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According to Wikipedia, the solubility of Potassium Dichromate increases very significantly with the temperature.

At 20°C, 13g/100ml, and at 100°C, 104g/100ml.

Therefore, it should be possible to dissolve about 520g of potassium dichromate in 500 ml of boiling water, but only about 65g in the same water, but at room temperature.

Today, I acquired about one kg of technical potassium dichromate and tried recrystallizing it, by weighing 507g of raw dichromate and dissolving it in about 500ml of water, room temperature. I expected it to no dissolve fully until the water was heated to boiling.

However, as soon as the water got slightly warm (45°C-50°C) everything dissolved. I added more raw material, and it kept dissolving until in the end everything (~970g) was added and dissolved before reaching boiling temperature.

Now, it is very slowly cooling down, and at about 35°C I seeded some small crystals and they're growing.

What went wrong in my calculations of the solubility? Anything I should've considered and missed?

Thanks.

PS.: Things I've already considered and ruled out:

  • Crystallization water: As far as I reseached, potassium dichromate doesn't have crystallization water, unlike, for example, copper (II) sulphate.
  • High humidity: My raw material was a free flowing powder, < 1mm particle size, uniformly deep orange colored. Had I high humidity, it'd clumping and not free flowing.
  • Contamination with potassium chromate, sodium chromate: Any of those wouldn't be soluble enough to explain what I observed.
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      $\begingroup$ This may be due to equilibrium with chromate, which is better soluble in low temp. If that's the case, you should acidify water. $\endgroup$ – Mithoron Feb 2 at 23:58
    • $\begingroup$ There are lots of images online of the color difference between dichromate and chromate solutions. Maybe you can tell from the color if that equilibrium was a big factor? $\endgroup$ – Andrew Feb 3 at 0:16
    • $\begingroup$ Nice idea. I will try acidifying the water, but I believe there was little to no chromate, since the color is a deep orange, not the strong yellow of chromates. Granted, the yellow may be masked by the orange. Collected crystals formed overnight (~25°C) and they weight at around 260g, still humid, guessing 230g dry. $\endgroup$ – Flyingfenix Feb 3 at 14:46
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    What happened was very simple.

    I've been duped.

    The seller sold me (intentionally or not) sodium dichromate and not potassium dichromate, or at least very contaminated with sodium dichromate.

    According to online source (https://en.wikipedia.org/wiki/Solubility_table#S), the solubility of sodium dichromate in water is about 198g/100ml of water at 30°C, and a very high 415g/100ml at 100°C.

    That explains clearly why everything dissolved promptly while the water was only slightly warm (40°C? Didn't come to mind put a thermometer in since I was expecting it to dissolve only near 100°C), since 500ml of solution would be able to dissolve 970g (my raw material) at only 30°C.

    Also, when I crystallized some of it overnight on a cold (~15°C?) water bath, the solubility would drop to about 175g/100ml, causing the precipitation of about 100g of salt.

    Finally, to settle matters about the presence of sodium dichromate, I dropped some saturated potassium chloride on a small sample of the solution in a test tube, and observed the instant formation of an orange precipitate.

    The same occurred when I picked some of the crystals I separated before and made a small amount of saturated solution - therefore the crystals were made of sodium chromate and not dichromate - since they have much higher solubility they shouldn't've precipitated at all before potassium dichromate.

    Also, I picked some potassium dichromate of known procedence and made some saturated solution and added a few drops of a saturated solution of potassium chloride, and nothing happened.

    In the end, everything said, for what I intend to make, sodium dichromate or potassium dichromate are interchangeable after adjusting quantities, but I'll have to execute extra steps to purify it from this solution.

    As an piece of advice (for me and everyone else), I'd test any new potassium/sodium dichromate from unknown sources before doing anything with it, simply by making some saturated solution and adding saturated potassium chloride.

    I'll have to devise a procedure for separating one from another given unknown proportions of either.

    Thanks for everyone time.

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    • $\begingroup$ RE: "... I'll have to devise a procedure for separating one from another given unknown proportions of either. ..." -- The best way would be with an ion exchange column. I'll add though that I did a lot of x-ray fluorescence. You never get just a K or Na salt alone. There is always a detectable amount of the other cation. $\endgroup$ – MaxW Feb 3 at 19:17
    • $\begingroup$ Sure, truly pure Na or K salts are hard to obtain, if their original production source involved both. For me, if it's > 99% of one or another, it is fine. Careful multiple recrystallizations are able to achieve that. Ion exchange columns are probably just outside of the reach of an amateur lab. $\endgroup$ – Flyingfenix Feb 5 at 13:26

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