Gibbs Free energy vs Gibbs Free energy at standard state

In summary, ΔG° measures the change in Gibbs Free energy at 1 bar with stoichiometric amounts based on the chemical equation. This is different from ΔG, which is a more general term for Gibbs Free energy that may not have the same properties as ΔG°. For equilibrium reactions, ΔG=0 and the graph of the reaction would show a certain composition of reactants and products. In questions involving phase changes, ΔH° and ΔS° are given to find the temperature at which equilibrium occurs using the formula ΔG°=ΔH°-TΔS°. However, this may be confusing as ΔG° should not be used with 100% complete reactions.
  • #1
sgstudent
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ΔG° is the measure of Gibbs Free energy change at 1 bar but no specified temperature and also the stoichiometric amounts depending on the equation of the chemical reaction.

For example, if X ->2Y then the ΔG° would be equal to ΔG°f(2Y)-ΔG°f(X). While for ΔG it is a general term for Gibbs Free Energy that might not have the same properties as the ΔG° (1 bar, stoichiometric amounts)

So I was wondering, for a equilibrium reaction the ΔG=0 as when it reaches the lowest point the gradient=0. At this point, there is a certain composition of reactants to products that would not be 0% reactants and 100% products as that would mean the reaction is irreversible and reaches a lowest point at 100% products.

The graph for the reversible reaction would look like this: http://postimg.org/image/r9vtp0fh3/

So now in most questions, such as in phase changes questions they would give us the ΔH° and ΔS° and they then would want us to find the temperature at which equilibrium takes place. So using the formula ΔG°=ΔH°-TΔS° we would let ΔG° be 0 and solve for T. The part that i don't get is that ΔG° is a non-zero value and that we shouldn't be able to use ΔH° or ΔS° to find ΔG because either ΔH° or ΔS° represents 100% complete reaction. So I'm not sure why we are allowed to do those thing actually.

Thanks in advance for the help :)
 
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  • #2
For the case of a. phase change, the ##\Delta\ zeros## are not at 1 bar. So their definitions are totally different.
 

1. What is Gibbs Free energy?

Gibbs Free energy is a thermodynamic quantity that measures the amount of energy available in a system for performing useful work at a constant temperature and pressure.

2. What is the difference between Gibbs Free energy and Gibbs Free energy at standard state?

Gibbs Free energy at standard state refers to the Gibbs Free energy of a substance under standard conditions, which includes a temperature of 298 K (25°C), a pressure of 1 bar, and a concentration of 1 mol/L. The Gibbs Free energy of a substance under non-standard conditions may differ from its value at standard state due to changes in temperature, pressure, or concentration.

3. How is Gibbs Free energy related to spontaneity?

Gibbs Free energy is a measure of a system's spontaneity, with a negative value indicating that a reaction or process will occur spontaneously. A positive value indicates that the reaction or process will not occur spontaneously and requires the input of energy.

4. Can Gibbs Free energy be used to predict the direction of a reaction?

Yes, the sign of the Gibbs Free energy change (ΔG) can be used to predict the direction of a reaction. A negative ΔG indicates a spontaneous reaction in the forward direction, while a positive ΔG indicates a non-spontaneous reaction that will proceed in the reverse direction.

5. How can Gibbs Free energy be calculated?

Gibbs Free energy can be calculated using the equation ΔG = ΔH - TΔS, where ΔH is the change in enthalpy, ΔS is the change in entropy, and T is the temperature in Kelvin. Alternatively, Gibbs Free energy can also be calculated using tabulated standard Gibbs Free energy values for each reactant and product and applying the equation ΔG° = ΣnΔG°(products) - ΣmΔG°(reactants), where n and m are the stoichiometric coefficients for each reactant and product, respectively.

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