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Topic: NaCl and AgCl solutivity question  (Read 31622 times)

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Offline Perrin

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NaCl and AgCl solutivity question
« on: August 09, 2008, 08:59:26 PM »
Hi, I've been learning chemistry alone from 0 about 2 weeks now.
While learning, my book said that NaCl will make a solution with water, and the ionic bond will break, while AgCl will not make a good solution. Now, NaCl has a strong ionic bond, and AgCl has a weaker ionic bond, based on the radius of Na being much shorter than the radius of AgCl. My question is then, why will the NaCl bonds break, being strong, while the AgCl will not? Is there some way to predict the result of these kinds of problems? And what about ionic molecules and non-water solvents? Like, for example, KBr in a C8H18(l) solution?
This subject has me a bit confused.
Thanks. Perrin.

Offline Ak

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Re: NaCl and AgCl solutivity question
« Reply #1 on: August 10, 2008, 09:56:31 AM »
http://en.wikipedia.org/wiki/Solubility#Solubility_of_ionic_compounds_in_water
-wiki tells you how you can calculate solubility and it gives a small chart of rules
http://www.ausetute.com.au/solrules.html
- this site gives you a larger chart of rules

http://www.elmhurst.edu/~chm/vchembook/171solublesalts.html
-From this site: "When an ionic crystal such as NaCl is placed in water, a dissolving reaction will occur. Initially, the positive and negative ion are only attracted to each other. The water molecules are hydrogen bonded to each other. If the crystal is to dissolve, these bonds must be broken.

Negative chloride ions on the surface are attracted by neighboring positive sodium ions and by the partially positive hydrogen atom in the polar water molecule (See the graphic on the left).

Similarly, the positive sodium ions are attracted by both chloride ions and the partially negative oxygen atom in the polar water molecule.

A "tug-of-war" occurs for the positive and negative ions between the other ions in the crystal and the water molecules. Several water molecules are attracted to each of the ions.

Whether the crystal dissolves is determined by which attractive force is stronger. If the internal ionic forces in the crystal are the strongest, the crystal does not dissolve. This is the situation in reactions where precipitates form. If the attractions for the ions by the polar water molecules are the strongest, the crystal will dissolve. This is the situation in sodium chloride."



Offline Perrin

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Re: NaCl and AgCl solutivity question
« Reply #2 on: August 10, 2008, 10:26:03 AM »
Still, the bond of NaCl is stronger than the bond of AgCl, and still it dissolves in water while AgCl doesn't. What's the explanation to it (other than the Ksp)?

Offline enahs

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Re: NaCl and AgCl solutivity question
« Reply #3 on: August 10, 2008, 10:56:45 AM »
The simplest explanation of this involves Enthalpy, Entropy, Crystal Lattice structure and energy, Hydration energies and solvation energies. I will be happy to type it up for you if you think you will understand it, but at the same time I can not explain in detail everything of those things as well. Other wise I might as well start working on a book!

Offline Perrin

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Re: NaCl and AgCl solutivity question
« Reply #4 on: August 10, 2008, 11:07:31 AM »
I might as well give it a try, if it's not too much problem! Thanks

Offline enahs

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Re: NaCl and AgCl solutivity question
« Reply #5 on: August 10, 2008, 12:23:01 PM »
Keep in mind this is a very simplified explanation. There are books on just this subject, and any good Inorganic Chemistry book should have a chapter devoted to this type of stuff.


Solutions of non-polar solutes in non-polar solvents demonstrate the simplest type.The forces involved in all interactions are London-Dispersion forces, which are relatively weak. The presence of these forces producing a condensed phase is the only difference of the mixing of an ideal gas. As with the latter case, the driving force is entropy of mixing. In an ideal solution at constant temperature, ΔHmixing = 0, and so the free energy will be comprised of nothing but Entropy terms:
ΔG = ΔH - TΔS    , if ΔH = 0 ---> ΔG = -TΔS


At the other extreme of the ideal solution of non-polar substances can be considered as solutions of ionic compounds in a very polar solvent, such as water, and thus starts to address your question. The entropy change for such a process can be negative or positive, unlike discussed above. As the ions go from solid into solution there will be an increase in disorder as expected, but there will also be an ordering of the solvent molecules as the ions become solvated. For large ions of low charge, the positive term will dominate, but for ions that interact strongly with water (small size, high charge) the negative term becomes more pronounced.

For an ionic compounds to dissolve, the Mandelung energy between ions in the lattice must be overcome. In a solution in which the ions are separated by molecules of solvent with a high dielectric constant, the attractive forces will be considerably less.

