section_path: Thermodynamics
section_name: Thermodynamics
paragraph_text: Thermodynamics
  
  Chemical reactions are determined by the laws of thermodynamics. Reactions can proceed by themselves if they are exergonic, that is if they release free energy. The associated free energy change of the reaction is composed of the changes of two different thermodynamic quantities, enthalpy and entropy:
  
  Reactions can be exothermic, where Δ H is negative and energy is released. Typical examples of exothermic reactions are combustion, precipitation and crystallization, in which ordered solids are formed from disordered gaseous or liquid phases. In contrast, in endothermic reactions, heat is consumed from the environment. This can occur by increasing the entropy of the system, often through the formation of gaseous or dissolved reaction products, which have higher entropy. Since the entropy term in the free-energy change increases with temperature, many endothermic reactions preferably take place at high temperatures. On the contrary, many exothermic reactions such as crystallization occur preferably at lower temperatures. A change in temperature can sometimes reverse the sign of the enthalpy of a reaction, as for the carbon monoxide reduction of molybdenum dioxide:
  
  This reaction to form carbon dioxide and molybdenum is endothermic at low temperatures, becoming less so with increasing temperature. Δ H ° is zero at 1855 K, and the reaction becomes exothermic above that temperature.
  
  Changes in temperature can also reverse the direction tendency of a reaction. For example, the water gas shift reaction
  
  is favored by low temperatures, but its reverse is favored by high temperatures. The shift in reaction direction tendency occurs at 1100 K.
  
  Reactions can also be characterized by their internal energy change, which takes into account changes in the entropy, volume and chemical potentials. The latter depends, among other things, on the activities of the involved substances.