First Two Laws of Thermodynamics

Let us briefly review some fundamentals of thermodynamics. All organisms require energy to persist and to replace themselves, and the ultimate source of practically all Earth's energy is the Sun. One can think of our Sun as "feeding" the earth via its radiant energy. But 99 percent or more of this incident solar radiation goes unused by organisms and is lost as heat and heat of evaporation; only about 1 percent is actually captured by plants in photosynthesis and stored as chemical energy. Moreover, energy available from sunlight varies widely over the earth's surface both in space and in time.

Physics and chemistry have produced two basic laws of thermodynamics that are obeyed by all forms of matter and energy, including living organisms.

The first law is that of "conservation of matter and energy, " which states that matter and energy cannot be created or destroyed. Matter and energy can be transformed, and energy can be converted from one form into another, but the total of the equivalent amounts of both must always remain constant. Light can be changed into heat, kinetic energy, and/or potential energy. Whenever energy is converted from one form into another, some of it is given off as heat, which is the most random form of energy. Indeed, the only energy conversion that is 100 percent efficient is conversion to heat, or burning. Aliquots of dried organisms can be burned in "bomb calorimeters" to determine how much energy is stored in their tissues. Energy can be measured in a variety of different units such as ergs and joules, but heat energy or calories is the common denominator.

The second law of thermodynamics states that energy of all sorts, whether it be light, potential, chemical, kinetic, or whatever, tends to change itself spontaneously into a more dispersed, random, or less organized, form. This law is sometimes stated as "entropy increases" - entropy being random, unavailable energy. Suppose you heat a skillet to cook an egg, and after finishing you leave it on the stove. At first, heat energy is concentrated near the skillet, which is, relative to the rest of the room, hot and quite nonrandom. But by the next morning the skillet has cooled to air temperature, and the heat energy has radiated throughout the room. That heat energy is now dispersed and unavailable for cooking; the system of the skillet, the room, and the heat has gone toward equilibrium, become more random, and entropy has increased. Unless an outside source of energy such as a stove, with fuel or electricity, is continually at work to maintain a non-equilibrium state, dispersion of heat results in a random equilibrium state. The same is true for all kinds of energy. According to this law, our solar system and presumably the entire universe should theoretically become a completely random overdispersed array of molecules and heat in the far distant future.