What is the order of solubility of carbonates ,sulphates & hydrogencarbonates in metal carbonates , metal sulphates and metal hydrogencarbonates...???

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The simplest approach is to consider thetwo enthalpy change trends.

• Theprocess of dissolving can be analysed in terms of two theoretical stagese.g. for simple cation–anion ionic compound.

  • In the arguments outlined belowMⁿ⁺could be Gp1 or Gp2 metal cation etc.,X^n–could be halide, oxide, hydroxide, sulphate, carbonate anion etc., and n is the charge on ion – the n's may be different or the same):

• (1)Mⁿ⁺_aX^n–_b(s)==> aMⁿ⁺_(g)+ bX^n–_(g)(breaking the lattice apart into its constituent ions)

• This process is always endothermic, and is called thelattice enthalpy. Its usually defined in the opposite direction by saying it is 'the energy released when 1 mole of an ionic lattice is formed from its constituent gaseous ions' (at 298K, 1 atmos./101kPa pressure).

• *The lattice enthalpy decreases down the group as the cation radius increases (anion radius constant for a particular seriese.g. sulphates). Therefore, energetically, the solvationin terms of lattice energy is increasingly favoured down the group.

• (2) Mⁿ⁺_(g)+ aq==>Mⁿ⁺_(aq)andX^n–_(g)+ aq==>X^n–_(aq)

  • Representing the solvation–hydration of ions.

• The equations above represent to the two 'hydration enthalpies', the heat released when an isolated gaseous ion becomes solvated by water to form an aqueous solution (1M, 298K, 1 atmos./101kPa pressure)

• *The hydration enthalpy for the cation decreases down the group as the radius gets larger. Therefore, energetically, the solvation is less favoured down the group as the cation radius increases.

* In both cases the numerical enthalpy value increases the smaller the radii as charges closer,and the greater the ionic charge (constant for a series), both factors increase the electrical attraction of either cation–anion in the crystal or ion–water in aqueous solution.