why is magnesium sulphate more soluable than barium sulphate? Watch

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john !!
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my thoughts on this:

it depends on lattice enthalpy and hydration enthalpy.
magnesium cation has a smaller ionic radius, therefore the sulphate ion can be held more tightly so the lattice enthalpy is higher.

however the hydration enthalpy of the magnesium sulphate is even higher because the delta +ve oxygen in the water molecule is more attracted to the small cation.

however, I then thought

the bariums electrons on the outer shell are held less tightly so they can be lost easier to the sulphate ion, so the lattice enthalpy is higher.

very confused, everything is contradicting in my head.
can someone put me straight please.

thanks very much, rep if you want
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BCHL85
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Lattice enthalpy of MgSO4 > that of BaSO4 (I dont consider the negative signs)
Hydration enthalpy of MgSO4 > that of BaSO4(negative values as well)
Enthalpy of solution = hydration enthalpy - lattice enthalpy
Hydration enthalpy changes more rapidly than lattice enthalpy -> MgSO4 is more soluble.
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john !!
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but why ?
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Golden Maverick
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(Original post by mik1w)
but why ?
I don't have a data book so can't check this but:
Lattice enthalpy:
barium sulphate > magnesium sulphate
-sulphate is a large anion and barium is larger than magnesium, therefore the packing of the ions will be closer in barium sulphate so the bonds will be stronger

Hydration enthalpy:
magnesium sulphate > barium sulphate
-the hydration enthalpy of the sulphate anion can be ignored in comparison, but magnesium as you said has a higher charge density so will have a higher hydration enthalpy

Solubility is proportional to (hydration enthalpy - lattice enthalpy)
in this case. So Magnesium sulphate will be more soluble. If you have a data book you could check this.
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Golden Maverick
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With the lattice enthalpy above, it may be that the increased charge density does serve to increase the bond strength in magnesium sulphate, and so it may as BCH be that it has a higher lattice enthalpy, but the steric considerations I talked about will have an effect and will contribute to magnesium sulphate being more soluble
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john !!
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are you sure barium sulphate bonds are stronger than magnesium sulphate?
I just want to check because I thought that the smaller the cation radius the closer the sulphate could get and the stronger the bond.
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Golden Maverick
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(Original post by mik1w)
are you sure barium sulphate bonds are stronger than magnesium sulphate?
I just want to check because I thought that the smaller the cation radius the closer the sulphate could get and the stronger the bond.
In lattice enthalpy there is a tradeoff between the charge density of the cation and it's size. I don't know if the bonds will actually be stronger in barium sulphate.
There is an eqution to express this I know but I cant remember it or find it! Hmm
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Golden Maverick
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Found it:
http://www.uh.edu/~chembi/lecture3.PDF

L = const x (z+ * z-)/(r+ + r-)

where z+ and z- are the charge densities and r+ and r- the radii.
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BCHL85
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Size of Mg(2+) is smaller than that of Ba(2+), so it polarises SO4(2-) more than Ba(2+). Therefore MgSO4 is more polarised, which makes it less stable. So heat needed to break the bond MgSO4 is smaller than that of BaSO4
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Golden Maverick
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Another page:
http://www.teachmetuition.co.uk/Chem...vision_ans.htm

6a. What happens to the solubility of group II Sulphates as you go down the group?

It decreases.
(1 mark)
6b. Explain this trend using the terms hydration enthalpy and lattice enthalpy in your answer.

Lattice enthalpy decreases down the group, and so does hydration enthalpy. It gets easier to break the ions free from the crystal lattice and also to hydrate them. However the hydration enthalpy decreases at a slower rate therefore solubility decreases.
What BCH said.
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