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Consider the dissolution of NH4NO3 (commonly use in fertilizer) in a chemical cold pack. Is the process spontaneous or non spontaneous?Is the process endothermic or exothermic?
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Nicolae Palibroda
Consider the dissolution of NH4NO3 (commonly use in fertilizer) in a chemical cold pack. Is the process spontaneous or non spontaneous?Is the process endothermic or exothermic?
It is spontaneous and endothermic. Salts like ammonium nitrate have two factors driving the enthalpy of dissolution. The energy required to break the crystalline ionic bonds and the energy released in the hydration of the resulting ions. Some salts have ionic bonds that are just too strong for the water to break, and they do not dissolve. Or there is an equilibrium between the concentration of the dissolved salt and the crystal in a saturated solution.
It is spontaneous and endothermic. Salts like ammonium nitrate have two factors driving the enthalpy of dissolution. The energy required to break the crystalline ionic bonds and the energy released in the hydration of the resulting ions. Some salts have ionic bonds that are just too strong for the water to break, and they do not dissolve. Or there is an equilibrium between the concentration of the dissolved salt and the crystal in a saturated solution.
Well, would we not consider DISSOLUTION of a SALT in water an endothermic process? That is, a priori, dissolution of a salt is a bond-breaking chemical process, and thus should give rise to an endotherm. And the process is thus ENTROPY-DRIVEN rather than driven by enthalpy change…
Well, would we not consider DISSOLUTION of a SALT in water an endothermic process? That is, a priori, dissolution of a salt is a bond-breaking chemical process, and thus should give rise to an endotherm. And the process is thus ENTROPY-DRIVEN rather than driven by enthalpy change…
It is spontaneous and endothermic. Salts like ammonium nitrate have two factors driving the enthalpy of dissolution. The energy required to break the crystalline ionic bonds and the energy released in the hydration of the resulting ions. Some salts have ionic bonds that are just too strong for the water to break, and they do not dissolve. Or there is an equilibrium between the concentration of the dissolved salt and the crystal in a saturated solution.
It is spontaneous and endothermic. Salts like ammonium nitrate have two factors driving the enthalpy of dissolution. The energy required to break the crystalline ionic bonds and the energy released in the hydration of the resulting ions. Some salts have ionic bonds that are just too strong for the water to break, and they do not dissolve. Or there is an equilibrium between the concentration of the dissolved salt and the crystal in a saturated solution.
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Well, would we not consider DISSOLUTION of a SALT in water an endothermic process? That is, a priori, dissolution of a salt is a bond-breaking chemical process, and thus should give rise to an endotherm. And the process is thus ENTROPY-DRIVEN rather than driven by enthalpy change…
[math]NH_{4}^{+}NO_{3}^{-}(s)+Delta stackrel{H_{2}O(l)}longrightarrow NH_{4}^{+}+NO_{3}^{-}[/math]
Well, would we not consider DISSOLUTION of a SALT in water an endothermic process? That is, a priori, dissolution of a salt is a bond-breaking chemical process, and thus should give rise to an endotherm. And the process is thus ENTROPY-DRIVEN rather than driven by enthalpy change…
[math]NH_{4}^{+}NO_{3}^{-}(s)+Delta stackrel{H_{2}O(l)}longrightarrow NH_{4}^{+}+NO_{3}^{-}[/math]
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