The other factor that we have mentioned that speeds up reaction is temperature. Changing the temperature can do more than just accelerate a reaction; it can also affect the nature of the products. As an example, the synthesis of ammonia is important as a means of fixing atmospheric nitrogen for use in fertilizers and explosives:
N(g) + 3H(g) 2NH(g)
G = -7.95 kcal per 2 moles of NH
If the reaction is run at room temperature, the final mixture is almost entirely NH with very little N and H left. A disadvantage is that the reaction is extremely slow, but it can be speeded up with an iron-manganese catalyst. Trying to accomplish the same result by raising the temperature leads only to trouble, since at 450K, the product is no longer virtually pure NH, but is a mixture of N, H and NH in roughly equal proportions. The standard free energy change at this temperature is zero. (The standard starting condition of 1 atm partial pressure for each gas is, in fact, the equilibrium condition at 450K.) Even worse, at 1000K the standard free energy change is +29.6 kcal (compared with -7.95 kcal at 298K), and almost no ammonia is formed.

From these examples, we can make two observations:
1. Not all chemical reactions go to completion. Even after an infinitely long time, some systems remain mixtures of reactants and products.
2. Some reactions that are highly spontaneous by free energy criteria do not proceed at a measurable rate. Catalysis or heat sometimes can help.