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Hello fellow fish huggers. My well water is at 8.0 . Has anybody use a acid they recommend for lower the PH. My system hasnt been run yet as it is in its final stage but the water is sure is high for starting off. My system is 1000 gal so an economical option would be great.

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Adjust it using Hydrochloric acid... use with care...

If your water is highly buffered... you may well see the pH change.. only to see it bounce back again...

It will continue to do so until all the carbonate buffer is used up...

You water is a little high for starting out but not as drastic as you think.  If you can adjust it down into the 7.2-7.6 range you might find that cycling may take quicker.  When I started up my 300 gallon system I decided to adjust the pH before starting the fishless cycling and I just got a container of pH down like is used for spas or pools (sulfuric acid) And I believe the directions on the bottle will be easy to figure for a 1000 gallon tank.  It took me several doses to get the pH down and it tended to pop back up to 7.6 so I finally left it and allowed the cycling to finish bringing it down below 7.  And the bio-filter action will bring the pH down provided your media or some other material in the system doesn't buffer the system up.

I wouldn't recommend using the sulfuric acid long term or anything but initial pre cycling application doesn't seem to have messed anything up.  The pool acid was probably only $5 at the grocery store here and you are not going to use that much of it. Other options include muratic acid aquarium or pond pH down phosphoric acid (like from a hydroponics shop) vinegar might work but being organic the effect tends not to last so probably not the best choice citric acid has antibacterial properties so I wouldn't recommend that There are probably others but these are the ones I'm able to recall off the top of my head. Two acids not mentioned above that I personally work with is formic acid (HCOOH) and nitric acid (HNO3). Formic acid is a little weaker than the nitric, but both break down to or contain substances that I'm totally happy with in my system. Thank you Kobus. Where do you purchase your chemicals. Kobus Jooste said: Two acids not mentioned above that I personally work with is formic acid (HCOOH) and nitric acid (HNO3). Formic acid is a little weaker than the nitric, but both break down to or contain substances that I'm totally happy with in my system. I'm not from your part of the world, but the acids are easily available at any dedicated chemical supply store here in South Africa. They are widely used in industry and education, and I'm sure that you would be able to find some if you look for dealers that supply to those sectors. The only issue is to find people that will sell in small volumes here - typically they do barrels and I only need a few liters per year for more than one system. Thanks for the Tip Rupert. I started the system today and added fish so the fun begins. RupertofOZ said: Adjust it using Hydrochloric acid... use with care... If your water is highly buffered... you may well see the pH change.. only to see it bounce back again... It will continue to do so until all the carbonate buffer is used up... Thank you TCL I am still learning how to manuver this site. You info was a great help . TCLynx said: You water is a little high for starting out but not as drastic as you think. If you can adjust it down into the 7.2-7.6 range you might find that cycling may take quicker. When I started up my 300 gallon system I decided to adjust the pH before starting the fishless cycling and I just got a container of pH down like is used for spas or pools (sulfuric acid) And I believe the directions on the bottle will be easy to figure for a 1000 gallon tank. It took me several doses to get the pH down and it tended to pop back up to 7.6 so I finally left it and allowed the cycling to finish bringing it down below 7. And the bio-filter action will bring the pH down provided your media or some other material in the system doesn't buffer the system up. I wouldn't recommend using the sulfuric acid long term or anything but initial pre cycling application doesn't seem to have messed anything up. The pool acid was probably only$5 at the grocery store here and you are not going to use that much of it.

Other options include

muratic acid

aquarium or pond pH down

phosphoric acid (like from a hydroponics shop)

vinegar might work but being organic the effect tends not to last so probably not the best choice

citric acid has antibacterial properties so I wouldn't recommend that

There are probably others but these are the ones I'm able to recall off the top of my head.

My well water is 8.6 and I have to add some to my system every week.  Is there a buffering agent I can put in my tanks that will lower the PH.  Not just an acid.  A powder or something that will keep lowering it if the water is too high.
I'm about to start experimenting with silica sand / shell grit to see what it does to the system.  My pool shop told me not to dump used pool filter sand in the garden as it will acidify the soil - I'm interested to see if it will do anything to water though.  Will let you know.  In the past, I used to pre-balance alkaline pond water with nitric acid in a seperate drum before adding it to my AP system at a nursery.  Not sure if you want to have to keep on adjusting your top-up water though.

M Cosmo said:
My well water is 8.6 and I have to add some to my system every week.  Is there a buffering agent I can put in my tanks that will lower the PH.  Not just an acid.  A powder or something that will keep lowering it if the water is too high.

Uh, well I don't know of any "buffering agent" that would work opposite since buffers usually work to bring pH up if it drops below a certain point.

However there are some methods that have been used to acidify water that is too alkali.  I believe Travis Hughey has used pine needles soaking in the water to help acidify it.

Now there is one media I have heard of that might affect pH to the acidic side and that would be Diatomite which is sold under the trade name Maidenwell in OZ and I believe there is an American version out there too but I don't remember the name off hand.  It should be noted that if such media is allowed to bring the system pH as low as it will try to (well below 6) that there could be a dangerous mobilization of zinc and copper and I've heard of some people running into issues where they started a system with such media and were trying to grow trout and this did cause a problem.  They wound up needing to put large amounts of shell grit in the sump and had to remember to stir it often if the pH got too low.

