<Dd> G: free energy, H: enthalpy, T: temperature, S: entropy, Δ: difference (change between original and product) </Dd> <P> Reactions can be exothermic, where ΔH is negative and energy is released . Typical examples of exothermic reactions are 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 reaction products, which have high entropy . Since the entropy increases with temperature, many endothermic reactions preferably take place at high temperatures . On the contrary, many exothermic reactions such as crystallization occur at low temperatures . Changes in temperature can sometimes reverse the sign of the enthalpy of a reaction, as for the carbon monoxide reduction of molybdenum dioxide: </P> <Dl> <Dd> 2 CO (g) + MoO 2 (s) ⟶ 2 CO 2 (g) + Mo (s) (\ displaystyle (\ ce (2CO (g) + MoO2 (s) -> 2CO2 (g) + Mo (s)))); Δ H o = + 21.86 kJ at 298 K (\ displaystyle \ Delta H ^ (o) = + 21.86 \ (\ text (kJ at 298 K))) </Dd> </Dl> <Dd> 2 CO (g) + MoO 2 (s) ⟶ 2 CO 2 (g) + Mo (s) (\ displaystyle (\ ce (2CO (g) + MoO2 (s) -> 2CO2 (g) + Mo (s)))); Δ H o = + 21.86 kJ at 298 K (\ displaystyle \ Delta H ^ (o) = + 21.86 \ (\ text (kJ at 298 K))) </Dd>

Which process converts an atom from one element to another