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[daddythicclegs]

Could someone explain what ion-dipoles are and the relevance they have in dissociation and Ksp

 

[wizzkids]

I will try to explain this in the context of Year 12 Chem, and Module 5 solubility equilibrium. Solubility and Ksp increases in the presence of attractive forces between ions and polar solvent molecules. Ion-dipole attraction explains for example, the enhanced solvation of ions by a polar solvent such as water. Compare the solubility of NaCl in ethanol (0.65 g per kg of solvent) with the solubility of NaCl in water (350 g per kg of solvent). You can see the solvent effect. You can also have ion-dipole attraction within coordination complexes. These are typically a central positively-charged metal ion surrounded by an array of bound molecules or other ions. The bonds within these coordination complexes can be weak or strong, and this influences the stability of the complex, and in turn this can influence the Ksp of a sparingly soluble solid which is in equilibrium with the complex.
Does that make sense?

 

[CM_Tutor]

I will try to explain this in the context of Year 12 Chem, and Module 5 solubility equilibrium. Solubility and Ksp increases in the presence of attractive forces between ions and polar solvent molecules. Ion-dipole attraction explains for example, the enhanced solvation of ions by a polar solvent such as water. Compare the solubility of NaCl in ethanol (0.65 g per kg of solvent) with the solubility of NaCl in water (350 g per kg of solvent). You can see the solvent effect. You can also have ion-dipole attraction within coordination complexes. These are typically a central positively-charged metal ion surrounded by an array of bound molecules or other ions. The bonds within these coordination complexes can be weak or strong, and this influences the stability of the complex, and in turn this can influence the Ksp of a sparingly soluble solid which is in equilibrium with the complex.
Does that make sense?

I recommend avoiding the phrase "K_sp increases" or anything similar for any type of equilibrium constant because it has a high risk of leading a marker to think that you think equilibrium constants change value, which can only occur with changes in temperature. I appreciate that you mean K_sp values for a chosen compound will be different in different solvents, and that is true because they are different equilibrium systems. However, I have seen a lot of answers lose marks because they treat equilibrium constants as non-constant.

Ion-dipole interactions are an intermolecular / between molecules and ions force, similar to a dipole-dipole interaction except that one of the dipoles (an uneven distribution of charge in a polar but net-uncharged molecule) has been replaced by a monopole, an ion with an overall charge. They are electrostatic in nature but manifest with an ion surrounded by polar molecules, all with their dipole vectors converging on the ion.

I also strongly recommend that you do not describe coordination complexes as examples of ion-dipole interactions as they are related but different phenomena. Take the complex of copper(II) with water, for instance. A coordination complex [Cu(H₂O)₆]²⁺ forms, with the six oxygen atoms of the water molecules being located at the corners of an octahedron with the copper(II) ion at the centre. The bonding is largely covalent in nature and the geometry is octahedral. Surrounding this overall positively charged ion will be another layer / shell of water atoms interacting with the coordination complex through ion-dipole interactions. The geometry of the positions of these water molecules is not fixed, nor is their number, but they will be aligned with their oxygen atoms able to hydrogen bond with the bound water molecules in the coordination complex and with their dipoles pointing to the centre of the complex. We are often careless in using terminology for describing this hydration system and there are ions (acetate, for example) where there is only the loose hydration sphere of water molecules with no coordination complex at the centre (though even here the oxygen atoms of the acetate ion will provide hydrogen bond acceptors to the solvent).

As far as complexes altering solubility, this is better recognised as a series of interconnected equilibria. Silver chloride has low solubility as

AgCl (s) <----in equilibrium----> Ag⁺ (aq) + Cl^-​

lies far to the left, as confirmed by its small K_sp. In the presence of ammonia, a second (complexation) equilibrium is established.

Ag⁺ (aq) + 2NH₃ (aq) <----in equilibrium----> [Ag(NH₃)₂]⁺ (aq)​

This equilibrium lies to the right, as confirmed by its comparatively large K_stab, and so causes the concentration of silver(I) ions in the solution to fall. Applying Le Chatelier's Principle, this leads to a shift to the right of the first equilibrium and increased solubility. It does not, however, cause the K_sp of silver(I) chloride to change.

You can combine this to describe the equilibrium as

AgCl (s) + 2NH₃ (aq) <----in equilibrium----> [Ag(NH₃)₂]⁺ (aq) + Cl^- (aq)​

which has its own equilibrium constant

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