# La Chatlier & Gibbs Free Energy Contradiction?

• Electric to be
In summary, the conversation discusses the relationship between temperature and the direction of a reaction, as well as the concepts of Le Chatlier's principle and Gibbs free energy. It is explained that while increasing temperature may cause reactions to shift in a certain direction according to Le Chatlier's principle, the change in Gibbs free energy is ultimately what determines the direction of the reaction. Furthermore, it is noted that the change in Gibbs free energy is dependent on the concentration of products and reactants, as well as temperature and pressure.

#### Electric to be

Say I have an exothermic reaction, whose change in Entropy is positive. (not the most common of reactions, but it can still happen)

If I increase the temperature, by La Chatlier's principle, the reaction should move to the left.

However, by Gibbs free energy, if I increase the temperature, the second term (dH - TdS) TdS, becomes more positive (and thus Gibb's free energy becomes more negative). So the reaction should move to the right?

So is La Chatlier's just a guiding tool for the majority of reactions, and not always correct? Like in this situation.

If you have a system that can move to a state of of higher entropy and still give off energy (exothermic), it is likely that it is not in equilibrium, and thereby Le Chatlier's principle would not apply. It's already in a reactive state. Adding heat should speed up the process.

If you have a system that can move to a state of of higher entropy and still give off energy (exothermic), it is likely that it is not in equilibrium, and thereby Le Chatlier's principle would not apply. It's already in a reactive state. Adding heat should speed up the process.

Okay so basically Gibbs free energy always applies and in this case adding heat would speed up the reaction.

However, for a system in equilibrium increasing the temperature will shift the equilibrium constant so that the endothermic reaction occurs. This makes sense to me if you kind of think of heat as one of the reactants and products, but how is this actually derived? Is it purely experimental?

Electric to be said:
However, for a system in equilibrium increasing the temperature will shift the equilibrium constant so that the endothermic reaction occurs. This makes sense to me if you kind of think of heat as one of the reactants and products, but how is this actually derived? Is it purely experimental?

This fact can be derived from the equation that relates the equilibrium constant to the change in Gibbs free energy. How does increasing the temperature alter the ΔG of an endothermic reaction? How does increasing temperature alter the ΔG of an exothermic reaction?

Ygggdrasil said:
This fact can be derived from the equation that relates the equilibrium constant to the change in Gibbs free energy. How does increasing the temperature alter the ΔG of an endothermic reaction? How does increasing temperature alter the ΔG of an exothermic reaction?

Well that depends on whether the Entropy change is positive or negative right? delta G = delta H - T deltaS

That was the original source of my confusion

Ygggdrasil said:
...

If I could ask one last thing (just quoting to notify), it's about Gibbs free energy. By the equation Delta G = DeltaG(standard) + RTln(Q) where Q is the reaction quotient, DeltaG is clearly dependent on the concentration of products and reactants.

However, the basic definition of it is DeltaG = DeltaH -TDeltaS. But isn't it true that DeltaH and DeltaS only depend the amount of product that was formed, and not the concentration of the reactants and products during formation. For example, if 1 mol of some substance was formed while the concentration of the reactants was 3x larger than if a mol of the same substance was formed with less concentration, the enthalpy change and heat released should be the same.

So under this definition of DeltaG, how does DeltaG change depending on the concentrations if the same amount of product is formed?

Electric to be said:
If I could ask one last thing (just quoting to notify), it's about Gibbs free energy. By the equation Delta G = DeltaG(standard) + RTln(Q) where Q is the reaction quotient, DeltaG is clearly dependent on the concentration of products and reactants.

However, the basic definition of it is DeltaG = DeltaH -TDeltaS. But isn't it true that DeltaH and DeltaS only depend the amount of product that was formed, and not the concentration of the reactants and products during formation. For example, if 1 mol of some substance was formed while the concentration of the reactants was 3x larger than if a mol of the same substance was formed with less concentration, the enthalpy change and heat released should be the same.

So under this definition of DeltaG, how does DeltaG change depending on the concentrations if the same amount of product is formed?
I am no expert on thermodynamics, but ## \Delta S ## of the reaction is ## S_{final}-S_{initial} ## and both of these quantities will depend on the concentrations of each and all of the substances that take part in the reaction=the initial concentrations and the final ones. I believe entropy is a quantity that is very much concentration dependent.

In the equation ##\Delta G=\Delta G^0+RT\ln Q##, ##\Delta G## refers to the change in free energy in going from stoichiometric molar quantities of pure reactants at temperature T and prescribed pressures (not necessarily 1 bar) to stoichiometric molar quantities of pure products at temperature T and prescribed pressures (not necessarily 1 bar).

Pepper Mint
Chestermiller said:
In the equation ##\Delta G=\Delta G^0+RT\ln Q##, ##\Delta G## refers to the change in free energy in going from stoichiometric molar quantities of pure reactants at temperature T and prescribed pressures (not necessarily 1 bar) to stoichiometric molar quantities of pure products at temperature T and prescribed pressures (not necessarily 1 bar).

Yes, I understand this. However my question is, depending on the current concentration of products and reactants, does either deltaH or deltaS change? I wouldn't think deltaH would, possibly deltaS would. Obviously at least one of them would have to change, since deltaG changes as Q does.

