Why does the dissolution of NH4NO3 happen if the ∆H is 28.1?

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The dissolution of NH4NO3 occurs despite a positive ∆H of 28.1 because the process is driven by an increase in entropy, making it spontaneous. The Gibbs free energy equation (∆G = ∆H - T∆S) illustrates that a favorable entropy change can offset the enthalpy increase. Strong bases can deprotonate H3PO4 after the third dissociation due to the unfavorable electronic environment, which requires a stronger base to facilitate the reaction. The strength of PO4^2- as a base is greater than that of H2PO4^-1, indicating its higher tendency to accept protons. Understanding these concepts is essential for analyzing acid-base reactions and their spontaneity.
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1. Why does the dissolution of NH4NO3 happen if the ∆H is 28.1? I'm supposed to be detailed with an equation, and explaining spontaneity.

Does it happen because of the energy created from protonation?

2. And why is only a strong base able to make H3PO4 act as an acid after the 3rd acid dissociation?

Is it because it is electronically unfavorable?
 
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1) at constant temperature and pressure what concept relates to sponaneity (depends on both entropy and enthalpy of process).

2)In terms of a base, how strong is PO4^2- compared to H2PO4^-1? What is the Arrehnius concept of an acid base neutralization in perspective of the conjugate acid/bases?
 
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