When I was reading my notes I came up with a problem. I read that for the s block element salts, when checking the solubility, we consider the hydration enthalpy so we decide their stability in water. But then there is this exception: hydroxides don't follow this. They depend on lattice enthalpy for their stability. Which has a greater dominance and why? Is it the lattice energy or the hydration enthalpy? How can we differentiate them, and where to use these concepts for deciding stability? As in the case of hydroxides lattice energy plays a dominant role, not the hydration energy: why is it so?


2 Answers 2


The dominant factor depends on situation to situation.Here is an explanation i found helpful while predicting solubilities

The hydration energy is of Form A÷r1 + B÷r2 .

The lattice energy is of form C÷(r1+r2). Here r1 and r2 are radius of cation and anion while A,B,C are constants.

When Size of anion is large example Iodine,Lattice energy does not decrease with increase in radius but hydration energy decreases drastically since r2>>r1(See formula) Therefore in such cases,Solubility of Salts decreases with radius of cation. Order- LiI is more than NaI is more than KI(in terms of solubility)

Now when anion size is comparable(eg Fluorine),The lattice energy decreases much more rapidly than hydration energy(see formula).Hence Solubility increases with increase in cationic size.Thefore order is LiF is less than NaF is less than KF


lets say i am comparing solubility of sulfate and hydroxides of alkaline earth metals.when comparing solubility two terms come into effect.Lattice energy and the hydration energy both contribute to the solubility but sometimes there is a dominance of one term over the other.In order to account for this dominance a new term was introduced. The resultant of two effects i.e.

ΔHsolution = ΔHlattice - ΔHHydration

This term plays the major role in deciding the stability of salts inside water and hence their solubility. Lets first get to know the definition of lattice energy. The lattice energy of a crystalline solid is usually defined as the energy of formation of a crystal from infinitely-separated ions and as such is invariably negative. The precise value of the lattice energy may not be determined experimentally, because of the impossibility of preparing an adequate amount of gaseous ions or atoms and measuring the energy released during their condensation to form the solid. However, the value of the lattice energy may either be derived theoretically from electrostatics or from a thermodynamic cycling reaction, the Born–Haber cycle.

By the formula given for the lattice energy we can see that lattice energy is directly proportional to the charge present on ions involved in bonding and inversely proportional to the size of the ions.It is a major contributor in deciding melting points and thermal stability in ionic salts.

Coming to hydration enthalpy,Hydration energy (also hydration enthalpy) is the amount of energy released when one mole of ions undergo hydration which is a special case of solvation. It is a special case of dissolution energy, with the solvent being water.

For example, upon dissolving a salt in water, the outermost ions (those at the edge of the lattice) move away from the lattice and become covered with the neighboring water molecules. If the hydration energy is equal to or greater than the lattice energy, then the salt is water-soluble. In salts for which the hydration energy is higher than the lattice energy, solvation occurs with a release of energy in the form of heat. For instance, CaCl2 (anhydrous calcium chloride) heats the water when dissolving. However, the hexahydrate, CaCl2·6H2O cools the water upon dissolution. The latter happens because the hydration energy does not completely over come the lattice energy, and the remainder has to be taken from the water in order to compensate the energy loss.

Coming to the question now that we have understood the terms involved,we know that sulfate ion is larger in size compared to the hydroxide ion. Coming to hydroxides of alkaline earth metals

The solubility of the alkaline earth metal hydroxides in water increases with increase in atomic number down the group. This is due to the fact that the lattice energy decreases down the group due to increase in size of the alkaline earth metals cation whereas the hydration energy of the cation remains almost unchanged.Due to its small size the lattice energy plays a dominant role in deciding its solubility whereas hydration energy almost remains constant.

Coming to sulfates, solubility:The solubility of the sulfates in water decreases down the groups i.e. Be > Mg > Ca > Sr > Ba. Thus BeSO4 and MgSO4 are highly soluble, CaSO­4 is sparingly soluble but the sulphates of Sr, Ba and Ra are virtually insoluble. Reason The magnitude of the lattice energy remains almost constant as the sulfate is so big that small increase in the size of the cation from Be to Ba does not make any difference. However the hydration energy decreases from Be+2 to Ba+2 appreciably as the size of the cation increase down the group. Hence, the solubility of sulfates of alkaline earth metals decrease down the group mainly due to the decreasing hydration energies from Be+2 to Ba+2. The high solubility of BeSo4 and MgSO4 is due to high hydration energies due to smaller Be+2 and Mg+2 ions.

This effect can also be explained by the fact that ionic character increases down the group.By fajan's rule more covalent character would mean smaller size of cation hence more polarization effect in solution state.More polarization leads to dissociation of compound in water easily hence the solubility increases from bottom to top for the sulfates.


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