The standard thermodynamics also says that the actual Gibbs free energy Δ G is related to the free energy change under standard state Δ G by the relationship: Δ Δ G = Δ Δ G ⊖ ⊖ + R T ln ⁡ ⁡ Q r where Q r is the reaction quotient and R is the universal ideal gas constant. The cell potential E associated with the electrochemical reaction is defined as the decrease in Gibbs free energy per coulomb of charge transferred, which leads to the relationship Δ Δ G = − − z F E . The constant F (the Faraday constant) is a unit conversion factor F = N A q, where N A is the Avogadro constant and q is the fundamental electron charge. This immediately leads to the Nernst equation, which for an electrochemical half-cell is E red = E red ⊖ ⊖ − − R T z F ln ⁡ ⁡ Q r = E red ⊖ ⊖ − − R T z F ln ⁡ ⁡ a Red a Ox . For a complete electrochemical reaction (full cell), the equation can be written as E cell = E cell ⊖ ⊖ − − R T z F ln ⁡ ⁡ Q r