Chemguide: Support for CIE A level Chemistry


Learning outcome 10.1(g)

This statement asks for the trend in the solubilities of the Group 2 sulphates and hydroxides, and an explanation for that trend.

Before you go on, you should find and read the statement in your copy of the syllabus.

Notice that the statement is in bold type, which means that it will only be examined at the end of a two year course.


The Chemguide pages about this are far more complicated than you will need for the sort of questions that CIE ask, and the answers they expect, so I will cover most of what you need to know below.

The explanation given on this page is only a part of the proper explanation. If you go on to do Chemistry at a higher level, you should be aware that this explanation is very over-simplified. It is, however, what CIE expect.

You won't be able to follow this explanation unless you have already covered enthalpies of solution, lattice enthalpies and hydration enthalpies from the energetics part of the syllabus (section 5).


The trends

Sulphates

The sulphates become less soluble as you go down the group.

  • Magnesium sulphate is soluble in water.

  • Calcium sulphate is only very slightly soluble in water.

  • Strontium and barium sulphates are virtually insoluble in water.

You almost certainly know enough simple chemistry to be able to work this trend out.

At some point, you will certainly have met the reaction between magnesium and dilute sulphuric acid to give hydrogen and a solution of magnesium sulphate. You know that magnesium sulphate is soluble.

You will also have come across the test for a sulphate by adding barium chloride (or nitrate) solution to a solution of a sulphate. You get a white precipitate of barium sulphate. So you know that barium sulphate is insoluble.

Sulphates become less soluble as you go down the group.


Hydroxides

The hydroxides become more soluble as you go down the group.

None of them are very soluble, but the solubility increases as you go down the Group. See the beginning of the page about the solubility of the hydroxides (etc).


The explanation for sulphates

The simple explanation is in terms of the changes which occur when an ionic compound dissolves in water. Energy has to be supplied to break up the lattice of ions, and energy is released when these ions form bonds with water molecules.

Breaking up the lattice

To break up an ionic lattice, you need to supply lattice dissociation enthalpy.

The size of the lattice dissociation enthalpy depends on the charges on the ions, and the distances between their centres. All of the Group 2 sulphates consist of 2+ ions attracting 2- ions, and so the only thing that matters is the distance between the ion centres.

As you go down the group, the energy needed to break up the lattice falls as the positive ions get bigger. The bigger the ions, the more distance there is between their centres, and the weaker the forces holding them together.

Releasing energy by forming bonds with water molecules

Energy is released as hydration enthalpy when water molecules cluster around the free metal ions and sulphate ions.

As the positive ions get bigger, the energy released as the ions bond to water molecules falls. Bigger ions aren't so strongly attracted to the water molecules.

The overall effect

Both lattice dissociation enthalpy and hydration enthalpy fall as you go down the group. What matters is how fast they fall relative to each other.

As you go down the group, the lattice dissociation enthalpies don't fall as much as the hydration enthalpies of the positive ions.

The size of the hydration enthalpy of a positive ion is due only to the size of that ion. But that isn't so for lattice dissociation enthalpy.

The lattice dissociation enthalpy is governed by the distance between the centres of the ions, and that is made up of the radius of the large sulphate ion, plus the radius of the smaller positive ion.

The effect of the change in size of the positive ion is being diluted by the presence of the large sulphate ion.

Putting some numbers on this

This is much easier to understand if you have got some numbers to work with. All the values in the table are in kJ per mole.


Note:  I have no confidence that the numbers I am going to use are reliable. There seem to be as many different values for these as there are sources of information (more about this at the bottom of the page).


lattice enthalpyhydration enthalpy M2+hydration enthalpy SO42-overall change
CaSO4+2653-1583-1137-67
SrSO4+2603-1450-1137+16

You can see that the lattice dissociation enthalpy has fallen by only 50 kJ, whereas the hydration enthalpy of the positive ion has fallen by 133 kJ. A bit less heat had to be put in in order to break the lattice, but quite a lot less was given out when the ions bonded to the water.

The net effect is that the overall process becomes less exothermic (or, in this case, actually becomes endothermic).


The explanation for hydroxides

The underlying explanation is still the same. Both lattice dissociation enthalpy and hydration enthalpy fall as you go down the group, and what matters is how fast they fall relative to each other.

Hydroxide ions are much smaller than sulphate ions, and so the size of the positive ion makes up a greater proportion of the distance between the positive and negative ions in the hydroxide case.

In this case, the lattice dissociation enthalpy falls by more than the hydration enthalpy as you go down the group.

As you go down the group, the energy you need to put in falls by more than the energy you get out. That makes the overall process more exothermic as you go from magnesium hydroxide to barium hydroxide.

There is a question involving some calculations about the relative solubilities of magnesium and strontium hydroxides on a past paper. See May /June 2010 paper 42 Q2(b) together with its mark scheme. Part (a) of that question asked about the solubilities of the sulphates.


What do CIE expect you to say?

A question asking about the solubility of the Group 2 sulphates would probably ask you to state and explain the trend in solubilities of the sulphates of Group 2 elements. Your answer would need to include:

For sulphates:

  1. Solubility decreases as you go down the group.

  2. The lattice dissociation enthalpy and hydration enthalpy both decrease as you go down the group.

  3. The hydration enthalpy decreases more than the lattice dissociation enthalpy.

  4. Therefore the enthalpy of solution becomes more endothermic (or less exothermic).

For hydroxides:

  1. Solubility increases as you go down the group.

  2. The lattice dissociation enthalpy and hydration enthalpy both decrease as you go down the group.

  3. The lattice dissociation enthalpy decreases more than the hydration enthalpy.

  4. Therefore the enthalpy of solution becomes more exothermic (or less endothermic).


This is based on mark schemes available at the time of writing. Notice that the depth of understanding they want is really limited. There isn't any need to explain why the lattice dissociation enthalpies and the hydration enthalpies change the way they do.


Read this if you are likely to do chemistry beyond A level:  I have no confidence whatsoever in this explanation, even though you will find it in lots of books, and all over the web. It looks to me like one of those explanations which "everybody knows", but which falls to pieces as soon as you look at the data. There seem to be a lot of these in inorganic chemistry at this level!

You will find the problems discussed in some detail on the page problems in explaining the solubility of Group 2 compounds. Don't even think about reading this unless your chemistry is really good.

The problem basically is that it is impossible to explain these patterns unless you include entropy in your explanation. Any explanation which doesn't include entropy is at best incomplete, and at worst, wrong.

I haven't been able to find any reliable data for this topic. Different data sources give different values both for lattice energies and hydration energies. Although values for calcium sulphate and strontium sulphate produce the same result whatever source you use (i.e. that strontium sulphate is likely to be less soluble than calcium sulphate), that doesn't hold true if you extend it to barium sulphate. Selecting values to fit your hypothesis, and ignoring others, is just bad science.

If anyone knows where I can get reliable values for the necessary lattice enthalpies and hydration enthalpies for all the Group 2 sulphates, could you let me know via the address on the about the CIE section page. I would also like to know why you think that particular set of values is reliable.




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© Jim Clark 2010 (last modified May 2014)