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# Entropy and Solubility

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1. Question: "Explain, using the concepts of lattice enthalpy and hydration enthalpies, why CaCO3 will not dissolve at 298K"

Lattice energy and hydration energy are both exothermic and the factors affecting the hydration enthalpy and lattice energy are ionic radius and charge density.
So as Ca is a 2+ ion, the charge density is bigger. This means a more positive exothermic lattice energy/hydration energy, so is this why is can dissolve?

I'm not sure what the ionic radius of the ions are though, but a larger radius means a smaller charge density - so the values are less exothermic. So does this increase the solubility?

Also, does charge density always refer to the number of electrons only?

Thanks
2. i think its because of entropy.
In lattice enthalpy the gaseous ions go from:
gaseous + gaseous ---> solid.

In hydration the ions go from
gaseous + gaseous ---> gaseous + aqueous.

Im both cases its decreased in entropy. meaning that when we use this equation:
ΔG=ΔH-TΔS
TΔS>ΔH so it wont dissolve
3. (Original post by lekha2611)
Question: "Explain, using the concepts of lattice enthalpy and hydration enthalpies, why CaCO3 will not dissolve at 298K"

Lattice energy and hydration energy are both exothermic and the factors affecting the hydration enthalpy and lattice energy are ionic radius and charge density.
So as Ca is a 2+ ion, the charge density is bigger. This means a more positive exothermic lattice energy/hydration energy, so is this why is can dissolve?

I'm not sure what the ionic radius of the ions are though, but a larger radius means a smaller charge density - so the values are less exothermic. So does this increase the solubility?

Also, does charge density always refer to the number of electrons only?

Thanks
sorry forgot to quote ^^
4. (Original post by asaaal)
i think its because of entropy.
In lattice enthalpy the gaseous ions go from:
gaseous + gaseous ---> solid.

In hydration the ions go from
gaseous + gaseous ---> gaseous + aqueous.

Im both cases its decreased in entropy. meaning that when we use this equation:
ΔG=ΔH-TΔS
TΔS>ΔH so it wont dissolve

I'm on the Edexcel Specification, so I'm guessing ΔG=ΔH-TΔS is not on there

Is G Gibbs energy (heard of this term, but never used it) ? Is TΔS the temperature x total entropy?
5. (Original post by lekha2611)

I'm on the Edexcel Specification, so I'm guessing ΔG=ΔH-TΔS is not on there

Is G Gibbs energy (heard of this term, but never used it) ? Is TΔS the temperature x total entropy?
in that case completely disregard my answer! Im on OCR and thats the only answer i could think of to do with my exam board.
The answer is most likely to be something my exam board hasnt done
good luck !
6. (Original post by asaaal)
in that case completely disregard my answer! Im on OCR and thats the only answer i could think of to do with my exam board.
The answer is most likely to be something my exam board hasnt done
good luck !
Could you explain your answer though please? I'm curious actually, maybe using your method I'll actually understand this part of the topic, because it's the only thing I'm really struggling with at the moment :P And then I'll use it to explain using Edexcel terms :P
7. (Original post by lekha2611)
Question: "Explain, using the concepts of lattice enthalpy and hydration enthalpies, why CaCO3 will not dissolve at 298K"

Lattice energy and hydration energy are both exothermic and the factors affecting the hydration enthalpy and lattice energy are ionic radius and charge density.
So as Ca is a 2+ ion, the charge density is bigger. This means a more positive exothermic lattice energy/hydration energy, so is this why is can dissolve?

I'm not sure what the ionic radius of the ions are though, but a larger radius means a smaller charge density - so the values are less exothermic. So does this increase the solubility?

Also, does charge density always refer to the number of electrons only?

Thanks
Lattice enthalpy is NOT exothermic when you are breaking the lattice, it is highly endothermic.

The hydration enthalpy of the ions is exothermic.

The term -TΔS is negative as the entropy increases on dissolution.

A substance is insoluble when the lattice enthalpy plus the hydration enthalpy is positive and bigger then -TΔS.

This can happen for small highly changed ions, however it´s not so easy to predict, as the association between ions and water also reduces the total entropy. A small ion cannot association water molecules very efficiently and so does not reduce the entropy so much.

From the above, the oxides and sulphides (both double charged) of double positive metal ions are usually insoluble.
8. (Original post by charco)
Lattice enthalpy is NOT exothermic when you are breaking the lattice, it is highly endothermic.
Lattice enthalpy is the formation of a solid ionic lattice from its gaseous ions. It's bond formation and therefore it is exothermic.
To break down a lattice it is endothermic, but that's not what lattice enthalpy is.
9. (Original post by Mathlete 4 the win)
Lattice enthalpy is the formation of a solid ionic lattice from its gaseous ions. It's bond formation and therefore it is exothermic.
To break down a lattice it is endothermic, but that's not what lattice enthalpy is.
There is no universal consensus regarding the definition of lattice enthalpy with different boards defining it in different ways.

I refer you to the IB for example:

10. (Original post by charco)
There is no universal consensus regarding the definition of lattice enthalpy with different boards defining it in different ways.

I refer you to the IB for example:

I see, the majority however will probably use the exothermic definitions : http://en.wikipedia.org/wiki/Lattice_energy
11. (Original post by Mathlete 4 the win)
I see, the majority however will probably use the exothermic definitions : http://en.wikipedia.org/wiki/Lattice_energy
I guess the message here is "don't trust wikipedia" ....

The most important thing is to find out the specifics of YOUR examinations board and approach the problems accordingly.
12. I dont get it, does solubility increase with an increase in entropy or not?

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