Gibbs free energy

1800s, Josiah Willard Gibbs, (1839-1903) submitted scientific papers which mathematically combined both enthalpy and entropy (the measure of energy release and disorder in a system respectively) that also incorporates the second law of thermodynamics:

Entropy can never decrease, only increase for a reaction to take place.

in order to measure the amount of free enerfy present within any given system.

<blockquoteWhy doesn't free energy = enthalpy - entropy?

Reasoning behind Gibbs equation 
The signs for enthalpy and entropy must be opposites to each other, "because one function tends to a maximum and the other tends to a minimum." As a consequence the one equation proposed for this "unknown energy function" could be:

X= U - S (where X = Function, U = Enthalpy, S = Entropy)

Although this must be carried out under standard conditions, so U must be substituted for an H to demonstrate a constant pressure. (of 1atm) In addition to this, the units are wrong in our current equation; as we know, enthalpy is measure in joules (J) of energy, entropy on the other hand is measured in joules per kelvin. (J K^-1) Thus, we must also multiply entropy (S) by temperature in Kelvin. (T) Giving us the following equation when delta symbols are incorporate to represent that this function is for free energy changes:

dG = dH - dTS (where G = Free energy, H = Enthalpy, S = Entropy, T = Temperature, d = Change in associated function)

Application

We can now determine each individual component of this equation, enthalpy change being determined via "calorimetric measurment" to give us our value for dH; entropy can then be found if dG and temperature are known, the opposite can be said for determining dG itself with dS being our known value instead. For a reaction to be possible, it has been stated that the entropy of the universe is always increased. Consequently for a reaction to take place, dG must always be negative, with dS in the the equation for free energy exceeding that of the enthalpy change dH.