In need of a Chemists help...

[Shadow_Arbor]

I have a question for any chemists in the audience. Why doesn't KH booster send the pH of our aquarium into the stratosphere? I'm in my first year of a bachelors in biology, in a week I'll have  my final exam in basic and analytical  chemistry. Now from what I understand the KH we measure is a buffer solution when HCO3 and H2CO3 are present. Well this poses an issue. Any KH booster that is sold to re-buffer RO water states that it has HCO3- ions. I'm currently putting together my own GH and KH plus for RO water(will post on that once testing is complete) and I can't figure out why these KH additives don't send the pH sky high. I did the math for 21 ppm of HCO3 added to 10 litres of water. This raises the KH to about 1 degree and the pH to 9.1. I calculated this through hydrolysis, found the concentration of OH- and then calculated the pOH and reduced that number from 14 to find the pH. Assuming that my math is correct (I can upload a picture if needed) there has to be something that I'm missing. Either something I don't know yet or a misunderstanding of mine on how buffers work.

 

Edit: To clear something up here, I'm sure our aquariums have plenty of acids that help create this buffer. I am focusing on remineralized RO water which should have nothing in it other than the added minerals.

Thanks in advance for the help.

Edited January 10, 2022 by Shadow_Arbor

[CT_]

I don't have the fully story for you (I'm not a chemist, but I do play one on TV) but there is an equilibrium for pH, carbonate(s), and CO2.  Without knowing how much CO2 is absorbed into your water you can't compute pH.  Also, from what I can tell CO2 concentrations in an aquarium are rarely at equilibrium with the atmosphere, even with an airstone.

You can read about it here: https://en.wikipedia.org/wiki/Carbonic_acid#Chemical_equilibria

 

On 1/10/2022 at 10:06 AM, Shadow_Arbor said:

21 ppm of HCO3 added to 10 litres of water. This raises the KH to about 1 degree and the pH to 9.1.

I'd also recheck your math here.  21ppm is a concentration, so it can't be "added to" a volume of water.  The conversion form KH to ppm is 17.9ppm per KH unit, but that's in calcium carbonate equivalent units.  ie it's the ppm of calcium carbonate required to raise carbonate hardness by one degree KH, so if mixing units you can to use the molecular weight of calcium carbonate. 

[Shadow_Arbor]

On 1/10/2022 at 8:36 PM, CT_ said:

I don't have the fully story for you (I'm not a chemist, but I do play one on TV) but there is an equilibrium for pH, carbonate(s), and CO2.  Without knowing how much CO2 is absorbed into your water you can't compute pH.  Also, from what I can tell CO2 concentrations in an aquarium are rarely at equilibrium with the atmosphere, even with an airstone.

You can read about it here: https://en.wikipedia.org/wiki/Carbonic_acid#Chemical_equilibria

 

I'd also recheck your math here.  21ppm is a concentration, so it can't be "added to" a volume of water.  The conversion form KH to ppm is 17.9ppm per KH unit, but that's in calcium carbonate equivalent units.  ie it's the ppm of calcium carbonate required to raise carbonate hardness by one degree KH, so if mixing units you can to use the molecular weight of calcium carbonate. 

Thanks for your reply! What do you act in? Is it a movie or show that I may know? What else have you played in, that's fascinating!

You are correct about CO2 levels not being at equilibrium with the atmosphere. If I remember correctly  I read somewhere that about 2 ppm of CO2 dissolves  into the water. But, that concentration  is higher in distilled water. I actually asked my chemistry professor and he explained that the amount of dissolved carbon dioxide is irrelevant to the calculations as it's a minuscule amount. I calculated for 21 ppm i.e. 21 mg/l. because I needed to convert this into Molar so I multiplied 21 mg\l by 10 l which comes back as 210 mg or 0.21 gr. From there I divided 0.21 by the molecular weight which gave me the amount of moles. That was divided by 10, as I'm adding that amount of moles to ten litres to find the Molarity. From there all that's left is the hydrolysis reaction and -log(OH-).

 

I will state my issue a bit clearer. If I perform a 50% WC weekly with remineralized RO water. According to my math (which may be off) the pH of the remineralized water should be around 9. Next water change I will test this with a test kit. This just seems wild to me, as such a fluctuation in pH should be detrimental. The only theory I have is that our aquariums have enough acids in them that enough HCO3- immediately becomes H2CO3 creating a buffer solution. The problem is a 100% water change that I have performed to no visual ill-effect.

By the way, even if my math is off, adding the conjugated base(HCO3-) of a weak acid to pure water should bring the pH up rather drastically.

[CT_]

On 1/10/2022 at 1:17 PM, Shadow_Arbor said:

Thanks for your reply! What do you act in? Is it a movie or show that I may know? What else have you played in, that's fascinating!

HA.  I don't know if you're trolling me or not, but I'm just riffing on the "I'm not a doctor but I play one on TV" line 😉 .  What I meant is I'm a researcher at a university and use chemistry and collaborate with chemists so I know a little, but can easily be very wrong; trust me as much as a TV doctor :). 

 

What do your equations say for 0ppm carbonate?  ph 7?  distilled water left out has a pH of 6ish (iirc).

 

Also I wouldn't call 2ppm CO2 trivial.  without bothering to do all the math CO2 has a molecular weight less than HCO3- (by one O and one H) so 2ppm CO2 and 20ppm HCO3- is >10% co2/HCO3-.  eyeballing that Bjerrum plot i linked above that's a pH just under 7. 

 

You can also find precalculated tables of co2, kh, carbonate ppm (in calcium carbonate equivalents).  You can see the effect of co2 is non-trivial even down in the 1-4ppm co2 range.

