SUMMARY
The dissolution of ammonium nitrate (NH4NO3) occurs despite a positive enthalpy change (∆H = 28.1 kJ/mol) due to the favorable increase in entropy during the process. The spontaneity of the dissolution can be explained using the Gibbs free energy equation, where the change in free energy (∆G) is negative when the increase in entropy (∆S) outweighs the positive enthalpy change. Additionally, the behavior of phosphoric acid (H3PO4) as an acid after its third dissociation is influenced by the strength of the base involved, with strong bases effectively facilitating this process due to their ability to stabilize the resulting conjugate base.
PREREQUISITES
- Understanding of Gibbs free energy and its relation to spontaneity
- Knowledge of enthalpy (∆H) and entropy (∆S) concepts
- Familiarity with acid-base theories, particularly the Arrhenius concept
- Basic chemistry of ammonium nitrate and phosphoric acid
NEXT STEPS
- Study the Gibbs free energy equation and its applications in chemical reactions
- Research the relationship between entropy and spontaneity in thermodynamics
- Explore the properties of strong bases and their role in acid-base reactions
- Investigate the dissociation constants of phosphoric acid and its conjugate bases
USEFUL FOR
Chemistry students, educators, and professionals interested in thermodynamics, acid-base chemistry, and the behavior of ionic compounds in solution.