The mechanism of the solution of ionic compounds in water may be analyzed by a Born-Haber type of cycle. The over all enthalpy is the enthalpy of  dissociating the ions from the lattice (lattice energy) and the enthalpy of introducing the dissociated ions to the solvent (solvation energy).

Two factors will contribute to the magnitude of the enthalpy of solvation. The ability of the solvent molecules to coordinate strongly to the ions. Polar solvents are able to coordinate through attraction of the solvent dipole to the solute ions. The second factor is the type of ion involved, and in this simplified explanation we will stick with just its size. The number and strength of the interactions between the ion and solvent molecules will depend on how large the latter is. The lattice energy of the solute also depends on the ionic size.  The forces of the lattice are stronger (ion-ion) that those holding the solvent molecules to the ion (ion-dipole), however there are several of the ion-dipole interactions for each ion. The result is that the enthalpy of solvation is roughly of the same order of magnitude as the lattice enthalpy, and so the total enthalpy of the solution can be negative or positive depending upon the particular compound. When the enthalpy of solution is negative and the entropy of solution is positive, the free energy of solution very favorable since then both enthalpy and entropy of solution(s) reinforce each other.

The fact that the solubility of a salt depends upon the enthalpy of solution raises an interesting question concerning the magnitude of this quantity. Obviously a large solvation entropy contributes towards a favorable enthalpy of solution. But we find that solvation enthalpy alone is not a very good and predicting solubility. Both water soluble and insoluble salts are known with large and small hydration energies. Countering the hydration energy is the lattice energy. Both lattice and hydration energies are favored by large Z (charge) and small r (size). The difference is in the nature of the dependence upon distance. Using the Born-Lande equation, and I am going to skip the math and writing the equations,  it can be shown that the lattice energy is inversely proportional to the sum of the raddii, whereas the hydration enthalpy is the sum of two quantites is inversely proportional individual radii.  To skip even more math I am going to go straight to describing a physical picture (but the math is the bases of the picture, keep studying chemistry!).

The lattice energy is favored when the ions are similar in size, the presence of either a much larger cation or anion can effectively reduce it. In contrast, the hydration enthalpy is the sum of the two individual ion enthalpies (in the math I skipped for simplification sake), and if just one is very large (from a single, small ion) the total may still be sizable even if the counter ion is unfavorable.

It has also been pointed out the relation between the enthalpy of solution and the difference between hydration energy enthalpy of the cation and anion. This difference will be largest when the cation and anion differ mostly in size. In this case, the enthalpy of solution tends to be large and negative and favor solution. When the hydration enthalpies and size and more nearly alike the crystal is favored. When entropy effects are added there is a nice correlation with solubility and free energy of solution.

And for bonus.

The insolubility of ionic compounds in non-polar solvents is a similar phenomenon. The solvation energies are limited to those from ion-induced dipole forces, where are considerably weaker than
 ion-dipole forces, and thus not large enough to overcome the very strong ion-ion forces of the lattice energy.

The reason for the insolubility of non-polar solutes in some polar solvents such as water is not as apparent. The forces holding the solute molecules together tend to be very weak London forces (the forces keeping the lattice from dissolving). The forces between the water and the solute are also weak (dipole-induced dipole) but expected to be somewhat stronger then that London forces. You might suppose that the small solvation energy plus the entropy of mixing would be sufficient to cause a nonpolar solute to dissolve. It does not because any entropy resulting from the disordering of the hydrogen bonded structure of the solvent water molecules is more then offset by the loss of energy from the breaking of the hydrogen bonds. You can say that the solute would willingly dissolve but the water is much more happy associating with its self.


This is all simplified, and probably poorly explained.
In short, it is complicated, learn more chemistry, it is fascinating stuff! And there are so many amazing properties of water that it is well, amazing, and why most people think life beyond the microscopic cellular level is not possible without it.


Offline Perrin

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Re: NaCl and AgCl solutivity question
« Reply #6 on: August 10, 2008, 06:49:11 PM »
Thanks for the detailed response, it is very much appreciated!
It took me a while to work through all that stuff, and there's still much more for me to learn. I do find chemistry quiet fascinating, so I will keep on going.

Offline Mitch

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Re: NaCl and AgCl solutivity question
« Reply #7 on: August 10, 2008, 07:32:52 PM »
enahs: Wow! 5 karma points awarded.

Perin: Although NaCl readily dissolves in water, one might suspect that there are solvents that AgCl will dissolve easier in than NaCl. Always remember, solvent plays a huge role in chemistry. Maybe that's why I'm a gas-phase chemist, one less thing to worry about.
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