Posted by Sheryl on Saturday, September 30, 2006 at 3:23pm.
Question: A student dissolved 0.30 g of a crude product in 10.5mL (the minimum amount required) ethanol at 25 degrees C. He cooled the solution in an ice-water bath for 15 minutes and obtained beautiful crystals. He filtered the crystals on a Hirsch funnel and rinsed them with about 2.0 mL of ice-cold ethanol. After drying, the weight of the crystals was found to be 0.015 g. Why was the recovery so low?
My answer: The dissolution occurred at nearly room temperature. The compound is too soluble in the solvent and most of it has been left in solution. A more appropriate solvent would be the one in which the compound is nearly insoluble at room temperature because the solubility-vs.-temperature curve is very steep.
Is this satisfactory?
Thanks from Sheryl
It may be ok but I think we are flying on a wing and a prayer with this one. First, there is nothing to indicate the mass of the crystals from the first pass. That would be nice to know. Second, the first pass was done with 10.5 mL ethanol and it states that is the minimum amount at 25 degrees C. (You note it might have been better to dissolve at an elevated temperature and that is true.) It was washed with only 2.0 cc OF ICE COLD ethanol. It is ice cold AND it is 1/5 as much as was used initially. So I wouldn't think 2.0 mL if ICE COLD ethanol would dissolve that much, especially if 10.5 mL of ice cold had crystals from it. IF the "beautiful crystals" from the first pass weighed only about 0.02 g or so then your answer holds up VERY WELL OR have you considered that 0.015 g may be the pure material with the remainder being another something. The problem states that it is 0.3 g CRUDE product so perhaps it just wasn't very good stuff.
Thanks a lot. I had to study that answer for a while.
Sheryl
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