# Difference between thermodynamic and kinetic stability

What is the difference between thermodynamic and kinetic stability? I'd like a basic explanation, but not too simple. For example, methane does not burn until lit -- why?

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What are your initial impressions about it? It will be helpful for those answering to be able to gauge your present understanding. –  jonsca Jan 6 '13 at 0:48

There is often a "reaction barrier" that needs to be overcome. Thus one must supply energy to get the reaction started, but once started it becomes self-sustaining.

One can think of it this way: one has to break bonds in a reaction before one can make bonds. Energy is needed to break the first bonds. Then making new bonds can supply enough energy to keep the reaction going.

Note that I'm talking about thermodynamically allowed reactions. If a reaction is not thermodynamically allowed, then even if some bonds are broken by adding energy, the reaction will not spontaneously continue.

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So you mean activation energy for example is kinetic stability. However I still do not understand Thermodynamic stability, does it have anything to do with entropy as I know which reactions are spontaneous by working out their Delta S Total. –  Cyrus Jan 6 '13 at 7:58
@Cyrus: Yes, activation energy has to do with kinetic stability. For example, diamonds are thermodynamically unstable relative to graphite, but don't hold your breath waiting for the spontaneous conversion. The second half is more complex. The thermodynamic quantity predicting the direction of a reaction depends on the conditions of the process being considered. For example at constant $T$ and $p$, the Gibbs Free energy $\Delta G$, must be negative. At constant $E$ and $V$, the entropy change $\Delta S$ must be positive. –  Paul J. Gans Jan 7 '13 at 2:15

Kinetics deals with all things that impact the rate at which a reaction occurs. One implication of that is that reaction kinetics are partially a function of the stability of the intermediate species (and transition states) that the pathway of a chemical reaction necessarily has to pass through en route to the formation of the final product. If those intermediates are highly unstable (i.e., highly energetic) relative to the initial reactants, then the initial barrier of activation energy that must be surpassed to set the reaction in motion will be relatively high, and the reaction will tend to proceed slowly (if at all).

Thermodynamic stability, on the other hand, is strictly a function of the change in free energy (ΔG), which is a state function, meaning that its value is determined exclusively by the difference between the initial state and the final state (or the free energy of the initial products and the final products). In other words, thermodynamic stability is totally independent of the pathway between reactants and products.

While it's true that thermodynamic stability is the ultimate arbiter of the spontaneity of a reaction, if the kinetics render the reaction too slow then, in practice, the reaction may not actually occur at all (under some given set of conditions). The classic textbook example is the conversion of diamond to graphite, which is thermodynamically favorable because the free energy of graphite is lower, but doesn't occur under ordinary conditions because the kinetics of the reaction (in the form of the immense activation energy required) are extremely unfavorable.

Imagine you have some set of reactants, A + B, and there are two different pathways through which A + B can react, both of which lead to final products that are lower in free energy than the reactants A + B. Pathway #1 has much higher activation energy than Pathway #2, but the final products at the end of Pathway #1 are lower in free energy (i.e., more thermodynamically stable) than those at the end of Pathway #2. This means that Pathway #1 is thermodynamically favored, but Pathway #2 is kinetically favored (that is, it poses a lower barrier in terms of activation energy and the reaction will proceed more rapidly via this pathway). The consequence of this is that the kinetic pathway, Pathway #2, will be the favored one, unless environmental conditions (temperature, pressure, concentrations of catalysts, etc.) are tuned in a such a way that the higher activation energy of Pathway #1 can be surpassed.

In summary, thermodynamic stability (in terms of free energy difference between reactants and products) will determine whether a given reaction could be theoretically spontaneous, but kinetic factors will decide whether the reaction occurs in practice under the given environmental conditions, as well as which pathway for a given reaction takes primacy.

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