A spontaneous process is one that proceeds by itself. Such a process may not occur without the addition of a certain amount of energy, such as the combustion of a fuel. However, once it has begun, the reaction will proceed until the reactants are exhausted or the system reaches a state of equilibrium. A non-spontaneous process is the opposite of a spontaneous process. It will only proceed if energy is supplied to the system. However, the moment the energy supply is removed, the reaction will cease. The amount of useful work that can be obtained from a spontaneous reaction is limited by the difference in free energy between the products and reactants. If the free energy of a system is negative, the reaction at constant temperature and pressure can produce useful work, and is said to be spontaneous. If the free energy change of a system is positive, work must be performed to carry out the reaction at constant temperature and pressure; the reaction in this direction is non-spontaneous but the reverse reaction is spontaneous. If the free energy change of the system is zero, the reaction is in equilibrium, meaning that the forward reaction and the reverse reaction proceed at the same rate. Standard free energy changes are those that accompany changes from reactants in their standard states to products in their standard states. For solids and liquids, the standard state is the pure solid or liquid under standard pressure. For gases, the standard state is the ideal gas at standard pressure. Thermodynamic relations make it possible to obtain information about the reactants and products at equilibrium from the standard free energy change.
A reaction that has a positive free energy change at 25°C and atmospheric pressure might be made to proceed by raising the temperature or lowering the pressure. The dissociation of calcium carbonate to calcium oxide and carbon dioxide at 25°C and atmospheric pressure provides an example; raising the temperature or lowering the pressure drives the reaction in the forward direction.
The standard free energy of formation of a pure substance is defined as the free energy change when one mole of the substance is formed from its constituent elements in their standard states.
The free energy change is proportional to the amount of substance that reacts or is produced in a reaction. It is equal in magnitude but opposite in sign to the free energy change for the reverse reaction. If a reaction can be regarded as the sum of two or more reactions, the free energy change for the overall reaction must equal the sum of the free energy changes for the other reactions.
For a reaction occurring under constant pressure, the free energy change is equal to the enthalpy change of the system minus the absolute temperature times the entropy change of the system; if the reaction is spontaneous, this quantity, known as the Gibbs free energy, must be less than zero. For spontaneous reactions occurring at constant volume, the equivalent quantity is the Helmholtz free energy, defined as the internal energy change minus the absolute temperature times the entropy change.
Because, experimentally, it is easier to control the pressure than the volume during a reaction, the Gibbs free energy has more practical value than the Helmholtz free energy. By studying the variation of the Gibbs function with pressure at constant temperature, and with temperature under constant pressure, it is possible to gather useful volume and entropy data, which are indispensible to understanding the kinetics of reactions.
The standard free energy change is also related to the equilibrium constant of a reaction. Thus, the free energy change can be used to find the value of the equilibrium constant and vice versa.
Thermodynamics can be used to determine whether a change can take place spontaneously and how far a change may proceed before the reaction reaches equilibrium. It does not, however, provide any information about how fast a change will take place or about how changes take place.
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