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### Course: MCAT>Unit 9

Lesson 3: Solubility equilibria

# Dissolution and precipitation

Defining solute, solvent, hydration, dissolution, precipitation, net ionic equation, and spectator ions. Looking at the molecular level interactions between water and ions in NaCl.  Created by Jay.

## Want to join the conversation?

• How do you know that the ionic crystal is AgCl(s). How come the solid is not NaNO3?
• solubility rules: AgCl will always be insoluble, but all alkali metal salts are soluble
• 1. Why does NaCl dissolve in water?. Isn't the electrostatic force to keep the ionic lattice greater than the attraction of the water molecules for the ions?
2. If all four compounds (NaCl, Ag(NO)3, AgCl and Na(NO)3) are ionic, why do NaCl, Ag(NO)3, Na(NO)3 all soluble in water but not AgCl?
• The exact reason is not simple and involves quantum mechanics. That said, many ionic solids, such as silver chloride (AgCl) do not dissolve in water because the forces holding the solid AgCl lattice together are too strong to be overcome by the forces favoring the formation of the hydrated ions, Ag+(aq) and Cl-(aq). Other salts and compounds that you mention will have different binding strength in a lattice, so some will dissolve, and others won't.
• why does he write 8- when he writes that Oxygen is partially negative and 8+ when he writes that Hydrogen is partially positive?
• It's really a lower-case Greek letter delta (δ). δ⁻ means "partially negative", and δ⁺ means "partially positive".
• What is a dipole at ?
• A dipole is a molecule in which a concentration of positive electric charge is separated from a concentration of negative charge.
Water is a dipole because it has a partially negative oxygen end, while the hydrogen atoms are at the positive end.
• When NaCl is dissolved in water, it dissociates into Na+ and Cl-. Does this mean, that Na and Cl are attracted more to water than to itself? Isn't ionic bond strong..? I understand that polar molecules dissolves in water, but the fact that ionic compound dissolves,, is not intuitive to me...
• Some ionic bonds are stronger than others - they're not all of equal strength. Thus the ionic bond between Ag+ and Cl- is stronger than between Na+ and Cl-.
• When NaCl bonds are broken in the solution, is any energy needed to do so? If yes,shouldn't the water/solution become colder?
• Nice point and for NaCl this will happen (although only by an extremely small amount, barely if at all noticeable)! If it takes more energy to separate the particles of the solute than is released when the water molecules bond to the particles, then the temperature goes down (endothermic).
• How come Ag+ and Cl- form AgCl but Na+ and NO3- do not form NaNO3?
Sodium and nitrate are also charged oppositely, shouldn't they attract?
• Some salts are soluble in water, and some salts aren't soluble in water. Added to which there are degrees of solubility, with some salts being more soluble than other salts.

It all depends on whether the ionic crystal state (AgCl or NaNO3) is more thermodynamically favoured than the solvated ions (Ag+ and Cl- or Na+ and NO3-). By solvated, I mean that water molecules surround the ions and stabilise them so that they prefer to be in solution rather than sitting as a crystals at the bottom of the beaker.

If ionic substances never dissolved in water then there would be no life as we know it.
• Is there any way to tell whether a compound will be solid or liquid based on molecular composition alone, or is some knowledge of the compound needed to determine its phase of matter during or after a reaction?