If delta G for the dissociation of Borax at 30 degrees C (higher than room temperature) is positive, why does it dissolve at room temperature?
It gets the needed energy from somewhere. Do you suppose the solution gets colder?
Here is a site that may be worth reading. Especially near the bottom where the discussion is on spontaneous versus non-spontaneous reactions.
http://www.800mainstreet.com/7/0007-003-free_energy.htm
The point is that non-spontaneous reactions (those with delta G = +) can take place if energy is added from elsewhere. My question was to suggest that the energy comes from the solvent. That is one reason why some solids dissolve and the solution gets cold; i.e., the energy to break the bonds comes from the solvent molecules and they get colder in the process. Of course, some molecule produce a lot of extra energy when the ions become hydrated; however, this is not always enough to break the bonds of the crystal and the energy comes from the solvent decreasing T.
To understand why Borax dissolves at room temperature even if the change in Gibbs free energy for its dissociation at a higher temperature is positive, we need to consider a few concepts.
Firstly, the change in Gibbs free energy (∆G) is an important thermodynamic parameter that determines the spontaneity of a reaction. If ∆G is negative, the reaction is spontaneous and likely to occur without any external input of energy. Conversely, if ∆G is positive, the reaction is non-spontaneous and generally requires an input of energy.
However, it's important to note that the value of ∆G also depends on other factors, such as temperature and concentration. The relationship between ∆G, concentration, and temperature is described by the equation:
∆G = ∆G° + RT ln(Q)
where ∆G° is the standard Gibbs free energy change, R is the gas constant, T is the temperature in Kelvin, and Q is the reaction quotient.
In the case of Borax dissolution at room temperature, even if ∆G for its dissociation at a higher temperature is positive, it is possible for Borax to dissolve. This is because at room temperature, the value of ∆G for the reaction at equilibrium will be negative.
The dissolution of Borax is an example of an endothermic process. When Borax dissolves, it absorbs heat from the surroundings, which increases the randomness or entropy of the system. This increase in entropy (∆S) has a significant impact on the overall value of ∆G.
At room temperature (∼25 degrees Celsius), the value of ∆S will generally contribute more to the overall value of ∆G than the endothermic ∆H term. As a result, even if the dissociation of Borax at a higher temperature (∼30 degrees Celsius) has a positive ∆G, the increased entropy at room temperature can drive the dissolution process and make it energetically favorable (∆G < 0).
In summary, although the change in Gibbs free energy for the dissociation of Borax might be positive at a higher temperature, the favorable increase in entropy at room temperature can drive the dissolution process and make it occur spontaneously.