The solubility of a given solute in a given solvent is function of temperature. Depending on the change in Gibbs free energy (ΔG) of the dissolution reaction, "i.e.", on the endothermic (ΔG > 0) or exothermic (ΔG < 0) character of the dissolution reaction, the solubility of a given compound may increase or decrease with temperature. The van 't Hoff equation relates the change of solubility equilibrium constant (Ksp) to temperature change and to reaction enthalpy change (ΔH). For most solids and liquids, their solubility increases with temperature because their dissolution reaction is endothermic (ΔG > 0). In liquid water at high temperatures, (e.g. that approaching the critical temperature), the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower dielectric constant results in a less polar solvent and in a change of hydration energy affecting the ΔG of the dissolution reaction.