Why Do I Require More CaCO3 to Raise pH than Predicted?
If you can remember back to a discussion and inquiry we exchanged last week concerning the use of sodium carbonate vs. calcium carbonate. If you'll recall we were theorizing as to why differing amounts of carbonate in a titration experiment were needed to raise the pH of a solution from 4.0 to 5.0.

When we assumed 100% solubility of Na2CO3 and CaCO3, we should get virtually the same amount of carbonate to be used. However, in the titration experiment of last week, I got a result of 0.246 grams of Na2CO3 to raise 1 liter of water from pH 4.0 to 5.0 and 0.579 grams of CaCO3 to do the same.

Am I correct in saying that a titration is a measure or reflection of a compounds instantaneous solubility. In other words, Na2CO3 has a sol. of 30 g/100 ml and CaCO3 has a sol. of .001 g/100 ml, so calcium carbonate is many many times less soluble than sodium carbonate.

Over the past couple of days, I have repeated the titration experiement and continue to get results that suggest that far less Na2CO3 is needed vs. CaCO3 to raise the pH of the water from 4.0 to 5.0.

For your information, this water has moderate concentrations of acetic acid and aluminum sulfate which I'm sure impede the upward pH adjustment.

However, with all of that said, can you think of any possible theories as to why more CaCO3 is being used than Na2CO3??? If I left the initial CaCO3 titration continuously stirred for a 24 hour period to promote greater dissolving, wouldn't this gave an impact??

What am I missing? I cannot imagine why my titration is using more Na2CO3. Wouldn't it be reasonable to assume that all of the CaCO3 would dissolve in a 24-48 hour period?

I am at a loss and need your help to theorize some possibilities. Also, please note that even the titration for Na2CO3 is nowhere near the emprical amount of Na2CO3 required to change pH from 4 to 5. The titration was typically between 0.19 to 0.26 grams per liter Na2CO3 vs. a theroteical of .00477 g/l. Why is this?