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I'm treating tap water to prepare fresh water suitable for the aquarium. I receive water with $\mathrm{pH}~7.6$ from a dechlorinator and I need to lower ammonia level.

For every unit decrease in $\mathrm{pH}$ the amount of unionized ammonium will decrease 10 times, so I need to lower the $\mathrm{pH}$ to at least $6.5$ to reduce the content of ammonia $(\mathrm{p}K_\mathrm{a}(\ce{NH3}) = 9.24)$ down to

$$\frac{[\ce{NH3}]}{[\ce{NH4+}]} = 10^{\mathrm{pH} - \mathrm{p}K_\mathrm{a}} = 10^{6.5 - 9.24} \approx 0.002\quad\text{or}\quad0.2\,\%.$$

I used citric acid to lower the $\mathrm{pH}$ down to $6,$ but after a couple of hours $\mathrm{pH}$ raised again to $7.$ The $\mathrm{pH}$ was measured using a $\mathrm{pH}$ meter prior to the contact with fish.

What is the reason for this raise? Citric acid overcomes waters buffering feature then raise the $\ce{H+}$ concentration in water. Why would $\mathrm{pH}$ raise again?

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  • $\begingroup$ The first part alone is hard to believe. Tap water is not much of a buffer, it's nearly impossible to get it to pH = 6 and not further. $\endgroup$ Dec 24, 2019 at 13:26
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    $\begingroup$ @Ivan Neretin In the Moravian carst area, the majority of water hardness is the bicarbonate hardness. CO2/HCO3- forms a diluted buffer with maximum capacity at pH=pKa1*=6.83, if I remember correctly. $\endgroup$
    – Poutnik
    Dec 24, 2019 at 13:56
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    $\begingroup$ It is possible citric acid ( in any form ) is being consumed by microbial activity, what would increase pH. $\endgroup$
    – Poutnik
    Oct 16, 2021 at 9:51
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    $\begingroup$ Also, note that rainwater in equilibrium with atmospheric $\ce{CO2}$ - assuming aqueous speciation of $10^{-3.5}$ mol / L and given a Henry's Law constant of $10^{-1.5}$ mol / (L atm) - clocks in with a pH of 5.65. This, btw, famously led researchers at Princeton University towards the (correct) "discovery" of acid rain for the wrong reasons (it's already acidic; the contributions of NOx and SOx were determined later). $\endgroup$ Oct 16, 2021 at 16:23

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If the water contains significant amount of bicarbonates, part of released carbon dioxide may escape and original bicarbonate is finally replace by the dihydrogen or hydrogen citrate.

$$\ce{ HCO3- + H3Citr -> H2Citr- + H2O + CO2 ^}$$ $$\ce{2 HCO3- + H3Citr -> HCitr^2- + 3 H2O + 2 CO2 ^}$$

So, the final $\mathrm{pH}$ may not differ much from the original one, compared to the temporary one.

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    $\begingroup$ So if I want to lower it how can I do it? Is there any other ways to lower pH and keep low? $\endgroup$ Dec 24, 2019 at 18:39
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    $\begingroup$ Try adding a citrate buffer Na2HCitr/Na3Citr. pKa3 = 6.39, You can prepare it by citric acid solution neutralisation to pH 6. $\endgroup$
    – Poutnik
    Dec 24, 2019 at 18:52
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Maybe citric acid have had a negligible effect on the $\mathrm{pH}$. Everybody knows that in contact with air, pure water is carbonated by $\ce{CO2}$ from the atmosphere. And the $\mathrm{pH}$ may go down to $5.5$. Later on, the water may loose its $\ce{CO2}$, and the $\mathrm{pH}$ goes back to $7$.

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Oxygenated tap water is rich in transition metals including Fe and Mn ions. Citric is a source of H+ and a good chelate and can drive a redox reaction in the presence of oxygen and H+ proceeding as follows: $$\ce{4 Fe^2+/Mn^2+ + O2 + 2 H+ --> 4 Fe^3+/Mn^3+ + 2 OH-}$$

There is also a likely equilibrium reactions that can be effective in recycling ions (to continue the reaction, as occurs in natural waters) in the presence of citrate acting a chelate for, say, ferric:

$$\ce{Fe^2+ + Mn^3+ <=> Fe^3+ + Mn^2+}$$ $$\ce{Cu+ + Fe^3+ <=> Fe^2+ + Cu^2+}$$

Note: The above are active charge equilibrium reactions, referred to as a metal redox couple, and support the role of transition metals in their lower valence states engaging in notable radical creation formation via Fenton and Fenton-type (non-iron case) reactions. Further, in the presence of lab or sunlight, some further recycling of metal ions is possible.

Reference: See my comments and sources cited here.

The reaction with oxygen is electrochemical in nature with an inception period and proceeds with time.

Note: $\ce{H+}$ is consumed, so the $\mathrm{pH}$ would be expected to rise.

Here is also an electrochemistry reference, in particular, Table 2, where the first listed half-cell reaction corresponds to the above after adding two $\ce{H+}$ to both sides (creating water as a product). However, as this reaction can produce basic salts, as I have discussed previously, I believe it is more informative. Further, I do prefer my rendition, as per my prior link supplied above, it could also display a possible radical chemistry underpinning to the underlying half-cell reaction mechanics.

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    $\begingroup$ I doubt chelation would take a couple of hours, in homogeneous aqueous solution it should be much faster. Also, I tried to follow you cited sources, but it appears like you are quoting yourself quoting Wikipedia on copper(I) oxidation, which has a questionable relation to this topic, so it would be nice if you duplicate your sources here, both for clarity and to protect the post from link rot. $\endgroup$
    – andselisk
    Dec 24, 2019 at 15:02
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    $\begingroup$ There is not enough metal in water. Even traces of metallic ions make the taste of water quite bad. This is not the correct reason. We have to confirm if the author is using citric acid indeed or not! $\endgroup$
    – AChem
    Dec 24, 2019 at 18:21
  • $\begingroup$ The cyclic nature of the chemistry (as noted in my metal redox couple comment and possible role of light) can overcome a seemingly small amount of transition metal presence with time. My argument also accounts for a reported pH rise in the presence of air. An alternate explanation concerning a relatively sudden loss (and not, say a gain, or an established equilibrium situation?) in the water's CO2 content, is in my opinion, not likely, as the pH of tap water is relatively stable. $\endgroup$
    – AJKOER
    Sep 21, 2020 at 4:41

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