In short, I think this is possible. I have never seen reference to this in the literature but one explanation I was told in grad school was that as you increase the temperature of an aqueous solution, the pH will decrease because the OH- ions have a greater vapor pressure than H+. I always have assumed that was true, but never really saw reference to that in any other literature, and I never really looked either. Along those lines it may be difficult to measure what you are looking for because there may be some temperature dependent effects of pH occurring.

Hello Kathryn, I regularly use carbonate buffers in the lab and can confirm that they become a lot more alkaline over time. I often observe a very significant pH change from 8.0 to 8.5/9.0 within a week or so. I never tried it, but base on these numbers it should be possible to estimate how much of the carbonate evaporates as CO2. Regarding acidification by the absorption of CO2 from the air, I have noticed that double-distilled water quickly acidifies in the lab. However, pure water has no buffering capacity, so I don't expect cultures that are grown in 5% CO2 will change in pH dramatically.

Thank you both for these answers. I agree that the temperature dependence would complicate calculations. It also has ramifications when trying to measure the pH with a meter, because as soon as media is taken out of the incubator it begins to cool as well as re-equilibriate with the lower CO2 concentration in the outside air.

In regards to acidification of culture medium in the incubator, I have found that our cultures do acidify under 5% CO2. Perhaps this is indicating (no pun intended) that I am not using enough buffer.

I found the observation that media can change so quickly in a week very interesting. These experiments I'm taking about were spaced three days apart so I guess there might have been large pH differences between the first and last.

You're halfway there, Kathryn. I'm going to guess you're using NaHCO3 as the complimentary base in your system. You got it right that HCO3 can pick up an acid to become H2CO3, which decomposes to H2O and CO2. But what about the other half of the reaction? Na does not sit alone.

NaHCO3 + H2O NaOH + H2CO3 then H2CO3 --> H2O + CO2(g)

You're out gassing the CO2, thus driving the reaction that is forming NaOH, a much stronger base. How to prevent this? You can't, but you can slow it down by keeping the caps on bottles tight, trapping the CO2 in the bottle. The reaction I wrote above is in equilibrium, so adding 5% CO2 causes acidification as more carbonic acid builds up in the media and the NaOH is nuetralized, leaving excess carbonic acid.

Hope this explanation helps.

Hi Charles,

Thank you very much for pointing this out. I am indeed using NaHCO3 as the bicarbonate source. I hadn't considered the salt + H20 part of the reaction.

So I guess the main reason media goes more alkaline in the fridge is because of the NaOH, not because protons are being mopped up by HCO3- ?

I'm sorry but would you mind explaining the last sentence your last paragraph with different wording... It probably makes sense as is but I don't quite understand! :P

My interpretation of what you have said is that you will just make more carbonic acid (or HCO3- + H+ when the proton dissociates) in the incubator. But wont the amount that is made depend on how much HCO3- is there in the first place?

Do you mind if I just ask a few direct questions?

(1) All equations/equilibriums considered, over time, do you think the bicarbonate concentration for stored media would go down? I've been thinking about this and I'm not sure anymore because if the cap is on tight its not like the air in the bottle is an infinite sink for CO2 so the bicarbonate wont just keep evolving off as CO2 right? Its a closed system so should come to equilibrium at some point.

(2) pH is in the incubator is related to HCO3- concentration by a modified version of the henderson-hasselbalch equation:

pH = 6.1 + log( [HCO3-]/(kH*PCO2) )

where kH is Henry's constant for CO2 dissolving into liquid (at 37degreesC) and PCO2 is the partial pressure of CO2 in the incubator.

When one uses 'aged' media in the incubator in which the HCO3- concentration has decreased over time, would this not imply that pH would be lower than for media made up fresh?

(3) Are you saying that it doesn't matter that NaOH was built up whilst the media is in the fridge because it just all gets neutralised once you put it in 5% CO2?

Sorry for bombarding everyone with long posts with lots of questions!

Thanks to everyone who has answered so far :)

Kathryn

Guus B Erkens Dear Dr Erkens, I was wondering how you deal with the problem of increasing alkalinity of your bicarbonate buffers? Besides making fresh buffers just before use, is there any way to preserve the pH of the buffers? What's the longest you keep a solution before the pH change is obvious? 3 days? I would be deeply grateful for your advice.