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Attraction between species during precipitate formation
QuiteThePredicament:
In this double replacement reaction AgCl(aq) + NaCl(aq) :rarrow: AgCl(s) + Na+ + NO3- the Na+ and NO3- ions are held apart by the ion-dipole interactions between Na-O and NO3-H. On the other hand, Ag+ and Cl- ions have low electronegativity difference, thus come together to form the covalent AgCl precipitate.
However, I understand that the ion-dipole forces pulling the Ag+ and Cl- apart isn't strong enough because of their low polarity, but how are they attracted to each other? If it was caused by the force of the negative and positive electric fields of the ions, wouldn't they also interact with the dipole moments of oxygen and hydrogen surrounding them, staying in solution like Na+ and NO3- ions? Or is there no force to move them to each other on their own spontaneously, instead they have to collide with their own momentum, caused by someone stirring the solution?
Enthalpy:
The logic looks fishy to me. Ionic salts can precipitate. Nor do I understand the "low polarity" of Ag+ and Cl-. One electron is the same charge on any ion. Where does this reasoning come from?
Besides that, I suppose you mean AgNO3 as a reactant.
Dissolution and precipitation are a matter of energy in the crystal and in the solution, as compared with the temperature. Or more cleanly, a matter of Gibbs energy. Considering only the energy in the solution is not enough.
And predictions based on qualitative properties or simple comparisons like the electronegativity are difficult. They use to fail in this case.
QuiteThePredicament:
Yes, AgNO3 is the reactant, mistyped it.
What you're saying is that, trying to predict solubility from how strong the ion-dipole interactions would be is wrong. What should be done, is to compare lattice and hydration energies and see if it's exothermic or endothermic. Only way to gather the energy values is by empirical evaluation, therefore trying to predict it is where the fault lies. Did I get it right?
Corribus:
Don't forget entropy.
Enthalpy:
--- Quote from: QuiteThePredicament on March 05, 2019, 06:43:40 PM ---What you're saying is that, trying to predict solubility from how strong the ion-dipole interactions would be is wrong. What should be done, is to compare lattice and hydration energies and see if it's exothermic or endothermic. Only way to gather the energy values is by empirical evaluation, therefore trying to predict it is where the fault lies. Did I get it right?
--- End quote ---
Yes, to the same reservation as Corribus pointed out, that G or ยต determine equilibria, not E or H. A simple reasoning about it is that if energy leaves somewhere, it arrives elsewhere, so minimizing the energy can't be a driving force. It's the distribution of energy that counts.
Yes too, solubility is what determines a precipitation. The mere existence of double displacement tells that the ease of hydration of lone ions is not the whole picture, since all ions were dissolved prior to the precipitation. How well two (or possibly more) ions match to build a solid counts too.
I know no qualitative nor simple quantitative way to determine a solubility. Bigger ions are easier to hydrolyse, for instance (NO3)- where the charge is diffuse (like a charge has a smaller energy on a bigger sphere), but for a crystal, the energy depends on arrangement details which I don't imagine to evaluate by hand.
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