 

Much of what you say surprised and bothered me at first too, so I've spent a lot of time reading, talking and asking questions too.  That doesn't make me right though, but hopefully if you find out something more or if/why I'm wrong you'll share it with me too :).

 

 

[mbwells]

@Shadow_Arbor I'm not a chemist either, but I am a microbial ecologist and have dealt with alkalinity *to a very limited extent*.  I have to second @CT_'s excellent reply that alkalinity refers to the concentration of CO2, bicarbonate, and carbonate in aqueous systems.  In anaerobic media for bacteria, bicarbonate and CO2 are common buffers for lower pH (a little less than neutral) media, whereas soda lake media (soda lakes are lakes where the salts present in the water come almost exclusively from bicarbonate and carbonate salts, and so they are extremely salty and quite alkaline) uses a bicarbonate and carbonate buffer to keep the pH at 9.5 or greater.  The bicarbonate and CO2 buffer is only possible in anaerobic systems because we can add a fixed concentration (usually 20%) in the atmosphere above the media.

Part of the trick is making sure you have the correct concentrations of CO2, bicarbonate, and carbonate salts to keep the pH of a solution at one of the pKa points of these carbon species.  The pKa points are points where the pH will remain quite stable, and so these are used to calculate the correct buffers to bacterial media so the pH doesn't change dramatically while the organisms are growing.

Mostly, I know that alkalinity is an incredibly difficult topic to learn well, and that it is only indirectly correlated with pH.  Here is a link to some powerpoint slides that may be useful for you.

 

https://www.soest.hawaii.edu/oceanography/courses/OCN623/Spring2012/CO2pH.pdf

[Shadow_Arbor]

On 1/11/2022 at 12:36 AM, CT_ said:

HA.  I don't know if you're trolling me or not, but I'm just riffing on the "I'm not a doctor but I play one on TV" line 😉 .  What I meant is I'm a researcher at a university and use chemistry and collaborate with chemists so I know a little, but can easily be very wrong; trust me as much as a TV doctor :). 

 

What do your equations say for 0ppm carbonate?  ph 7?  distilled water left out has a pH of 6ish (iirc).

 

Also I wouldn't call 2ppm CO2 trivial.  without bothering to do all the math CO2 has a molecular weight less than HCO3- (by one O and one H) so 2ppm CO2 and 20ppm HCO3- is >10% co2/HCO3-.  eyeballing that Bjerrum plot i linked above that's a pH just under 7. 

 

You can also find precalculated tables of co2, kh, carbonate ppm (in calcium carbonate equivalents).  You can see the effect of co2 is non-trivial even down in the 1-4ppm co2 range.

 

Much of what you say surprised and bothered me at first too, so I've spent a lot of time reading, talking and asking questions too.  That doesn't make me right though, but hopefully if you find out something more or if/why I'm wrong you'll share it with me too :).

 

 

Lol, not trolling. That sentence just went way over my head and I took it absolutely seriously! I will try and redo some of the math for 0 ppm carbonate and for 2 ppm CO2 as well as the HCO3. I will also try and get a more in-depth explanation from my professor. What you've written makes sense, and I a, sure that I'm missing something as our aquariums don't have a 9.1 pH for the most part.

Thanks for your help, ill update with my math and my professors answers when I get the chance!

On 1/11/2022 at 4:17 AM, mbwells said:

@Shadow_Arbor I'm not a chemist either, but I am a microbial ecologist and have dealt with alkalinity *to a very limited extent*.  I have to second @CT_'s excellent reply that alkalinity refers to the concentration of CO2, bicarbonate, and carbonate in aqueous systems.  In anaerobic media for bacteria, bicarbonate and CO2 are common buffers for lower pH (a little less than neutral) media, whereas soda lake media (soda lakes are lakes where the salts present in the water come almost exclusively from bicarbonate and carbonate salts, and so they are extremely salty and quite alkaline) uses a bicarbonate and carbonate buffer to keep the pH at 9.5 or greater.  The bicarbonate and CO2 buffer is only possible in anaerobic systems because we can add a fixed concentration (usually 20%) in the atmosphere above the media.

Part of the trick is making sure you have the correct concentrations of CO2, bicarbonate, and carbonate salts to keep the pH of a solution at one of the pKa points of these carbon species.  The pKa points are points where the pH will remain quite stable, and so these are used to calculate the correct buffers to bacterial media so the pH doesn't change dramatically while the organisms are growing.

Mostly, I know that alkalinity is an incredibly difficult topic to learn well, and that it is only indirectly correlated with pH.  Here is a link to some powerpoint slides that may be useful for you.

 

https://www.soest.hawaii.edu/oceanography/courses/OCN623/Spring2012/CO2pH.pdf

I will give those slides a look. What the both of you have said makes a lot of sense and I'm assuming there's just many things about buffers I have no clue on. Hopefully next semester Organic Chem will have some answers 😂

In the meantime I'll redo my math and look into what you both wrote.

Thanks to the both of you for sharing your experience @CT_ @mbwells

[Gideyon]

I'm not a chemist. Nor do I pretend to be one. 

 

But I'm curious... I know you're testing this with RO water. But is it in a "vacuum"?  ie nothing else in the aquarium or container you're doing these tests in? 

Just from my own experience... I raised KH with baking soda, and the pH did rise as well.   When my KH zero'd out, the pH crashed from 7.4 to lower than the test can go. 1/4 tsp of baking soda in the 10 gallon raised it up to the 7s again.   But the reason for the fluctuation was my substrate absorbing all the carbonate in the water.