..... Click the link for more information. and used as a measure of the relative stability of a physical or chemical system, i.e., the tendency of the system to react or change. If the change in free energy, ΔG, is negative, the transformation of the system will occur spontaneously, since transitions in which the energy decreases are favored, whereas those in which it increases (ΔG positive) are not. The change in free energy for a given process at a particular temperature depends on three factors, as seen from the equation ΔG = ΔH − TΔS, where ΔH is the change in the enthalpy enthalpy (ĕn`thălpē)
..... Click the link for more information. of the system, T is the temperature in degrees Kelvin, and ΔS is the change in entropy entropy (ĕn`trəpē), quantity specifying the amount of disorder or randomness in a system bearing energy or information.
..... Click the link for more information. . A negative value of the enthalpy change indicates a decrease in the heat content of the system and contributes to a favorable value of the free energy; a positive entropy change indicates a decrease in the orderliness of the system and also contributes to a favorable value of the free energy, since a system tends to go from more ordered to less ordered states. It may happen that the change in enthalpy for the reaction is favorable but that of the entropy is unfavorable, or vice versa; in such a case the temperature is the deciding factor since it determines how much weight is given to the entropy change. For example, in the transition of liquid water to ice, the enthalpy change is favorable because heat is released in the process but the entropy change is unfavorable because the transition is to the more ordered, crystalline state. Below a temperature of 32°F; (273°K;) the enthalpy term, ΔH, is larger and the process is spontaneous, but at higher temperatures the entropy term, TΔS, predominates, and the transition does not occur. Although the free energy indicates whether or not a given reaction will occur, it gives no information about the speed of such a reaction. The reaction of hydrogen with oxygen to form water has a favorable, negative, free energy, but the reaction rate is so slow that without the presence of a catalyst it is not observable. Scientists use tables listing the standard free energy, ΔG°, of various compounds; the standard free energy is the change in free energy when one mole of the compound is formed at 25°C; and 1 atmospheric pressure.
free energy
Measure of the total combined energies within a system, derived from heats of transformation, disorder, and other forms of internal energy (e.g., electrostatic charges). A system will change spontaneously to achieve a lower total free energy. Thus, free energy is the driving force toward equilibrium conditions. The change in free energy between an initial and a final state is useful in evaluating certain thermodynamic processes and can be used to judge whether transformations will occur spontaneously. There are two forms of free energy, with different definitions and applications: the Helmholtz (see Hermann von Helmholtz) free energy, sometimes called the work function, and the Gibbs (see J. Willard Gibbs) free energy.
free energy [¦frē ′en·ər·jē]
Free energy
A term in thermodynamics which in different treatments may designate either of two functions defined in terms of the internal energy E or enthalpy H, and the temperature-entropy product TS.
The function (E – TS) is the Helmholtz free energy and is the function ordinarily meant by free energy in European references. The Gibbs free energy is the function (H – TS). For the Lewis and Randall school of American chemical thermodynamics, this is the function meant by the free energy F. To avoid confusion with the symbol F as applied elsewhere to the Helmholtz free energy, the symbol G has also been used. Another development was the introduction of the name free enthalpy, with symbol G, for the Gibbs function. See Work function (thermodynamics)
For a closed system (no transfer of matter across its boundaries), the work which can be done
For constant temperature and pressure in a reversible process the decrease in the Gibbs function G for the system again corresponds to a free-energy change in the above sense, since it is equal to the work which can be done by the closed system other than that associated with its change in volume ΔV under the given constant pressure P. The relations shown in Eq. (2) can be formed since ΔH = ΔE + PΔV.
Each of these free-energy functions is an extensive property of the state of the thermodynamic system. For a specified change in state, both ΔA and ΔG are independent of the path by which the change is accomplished. Only changes in these functions can be measured, not values for a single state.
The thermodynamic criteria for reversibility, irreversibility, and equilibrium for processes in closed systems at constant temperature and pressure are expressed naturally in terms of the function G. For any infinitesimal process at constant temperature and pressure, -dG ≥ δwnet. If δwnet is never negative, that is, if the surroundings do no net work on the system, then the change dG must be negative or zero. For a reversible differential process, -dG > δwnet; for an irreversible process, -dG > δwnet. The free energy G thus decreases to a minimum value characteristic of the equilibrium state at the given temperature and pressure. At equilibrium, dG = 0 for any differential process taking place, for example, an infinitesimal change in the degree of completion of a chemical reaction. A parallel role is played by the work function A for conditions of constant temperature and volume. Because temperature and pressure constitute more convenient working variables than temperature and volume, it is the Gibbs free energy which is the more commonly used in thermodynamics. See Entropy, Thermodynamic principles
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