# Why does the pH of this solution rise instead of fall when exposed to the atmosphere?

I have a couple of big barrels with a $50\,\mathrm{L}$ demi-water solution of the following nutrients:

$$\begin{array}{rl} 200~\mathrm{\mu M} & \ce{NH4NO3} & +\\ 30~\mathrm{\mu M} & \ce{KCl} & +\\ 10~\mathrm{\mu M} & \ce{CaCl2} & +\\ 10~\mathrm{\mu M} & \ce{KH2PO4} & +\\ \\ 10~\mathrm{\mu M} & \ce{Fe-EDTA} & +\\ \\ 0.7~\mathrm{\mu M} & \ce{ZnSO4} & +\\ 0.8~\mathrm{\mu M} & \ce{MnCl2} & +\\ 0.2~\mathrm{\mu M} & \ce{CuSO4} & +\\ 0.8~\mathrm{\mu M} & \ce{H3BO4} & +\\ 0.008~\mathrm{\mu M} & \ce{(NH4)6Mo7O24}\\ \end{array}$$

After mixing those nutrients, I adjust the $\mathrm{pH}$ of the barrels to $5.5$ by carefully adding $\ce{HCl}$ and continuing mixing until the $\mathrm{pH}$ stabilizes at $5.5$.

Then, when I measure the $\mathrm{pH}$in the barrels again after four days, I notice that it has risen to about $5.8$-$5.9$. This is contrary to my expectation. I'd have expected $\ce{CO2}$ from the air to be dissolved in the solution, thus lowering the $\mathrm{pH}$ rather than raising it.

Could someone explain to me what could be happening here?

• Could the barrel itself be reacting with the contents of the solution? I can see a small amount of metal/oxide corrosion happening at the expense of some hydrogen ions. Unbuffered solution pH is very sensitive to small variations of acid/base near pH 7, too. – Nicolau Saker Neto May 13 '15 at 15:31
• Another scenario that I didn't consider at the time is that I might not have mixed the barrels vigorously enough while adding the HCl or that I might not have given the mixture enough time to settle, and since I always stuck the pH meter in the upper layer… Also, later, this stopped happening so much, which is why I think the fault must have been with the observer (me). – BigSmoke Jul 14 '15 at 7:33

At that pH, you're well outside the buffering capacity of both phosphate and boric acid, so a $\mathrm{pH}$ rise from $5.5$ to $5.9$ corresponds to a net increase in base of about $2\ \mathrm{\mu M}$. That implies to me that it's likely something going on with one of the first five ingredients in your recipe, or the $\ce{HCl}$ added to adjust the $\mathrm{pH}$.

I can think of two scenarios:

1. Similar to Brinn's answer, perhaps a small amount of $\ce{HCl}$ has evaporated out of solution.
2. If your $\ce{Fe-EDTA}$ was added as $\ce{Fe(II)-EDTA}$, oxidation to $\ce{Fe(III)-EDTA}$ could explain the pH rise, by reducing water to hydrogen gas and hydroxide ion:

$2\ \ce{Fe(II)-EDTA} \rightarrow 2\ \ce{Fe(III)-EDTA} + 2\ e^-$

$2\ \ce{H2O} + 2\ e^- \rightarrow \ce{H2} + 2\ \ce{OH-}$

I don't know offhand how thermodynamically or kinetically feasible #2 is. With all of the other metals present in solution, though, it's not unreasonable to imagine that some sort of catalysis might be possible.

• Although I think the problem might have simply been in my impatient mixing technique, I'm accepting your answer, because it's beautiful. – BigSmoke Jul 14 '15 at 7:35

Unless the water to make the solution was freshly boiled then it was probably saturated with carbon dioxide to begin with. It's likely that ammonia is being lost to the atmosphere.

• Mmmm, not at pH 5.5, ammonia evaporation over four days should be negligible. Anyways, the pH would still be going down: ammonium would have to deprotonate before escaping as molecular ammonia. – hBy2Py May 13 '15 at 14:40