Why can dissolving salt in water be exothermic/endothermic?

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pernero
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Why can dissolving salt in water be exothermic/endothermic?

Postby pernero » Fri Jul 31, 2009 6:59 pm UTC

To rephrase that,
Why can dissolving salt in water be exothermic OR endothermic?
I understand that the steps of dissolving salt in water: Breaking the ionic bonds of the salt, breaking the hydrogen bonds in the water, and bonding water molecules to the ions. But of these processes, which are endothermic, which are exothermic and why do some salts produce different results?

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BlackSails
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby BlackSails » Fri Jul 31, 2009 7:17 pm UTC

When you add a solute to a solvent, you are reducing the solvent-solvent interactions, and adding solute-solvent interactions. If the solute-solvent interactions are really favorable compared to the solvent-solvent interactions, the process is exothermic.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby pernero » Fri Jul 31, 2009 7:19 pm UTC

But what determines how favorable it is?

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby opsomath » Fri Jul 31, 2009 8:09 pm UTC

The answer: a lot of things. Basically, it's the difference between the stability (total energy) of a unit of the solid (in its crystal lattice) vs. that same unit in dissolved form (with its ions separated and surrounded by water molecules).

This is determined by a bunch of factors. For instance, dissolving calcium salts in water is typically very exothermic because the calcium ion is relatively small and has a charge of +2, enabling very, very strong interactions with the negative end of water molecules. These release a lot of energy when they form which winds up as heat. On the other end of the scale, you have ammonium cations, (NH4+) which are relatively large. They form okay interactions with water molecules, but because of their size and shape they bust up a lot of favorable water-water interactions, which takes a lot of energy. Thus, it's endothermic to dissolve most ammonium salts in water.

The short version; it's complicated.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby pernero » Fri Jul 31, 2009 8:35 pm UTC

Thank you!

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Game_boy » Fri Jul 31, 2009 8:52 pm UTC

Remember chemistry is an abstraction of quantum mechanics etc. . In some cases (for more complex molecules) we haven't expressed, let alone solved, the equations that determine energy stored or transferred during a reaction*. So, though we can measure the enthalpy change, sometimes we can't say why in simple terms.

However if you make the generalisations such as atoms are point particles with charge you can make useful predictions.

*Though they are entirely determined by basic laws we can write down. It's not arbitrary, just hard.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Charlie! » Sat Aug 01, 2009 12:32 am UTC

I think some of the replies aren't entirely in the right direction :D

The right direction is, of course, thermodynamics. In short, the second law of thermodynamics determines whether or not a process is favorable. Just ask yourself "does entropy increase?"

Entropy changes in dissolving in two ways. The entropy of the whole thing is raised or lowered as the temperature raises or lowers. Additionally, the entropy increases just because dissolved molecules are more "free." So for all favorable processes, the entropy from dissolving + the entropy from the temperature change > 0. This allows the temperature change to be negative, as long as it's balanced out by the entropy from dissolving.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Seraph » Sat Aug 01, 2009 5:51 am UTC

Charlie! wrote:I think some of the replies aren't entirely in the right direction :D

The right direction is, of course, thermodynamics. In short, the second law of thermodynamics determines whether or not a process is favorable. Just ask yourself "does entropy increase?"

I think your reply is the one that isn't entirely in the right direction. "Exothermic" and "Endothermic" describe the sign of the Enthalpy of a process. It has nothing to do with the entropy of a process. Freezing water is one example of a process where Enthalpy increses while entropy decreases.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Talith » Sat Aug 01, 2009 1:32 pm UTC

Seraph wrote:
Charlie! wrote:I think some of the replies aren't entirely in the right direction :D

The right direction is, of course, thermodynamics. In short, the second law of thermodynamics determines whether or not a process is favorable. Just ask yourself "does entropy increase?"

I think your reply is the one that isn't entirely in the right direction. "Exothermic" and "Endothermic" describe the sign of the Enthalpy of a process. It has nothing to do with the entropy of a process. Freezing water is one example of a process where Enthalpy increses while entropy decreases.

Only local entropy decreases, necessarily the entropy of the environment that the water is in has to increase more than the entropy of the water system decreases (through the transfer of energy to non-water particles).

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby andyisagod » Sat Aug 01, 2009 2:33 pm UTC

Seraph wrote:
Charlie! wrote:I think some of the replies aren't entirely in the right direction :D

The right direction is, of course, thermodynamics. In short, the second law of thermodynamics determines whether or not a process is favorable. Just ask yourself "does entropy increase?"

I think your reply is the one that isn't entirely in the right direction. "Exothermic" and "Endothermic" describe the sign of the Enthalpy of a process. It has nothing to do with the entropy of a process. Freezing water is one example of a process where Enthalpy increses while entropy decreases.


