000