Calcium oxalate is nearly insoluble in water, and only very slightly soluble in acetic acid, but is readily dissolved by the strong mineral acids.  This behavior with acids is explained by the fact that oxalic acid is a stronger acid than acetic acid; when, therefore, the oxalate is brought into contact with the latter there is almost no tendency to diminish the concentration of C_{2}O_{4}^{—­} ions by the formation of an acid less dissociated than the acetic acid itself, and practically no solvent action ensues.  When a strong mineral acid is present, however, the ionization of the oxalic acid is much reduced by the high concentration of the H^{+} ions from the strong acid, the formation of the undissociated acid lessens the concentration of the C_{2}O_{4}^{—­} ions in solution, more of the oxalate passes into solution to re-establish equilibrium, and this process repeats itself until all is dissolved.

The oxalate is immediately reprecipitated from such a solution on the addition of Oh^{-} ions, which, by uniting with the H^{+} ions of the acids (both the mineral acid and the oxalic acid) to form water, leave the Ca^{++} and C_{2}O_{4}^{—­} ions in the solution to recombine to form [CaC_{2}O_{4}], which is precipitated in the absence of the H^{+} ions.  It is well at this point to add a small excess of C_{2}O_{4}^{—­} ions in the form of ammonium oxalate to decrease the solubility of the precipitate.

The oxalate precipitate consists mainly of CaC_{2}O_{4}.H_{2}O when thrown down.]

[Note 5:  The small quantity of ammonium oxalate solution is added before the second precipitation of the calcium oxalate to insure the presence of a slight excess of the reagent, which promotes the separation of the calcium compound.]

[Note 6:  On ignition the calcium oxalate loses carbon dioxide and carbon monoxide, leaving calcium oxide: 

CaC_{2}O_{4}.H_{2}O —­> CaO + Co_{2} + Co + H_{2}O.

For small weights of the oxalate (0.6 gram or less) this reaction may be brought about in a platinum crucible at the highest temperature of a Tirrill burner, but it is well to ignite larger quantities than this over the blast lamp until the weight is constant.]

[Note 7:  The heat required to burn the filter, and that subsequently applied as described, will convert most of the calcium oxalate to calcium carbonate, which is changed to sulphate by the sulphuric acid.  The reactions involved are

CaC_{2}O_{4} —­> CaCO_{3} + Co,
CaCO_{3} + H_{2}so_{4} —­> CaSO_{4} + H_{2}O + Co_{2}.

If a porcelain crucible is employed for ignition, this conversion to sulphate is to be preferred, as a complete conversion to oxide is difficult to accomplish.]

[Note 8:  The determination of the calcium may be completed volumetrically by washing the calcium oxalate precipitate from the filter into dilute sulphuric acid, warming, and titrating the liberated oxalic acid with a standard solution of potassium permanganate as described on page 72.  When a considerable number of analyses are to be made, this procedure will save much of the time otherwise required for ignition and weighing.]