I dissagree his reply was in my opnion excellent, Entropy is at the heart of chemistry while quantum mechanics may allow you to calculate the energy of the bonds broken and formed in the reaction, it is the entropy which will allow the reaction to take place. This is important as otherwise the endothermic reaction would lower the entropy of the system and so you would never see such salts disolve.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby BlackSails » Sat Aug 01, 2009 3:23 pm UTC

andyisagod wrote:
I dissagree his reply was in my opnion excellent, Entropy is at the heart of chemistry while quantum mechanics may allow you to calculate the energy of the bonds broken and formed in the reaction, it is the entropy which will allow the reaction to take place. This is important as otherwise the endothermic reaction would lower the entropy of the system and so you would never see such salts disolve.


Its both. The gibbs free energy includes both enthalpy change and entropy change. The entropy of dissolution is (as far as I can see) always positive, because you have some bulk solid being distributed into a liquid.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Charlie! » Sat Aug 01, 2009 4:44 pm UTC

Seraph wrote:
Charlie! wrote:The right direction is, of course, thermodynamics. In short, the second law of thermodynamics determines whether or not a process is favorable. Just ask yourself "does entropy increase?"

I think your reply is the one that isn't entirely in the right direction. "Exothermic" and "Endothermic" describe the sign of the Enthalpy of a process. It has nothing to do with the entropy of a process. Freezing water is one example of a process where Enthalpy increses while entropy decreases.

I was replying to pernero's questions "Why can dissolving salt in water be exothermic OR endothermic?" and "But what determines how favorable it is?" rather than talking specifically about what determines the enthalpy (heat) of the process. I'm still pretty sure that was the right direction.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Tass » Sat Aug 01, 2009 5:42 pm UTC

BlackSails wrote:
andyisagod wrote:
I dissagree his reply was in my opnion excellent, Entropy is at the heart of chemistry while quantum mechanics may allow you to calculate the energy of the bonds broken and formed in the reaction, it is the entropy which will allow the reaction to take place. This is important as otherwise the endothermic reaction would lower the entropy of the system and so you would never see such salts disolve.


Its both. The gibbs free energy includes both enthalpy change and entropy change. The entropy of dissolution is (as far as I can see) always positive, because you have some bulk solid being distributed into a liquid.


It is actually entirely possible for an organic salt to have a negative entropy of solvation because of hydrophobic effect.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby AFedchuck » Tue Aug 04, 2009 5:25 pm UTC

Tass wrote:
BlackSails wrote:
andyisagod wrote:
I dissagree his reply was in my opnion excellent, Entropy is at the heart of chemistry while quantum mechanics may allow you to calculate the energy of the bonds broken and formed in the reaction, it is the entropy which will allow the reaction to take place. This is important as otherwise the endothermic reaction would lower the entropy of the system and so you would never see such salts disolve.


Its both. The gibbs free energy includes both enthalpy change and entropy change. The entropy of dissolution is (as far as I can see) always positive, because you have some bulk solid being distributed into a liquid.


It is actually entirely possible for an organic salt to have a negative entropy of solvation because of hydrophobic effect.

Don't even need the hydrophobic effect.
Remember the same argument that you used in the enthalpy. Coordination of a water molecule reduces its entropy, as its lost its translational freedom. If the coordination is sufficiently strong the entropy of solvation can be negative. I'll try to find some examples.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby meat.paste » Wed Aug 05, 2009 5:58 pm UTC

AFedchuck wrote:Don't even need the hydrophobic effect.
Remember the same argument that you used in the enthalpy. Coordination of a water molecule reduces its entropy, as its lost its translational freedom. If the coordination is sufficiently strong the entropy of solvation can be negative. I'll try to find some examples.


Maybe. Don't forget about the huge entropy increase with the phase change of the salt going from solid to solution. I would think that this entropy change would dwarf the entropy decrease from coordinating solvent molecules. Maybe if there was a high charge density salt (for maximum binding strength to water) that didn't stick to itself readily (to minimize the entropy gain upon dissolution), then the overall entropy change would be negative. Unfortunately, the high charge density salts, like Al3+ or PO43-, tend to want to stick to each other. I am very curious to see what solid would dissolve in water to give a negative entropy change. If you can find one that simultaneously cools down (+deltaH), I'll be really impressed :)

Back to the OP - Breaking bonds requires an energy input, making bonds gives energy back. So, breaking the ionic bonds of the salt and breaking the hydrogen bonds of the water both require energy. Making the ion-water bonds releases energy. Depending on how much energy it takes to break the ionic bond (higher charge states will require more energy) versus how much energy you get from the ion-water bonds (higher charge densities [charge state divided by ion size] bind more water and give back more energy) will determine how endo- or exothermic the reaction is.

