I am set out to purify a batch of commercial oxalic acid with the aim in mind to end up with a highly pure product (starting from $\pu{4,000 ppm}$ total impurities and doing so with basic lab equipment and no advanced analytical capacity).

Looking at the distribution of solubility vs. temperature in aqueous medium (which - despite the widely discrepancy in data from different sources - seems to generally adopt a somewhat exponential shape), I was thinking that the best may be to carry out dissolution/filtration at high temperature, followed by fractional recrystallization.

A few questions remain though:

  1. I have no idea how I could discriminate against Vanadium impurities with the basic equipment I have at my disposal.
    According to the data I've found $\ce{V2O5}$ has a fairly low solubility at ~$\pu{10^{-1} g}/\pu{100ml}$ at $\pu{25 ^\circ C}$, a level which makes it much too soluble for proper filtration, and too little for the recystalization to efficiently remove it.
    I found data for the solubility of Vanadyl Oxalate to be of ~$\pu{22 g}/\pu{100ml}$ at $\pu{20 ^\circ C}$, and that of vanadium oxalate complexes of ~$\pu{14.5 g}/\pu{100ml}$ at $\pu{25 ^\circ C}$, which makes it very close to the solubility of oxalic acid itself (though, unlike for oxalic acid, I don't have data for how the solubility of Vanadium evolves with temperature).
  2. I had the notion that slowly recrystalizing the oxalic acid would yield better results, but I found an old paper* on its purification which found/states exactly the opposite, for a reason I don't very well understand. If someone had enough understanding to explain to me why a faster recrystalization would yield better results and/or to confirm/refute the paper's statement on the subject, I would greatly appreciate.
  3. I read upon a preliminary search here on Chemistry Stack Exchange that upon excessive heating, oxalic acid decomposes into $\ce{CO2 + CO + H2O}$ (in the ~$\pu{103-185 ^\circ C}$ range), thus I was wondering if it would be worth it to attempt to further exploit the exponential shape of its water solubility vs temperature by carrying out the dissolution/filtration in a $\ce{CO2}$ atmosphere, maybe under a slightly increased pressure. Would that significantly decrease the rate of its decomposition or simply complicate things without significant advantage?

  • 1
    $\begingroup$ Convert all oxalate salts to calcium oxalate first, which is extremely insoluble in water. Then dissolve in dilute $\ce{HCl}$ to get back oxalic acid form and recrystallize. $\endgroup$ Jun 15, 2020 at 19:19
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    $\begingroup$ The purity of 99.6% (based on my comment about buying) means $\frac{40}{100} \times 1000 = \pu{400 ppm}$ of impurity. How pure you want your oxalic acid? $\endgroup$ Jun 15, 2020 at 20:02
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    $\begingroup$ You can get $\ce{CaCl2}$ in high purity and all soluble in water. No need to purify. If you don't have equipment to check the purity, then how you are going to know how pure your crystals, any way? $\endgroup$ Jun 15, 2020 at 20:07
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    $\begingroup$ What Methew asked, and also: Why? (I´m sure there could be good reasons, no sweat, but what are they?) $\endgroup$
    – Karl
    Jun 15, 2020 at 20:36
  • 1
    $\begingroup$ Here it is 99% pure $\ce{CaCl2}$. $\endgroup$ Jun 15, 2020 at 21:22


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