Main Page | See live article | Alphabetical index A spontaneous process in chemical reaction terms is one which occurs with the system releasing free energy in some form (often, but not always, heat) and moving to a lower energy, hence more thermodynamically stable, state. The sign convention is the same as for exothermic/endothermic reactions, with a release of free energy corresponding to a negative change in G (the Gibbs free energy). Thus, for a reaction at constant temperature where ΔG = ΔH -TΔS a negative ΔG would require a negative change in enthalpy (ie an exothermic reaction) and/or a positive change in the entropy. This is a direct consequence of the second law of thermodynamics, which states that for 'any' spontaneous process the overall change in entropy of the system must be greater than or equal to zero. Can a spontaneous chemical reaction have a product with a lower entropy than the reactants? Yes it can. This does not contradict the second law however, since such a reaction must have a sufficiently large negative change in enthalpy (heat energy) that the increase in temperature of the reaction surroundings (considered to be part of the system in thermodynamic terms) results in a sufficiently large increase in entropy that overall the change in entropy is positive. This is why so many spontaneous reactions are exothermic. The release of heat energy creates a large increase in entropy. A spontaneous endothermic reaction (one which absorbs heat) must have a sufficiently large increase in the entropy of the products that it offsets the decrease in the entropy of the surroundings. Many important reactions are not spontaneous, and require the external addition of energy to drive them forward eg aluminium smelting, iron smelting. This article is a stub. You can help Wikipedia by fixing it. This article is from Wikipedia. All text is available under the terms of the GNU Free Documentation License.