Calcium oxalate is insoluble in acetic acid, but not it's carbonate, oxide and hydroxide.

From internet, it is even considered as a test to distinguish calcium oxalate from it's carbonate and oxide. The reason from the resources which i had searched, only gave factual reasons such as

  1. High Lattice Energy and
  2. Low solvation energy of $Ca(C_2O_4)$

But intuitively the "short sized" compounds like $CaO$ and $Ca(OH)_2$ must have better packing and more lattice energy, then why much bigger calcium oxalate have higher lattice energy?

Also, i am clueless about how to compare solvation energy of different compounds. I think it has to do with polarity of the solute and the solvent. Any help will be appreciated.

Edit: I had omitted the relative basicity of oxide, hydroxide and carbonate of calcium but i am more interested in knowing how to compare the lattice energy and solubility of the given compounds.

  • 1
    $\begingroup$ You have omitted to consider acid-basis relations, particularly the relative acidity/alkalinity of respective cunjugated acid/base pairs. $\endgroup$
    – Poutnik
    Apr 21, 2020 at 8:55
  • $\begingroup$ Yes, but what may be the reason for other facts, i am more interested in knowing how to compare properties like solubility and lattice energy $\endgroup$
    – user91694
    Apr 21, 2020 at 10:18
  • 1
    $\begingroup$ Such things are more related to solubility product. For solubility itself, you have to involved the side acid-base equilibrium reactions. The play essential role. $\endgroup$
    – Poutnik
    Apr 21, 2020 at 12:14

1 Answer 1


Calcium oxide, hydroxide and carbonate react with acetic acid to form calcium acetate. Calcium oxalate does not react.

Edit: In real life, you'll of course will end up with a little oxalic acid and calcium acetate, as calcium oxalate is very slightly soluble and then you'll mix all available ions.

  • $\begingroup$ If you take a solution of calcium acetate and add oxalic acid, you will precipitate calcium oxalate, leaving acetic acid in solution. That's because oxalic acid is a stronger acid than acetic, as Poutnik suggested. $\endgroup$ Apr 22, 2020 at 12:56
  • $\begingroup$ "That's because oxalic acid is a stronger acid than acetic". Probably doesn't work like this. Hydrochloric acid is stronger that oxalic acid, but calcium oxalate won't dissolve in it. It's rather the solubility product of calcium oxalate that's driving the precipitation. $\endgroup$ Apr 22, 2020 at 13:24
  • $\begingroup$ Reverse the comparison: Since oxalic acid is stronger than acetic acid, acetic acid is weaker than oxalic acid, and hence won't protonate oxalate (at least, not very much). Since acetic acid won't protonate oxalate, it can't bring calcium oxalate into solution. You are correct that the solubility product drives the precipitation, but to undrive it requires protonation of oxalate anion. CRC Handbook says calcium oxalate is soluble in acids; stronger acids will do better because a strong acid could protonate oxalate and drive down the concentration of oxalate anion as CaC2O4 dissolves. $\endgroup$ Apr 22, 2020 at 18:41
  • $\begingroup$ You're right, James. There is the equilibrium between free oxalate and oxalic acid which will draw the calcium oxalate into solution at low pH. My bad. $\endgroup$ Apr 23, 2020 at 6:36

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