A spontaneous process is the time-evolution of a system in which it releases free energy (usually as heat) and moves to a lower, more thermodynamically stable energy state. The sign convention of changes in free energy follows the general convention for thermodynamic measurements, in which a release of free energy from the system corresponds to a negative change in free energy, but a positive change for the surroundings.
Depending on the nature of the process, the free energy is determined differently. For example, the Gibbs free energy is used when considering processes that occur under constant pressure and temperature conditions whereas the Helmholtz free energy is used when considering processes that occur under constant volume and temperature conditions.
Because spontaneous processes are characterized by a decrease in the system's free energy, they do not need to be driven by an outside source of energy.
For a process that occurs at constant temperature and pressure, spontaneity can be determined using the change in Gibbs free energy, which is given by:
where the sign of ΔG depends on the signs of the changes in enthalpy (ΔH) and entropy (ΔS), as well as on the absolute temperature (T). The sign of ΔG will change from positive to negative (or vice versa) where T = ΔH/ΔS.
In cases where ΔG is:
- negative, the process is spontaneous and may proceed in the forward direction as written.
- positive, the process is non-spontaneous as written, but it may proceed spontaneously in the reverse direction.
- zero, the process is at equilibrium, with no net change taking place over time.
This set of rules can be used to determine four distinct cases by examining the signs of the ΔS and ΔH.
- When ΔS > 0 and ΔH < 0, the process is always spontaneous as written.
- When ΔS < 0 and ΔH > 0, the process is never spontaneous, but the reverse process is always spontaneous.
- When ΔS > 0 and ΔH > 0, the process will be spontaneous at high temperatures and non-spontaneous at low temperatures.
- When ΔS < 0 and ΔH < 0, the process will be spontaneous at low temperatures and non-spontaneous at high temperatures.
For the latter two cases, the temperature at which the spontaneity changes will be determined by the the relative magnitudes of ΔS and ΔH.
The second law of thermodynamics states that for any spontaneous process the overall ΔS must be greater than or equal to zero, yet a spontaneous chemical reaction can result in a negative change in entropy. 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. That is, the ΔS of the surroundings increases enough because of the exothermicity of the reaction that it overcompensates for the negative ΔS of the system, and since the overall ΔS = ΔSsurroundings + ΔSsystem, the overall change in entropy is still positive.
Another way to view the fact that some spontaneous chemical reactions can lead to products with lower entropy is to realize that the second law states that entropy of an isolated system must increase (or remain constant). Since a negative enthalpy change in a reaction means that energy is being released to the surroundings, then the 'isolated' system includes the chemical reaction plus its surroundings. This means that the heat release of the chemical reaction sufficiently increases the entropy of the surroundings such that the overall entropy of the isolated system increases in accordance with the second law of thermodynamics.for a irreversible process change in entropy always >0
Spontaneity does not imply that the reaction proceeds with great speed. For example, the decay of diamonds into graphite is a spontaneous process that occurs very slowly, taking millions of years. The rate of a reaction is independent of its spontaneity, and instead depends on the chemical kinetics of the reaction. Every reactant in a spontaneous process has a tendency to form the corresponding product. This tendency is related to stability. Stability is gained by a substance if it is in a minimum energy state or is in maximum randomness. Only one of these can be applied at a time. e.g. Water converting to ice is a spontaneous process because ice is more stable since it is of lower energy. However, the formation of water is also a spontaneous process as water is the more random state.
- Endergonic reaction reactions which are not spontaneous at standard temperature, pressure, and concentrations.
- Diffusion spontaneous phenomenon that minimizes Gibbs free energy.