I wonder how much of the endothermicity of ammonium dissolution is from the ease of dissociation into NH3. As an example, ammonium nitrate is a very endothermic dissolution. This could be because the salt can dissociate as NH4+ and NO3- or as NH3 and HNO3. Clearly, the second reaction would not have electrostatic bonds to overcome.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby oxoiron » Wed Aug 05, 2009 6:32 pm UTC

meat.paste wrote:I wonder how much of the endothermicity of ammonium dissolution is from the ease of dissociation into NH3. As an example, ammonium nitrate is a very endothermic dissolution. This could be because the salt can dissociate as NH4+ and NO3- or as NH3 and HNO3. Clearly, the second reaction would not have electrostatic bonds to overcome.
The amount of NH3 and HNO3 present in solution would be so small as to be immeasurable. I can't imagine how it could significantly affect the thermodynamics of the situtation.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby meat.paste » Thu Aug 06, 2009 3:22 pm UTC

The NH3 concentration may not be. It's a pretty weak base. I had a brain fart on the HNO3. You are correct on that.

Still, the relatively low bond strength (judging from ammonium's tendency to cool off when dissolving in water) may be in part to the allowed pathway that does not involve breaking an ionic bond. Maybe something like:

NH4NO3 + H2O -> NH3 + H3O+ + NO3- -> solvated ammonia + solvated nitrate.

Once the ammonia is solvated, there will be the occasional proton exchange to make ammonium and hydroxide. From the Kb, which is ~10-5 IIRC, I think most of the ammonia will stay intact.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby oxoiron » Thu Aug 06, 2009 4:12 pm UTC

You are forgetting that initially [H3O+] = [NH3], so you are looking at this equilibrium:

H3O+ + NH3 <--> H2O + NH4+

I don't have the equilibrium constant in front of me, but my chemical intuition (my advisor always used to ask what my chemical intuition was telling me) says that it lies way to the right. If you have numbers showing otherwise, I would be happy to have my intuition proven wrong as has happened so many times before. :(
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby Tass » Thu Aug 06, 2009 5:02 pm UTC

Ammonium nitrate is acidic since nitric acid is a stronger acid than ammonia is as base (equivalently: nitrite is much weaker as a base than ammonium is as acid). Therefore pH is always going to be less than 7.

Since the pKb of ammonia is 5, it is going to be 99% on ammonium form at pH 7. So maximally 1% of the ammonium will dissociate, if the concentration is high enough to make the 1% matter to pH, then even less will dissociate.

Needless to say the protonation of nitrate will occur to an even smaller extend.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby meat.paste » Fri Aug 07, 2009 6:56 pm UTC

<grumble> You're going to make me get out a pen and paper to do this...

pKa of ammonium is 9.24 @ 25C. If the pH is 7, then [NH3]/[H4+] = .56%.

After plotting a curve of the heat of solvation versus the ionic radius for the nitrates of Li, Na, K, Rb, and Cs, the least squares fit for them is ~-10kcal/mol/angstrom. The NH4+ ion is 1.43A in diameter, and Rb+ is 1.47, yet the enthalpy difference is 2.58 kcal/mol (-6.14 kcal/mol for NH4+, -8.72 for Rb+). After normalizing the radii, NH4+ would be expected to need .40 kcal/mol less heat than Rb+, not 2.6. The difference is a lot more than .5%, so my theory is wrong.

Now I'm curious to know what the dealio is with NH4+.

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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby oxoiron » Fri Aug 07, 2009 7:58 pm UTC

If I'm understanding the last few posts, we are all in agreement that the ammonia concentration is so low that we don't really need to worry about it, so let's move on to meat.paste's rather interesting problem.

Since charge and size are essentially the same for Rb+ and ammonium, why are their solvation energies so different?

Again, I will call on my chemical intuition (don't fail me, baby!). We have three primary differences:

1) Effective nuclear charge
2) LUMO energy
3) Number of open coordination sites (NH4+ = 4 and Rb+ = 6)

I don't see how (1) will affect things, but because solvation depends in part on electron density donation, it occurs to me that (2) may be important. (3) may also produce an effect, because more bond formation means more energy released.

If I had to put my money on something, I'd go with (3), because it agrees with meat.paste's data. However, I'd be happy to hear arguments against my hypothesis as well as any other suggestions.
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Re: Why can dissolving salt in water be exothermic/endothermic?

Postby PM 2Ring » Sat Aug 08, 2009 4:39 pm UTC

Is the nitrogen inversion of ammonia (but not ammonium) relevant to this question?


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