Electric to be said:
Yes, I understand this. However my question is, depending on the current concentration of products and reactants, does either deltaH or deltaS change? I wouldn't think deltaH would, possibly deltaS would. Obviously at least one of them would have to change, since deltaG changes as Q does.
I am very puzzled by this question. If the temperature, pressure, and current concentration (whatever the current concentration means) are constant, how can H, S, and G change? Are you talking about a liquid phase reaction or a gas phase reaction? Are you talking about a van't Hoff equilibrium box, in which pure reactants are injected and pure products are removed?

Electric to be said:
If I could ask one last thing (just quoting to notify), it's about Gibbs free energy. By the equation Delta G = DeltaG(standard) + RTln(Q) where Q is the reaction quotient, DeltaG is clearly dependent on the concentration of products and reactants.

However, the basic definition of it is DeltaG = DeltaH -TDeltaS. But isn't it true that DeltaH and DeltaS only depend the amount of product that was formed, and not the concentration of the reactants and products during formation. For example, if 1 mol of some substance was formed while the concentration of the reactants was 3x larger than if a mol of the same substance was formed with less concentration, the enthalpy change and heat released should be the same.

So under this definition of DeltaG, how does DeltaG change depending on the concentrations if the same amount of product is formed?

I believe the correct equation here should be ΔG° = ΔH° - TΔS°.

Chestermiller said:
I am very puzzled by this question. If the temperature, pressure, and current concentration (whatever the current concentration means) are constant, how can H, S, and G change? Are you talking about a liquid phase reaction or a gas phase reaction? Are you talking about a van't Hoff equilibrium box, in which pure reactants are injected and pure products are removed?

Yes I am talking about a van't Hoff equilibrium box. What I'm considering are two separate equilibrium boxes. G will remain constant within the boxes, but be different between the two boxes. Both at the same temperature and pressure, but one box has a different ratio of reactants and products to the other box.

I know that G must be different, since the reactant quotient is different between the two boxes. But I'm wondering, it doesn't really make sense to me that DeltaH(for a unit amount of reactions) would have changed much if at all. Shouldn't some pair of molecules reacting to form a new molecule release the same amount of energy, regardless of the conditions the molecules are constrained to? (Temperature, pressure, and especially relative concentration). I'm willing to accept deltaS (for a unit amount of reaction times) might change depending on concentration, but I'm wondering why?

Electric to be said:
Yes I am talking about a van't Hoff equilibrium box. What I'm considering are two separate equilibrium boxes. G will remain constant within the boxes, but be different between the two boxes. Both at the same temperature and pressure, but one box has a different ratio of reactants and products to the other box.
If the contents of the two boxes are each at equilibrium, then the ratio of reactants and products in the two boxes is the same, and equal to the equilibrium constant.

But I'm wondering, it doesn't really make sense to me that DeltaH(for a unit amount of reactions) would have changed much if at all. Shouldn't some pair of molecules reacting to form a new molecule release the same amount of energy, regardless of the conditions the molecules are constrained to? (Temperature, pressure, and especially relative concentration). I'm willing to accept deltaS (for a unit amount of reaction times) might change depending on concentration, but I'm wondering why?
When you are talking about ##\Delta H##, ##\Delta S##, and ##\Delta G##, you are referring to the change in these functions between two thermodynamic equilibrium states. It is important in such cases that you precisely define the initial and final thermodynamic equilibrium states. So please, now define for us the two thermodynamic equilibrium states that you are referring to here (so that we can remove all ambiguity).

State 1: ...

State 2:...

Chet

## 1. What is the La Chatelier principle?

The La Chatelier principle, also known as the equilibrium principle, states that when a system in equilibrium is subjected to a disturbance, it will adjust in such a way as to minimize the effect of the disturbance and return to its original equilibrium state.

## 2. What is the Gibbs free energy contradiction?

The Gibbs free energy contradiction is a theoretical contradiction between the La Chatelier principle and the second law of thermodynamics. According to the La Chatelier principle, a system in equilibrium will respond to a disturbance by counteracting it and returning to equilibrium. However, the second law of thermodynamics states that all spontaneous processes must increase the overall entropy of the universe. This apparent contradiction arises when a system is subjected to a disturbance that increases its Gibbs free energy, which goes against the principle of equilibrium.

## 3. How does the Gibbs free energy relate to the La Chatelier principle?

The Gibbs free energy is a thermodynamic potential that measures the maximum amount of work that can be extracted from a system at constant temperature and pressure. According to the La Chatelier principle, a system will respond to a disturbance in a way that minimizes its Gibbs free energy. This means that if a disturbance causes an increase in the Gibbs free energy, the system will respond in a way that decreases its Gibbs free energy back to its original value, returning to equilibrium.

## 4. Can the Gibbs free energy contradiction be observed in real-world systems?

The Gibbs free energy contradiction is a theoretical concept and cannot be directly observed in real-world systems. However, it can be observed indirectly by studying systems that are in a state of dynamic equilibrium, where the rates of forward and reverse reactions are equal. In these systems, a disturbance can cause a temporary increase in Gibbs free energy, but the system will respond by shifting to a new equilibrium state with a lower Gibbs free energy.

## 5. How is the Gibbs free energy contradiction resolved?

The Gibbs free energy contradiction is resolved by considering the system and its surroundings together as a closed system. When a disturbance causes an increase in Gibbs free energy, the system will respond by decreasing its Gibbs free energy, but this will result in an increase in the Gibbs free energy of the surroundings. This overall decrease in Gibbs free energy of the closed system satisfies the second law of thermodynamics and resolves the apparent contradiction with the La Chatelier principle.