Cooking Carbonate Again

Previously, I've made some sodium carbonate for fun (even if the final experiments didn't work). This time, I need sodium carbonate for work.

The thing about my workplace is that they have minimal analytical reagents and equipment. An analytical balance that goes only to three decimal places. A pH meter with buffers. A few volumetric flasks. A burette*. There isn't even a single volumetric pipette in sight.

However, I want to carry out an acid-base titration**. The titrant and analyte are abundant in the lab, but I don't have any primary standards. Now what?

Possibly the most obvious solution would be to buy some standards. Things like KHP and sodium carbonate doesn't cost that much, even at analytical grade (I think). Unfortunately, it takes weeks for reagents to be delivered, and I don't have that much time.

In theory, of course, I can simply calculate the titrant concentration by measuring its pH. The accuracy and precision of the pH meter, however, leaves much to be desired, especially when the standard's pH will lie outside the usual calibration range of 4.01 to 10.00.

So here I am, cooking sodium carbonate so I can use it as a standard. Obviously, food grade reagents are not as pure as primary standards ought to be, but here's to hoping that they will be pure enough. I can't believe this, but I am actually nostalgic for a simple, KHP-NaOH-phenolphthalein titration. *Sigh*.

Which reminds me - the lab doesn't have any acid-base indicators except possibly phenolphthalein (which I'll have to borrow), either. At least there's the pH meter. *Another sigh*. I don't like monitoring titrations with pH because the amount of rinse water used during the titration will significantly influence results, whereas an indicator is somewhat immune to this dilution effect. The fact that a pH titration requires alot more readings doesn't help matters.

The lack of indicators in the titration makes me wonder, though. Usually, a carbonate-acid titration requires the reaction mixture to be boiled at indicator endpoint to expel CO₂ and to allow for a more accurate titration (otherwise the indicator changes colour prematurely due to dissolved carbonic acid). Given that I am monitoring the endpoint using pH, is there much point in me performing the boiling step? Will the inflection point be more obvious, or will there just be a little bump/spike on the pH versus volume curve right before the endpoint? (I am pretty sure the inflection point won't shift much - maybe by a drop of titrant or so.) This powerpoint suggests that yes, I ought to boil anyway, but laziness is a factor here, too...

Alternatively, what if I monitor the first endpoint with phenolphthalein instead? It's not a really good endpoint due to CO₂ absorption from the atmosphere, though. (Maybe if I perform the titration really fast...but that's optimism speaking here.)

The third alternative would be to do both - stick in both a pH probe and a few drops of phenolphthalein, and compare endpoints. The phenolphthalein will affect the pH a bit, but not much (ideally). The CO₂ absorption problem still applies, however.

And yes, I am well aware that my to-do list has things like "write essays" on it, not "toast baking soda" and "blog". I'll get back to work now...

*To it's credit, it's a *long* burette - 100 mL. Haven't seen that before. On the other hand, this means that I'll need a chair when taking readings. Dammit.
**It's a crappy way of determining pKa and molar mass (assuming monoproticity) of my analyte, but it's all I've got. Ideally, I want the temperature dependence of analyte pKa, but the van't Hoff equation needs information that I won't be able to get.