Diatomite sounds very much to me like something that was sourced from diatomacious earth.  Diatomaceous earth, in turn, is a sedimentary rock high in silica sourced from deposition of diatom shells rich in silica.  It also contains aluminium and iron due to the fact that there is often clay particles mixed in with it.  Rock that has an acidifying effect on water almost always contain silica, typically in the form of quartz, such as sandstone.  Swimming pool filter sand is almost pure course quartz - derived crystals here in South Africa, which is why I originally suggested its potential use as a pH lowering agent.  I agree with TC Lynx that any pH lowering agent will potentially lower pH well below what Cosmo is looking for, thus I want to trial a small amount of it in water to see what it does.  Typically quartz is insoluble and I'm not sure what effect pool sand will have on a water sample over the long term, if anything at all.

I will be cautious of pine needles.  In many cases, it can be very acidic too.  Some South African pine plantations have a soil pH of 4 and nothing grows there other than pine trees.  All the soil microbes have been destroyed, resulting in pine needles just piling up and not decomposing.

Here is a different thought line to explore.  You can make a weak acid - Carbonic acid (H2CO3) by bubbling CO2 through your water.  Anyone try this before?  Not sure what the total effect of this process will be, as you are potentially loading the water with extra CO2 as well as potentially creating carbonates and bicarbonates.  A water chemistry expert will probably be needed to highlight the potential issues.

Carbonic acid

Not to be confused with carbolic acid, an antiquated name for phenol.
Carbonic acid is also an archaic name for carbon dioxide
Carbonic acid
Identifiers
CAS number 463-79-6 Y
ChemSpider 747 Y
KEGG C01353 Y
ChEMBL CHEMBL1161632 Y
Jmol-3D images Image 1
Properties
Molecular formula H2CO3
Molar mass 62.03 g/mol
Density 1.0 g/cm3 (dilute soln.)
Melting point

n/a

Solubility in water Exists only in solution
Acidity (pKa) 6.352 (pKa1)
Y(what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Carbonic acid is the inorganic compound with the formula H2CO3 (equivalently OC(OH)2). It is also a name sometimes given to solutions of carbon dioxide in water, because such solutions contain small amounts of H2CO3. Carbonic acid salts forms two kinds of salts, the carbonates and the bicarbonates. It is a weak acid.

Chemical equilibria

When dissolved in water, carbon dioxide exists in equilibrium with carbonic acid:

CO2 + H2O H2CO3

The hydration equilibrium constant at 25 °C is called Kh, which in the case of carbonic acid is [H2CO3]/[CO2] = 1.70×10−3: hence, the majority of the carbon dioxide is not converted into carbonic acid, remaining as CO2 molecules. In the absence of a catalyst, the equilibrium is reached quite slowly. The rate constants are 0.039 s−1 for the forward reaction (CO2 + H2O → H2CO3) and 23 s−1 for the reverse reaction (H2CO3 → CO2 + H2O). Carbonic acid is used in the making of soft drinks, inexpensive and artificially carbonated sparkling wines, and other bubbly drinks. The addition of two equivalents of water to CO2 would give orthocarbonic acid, C(OH)4, which is unimportant in aqueous solution.

Addition of base to an excess of carbonic acid gives bicarbonate. With excess base, carbonic acid reacts to give carbonate salts.

Role of carbonic acid in blood

Carbonic acid is an intermediate step in the transport of CO2 out of the body via respiratory gas exchange. The hydration reaction of CO2 is generally very slow in the absence of a catalyst, but red blood cells contain carbonic anhydrase, which both increases the reaction rate and dissociates a hydrogen ion (H+) from the resulting carbonic acid, leaving bicarbonate (HCO3-) dissolved in the blood plasma. This catalysed reaction is reversed in the lungs, where it converts the bicarbonate back into CO2 and allows it to be expelled. This equilibration plays an important role as a buffer in mammalian blood.[1]

Role of carbonic acid in ocean chemistry

The oceans of the world have absorbed almost half of the CO2 emitted by humans from the burning of fossil fuels.[2]  The extra dissolved carbon dioxide has caused the ocean's average surface pH to shift by about 0.1 unit from pre-industrial levels.[3] This process is known as ocean acidification.

Acidity of carbonic acid

Carbonic acid is diprotic: it has two protons, which may dissociate from the parent molecule. Thus there are two dissociation constants, the first one for the dissociation into the bicarbonate (also called hydrogen carbonate) ion HCO3:

H2CO3 HCO3 + H+
Ka1 = 4.45×10−7 ; pKa1 = 6.352 at 25 °C.

With a pKa1 of 6.352, carbonic acid H2CO3 is almost 10x weaker acid than acetic acid.

The second for the dissociation of the bicarbonate ion into the carbonate ion CO32−:

HCO3 CO32− + H+
Ka2 = 4.69×10−11 ; pKa2 = 10.329 at 25 °C and Ionic Strength = 0.0.

Care must be taken when quoting and using the first dissociation constant of carbonic acid. In aqueous solution carbonic acid only exists in equilibrium with carbon dioxide, and the concentration of H2CO3 is much lower than the dissolved CO2 concentration. Since it is not possible to distinguish between H2CO3 and dissolved CO2 (referred to as CO2(aq)) by conventional methods, H2CO3* is used to represent the two species when writing the aqueous chemical equilibrium equation. The equation may be rewritten as follows (cf. sulfurous acid):

H2CO3* HCO3 + H+
Ka = 4.6×10−7(General Chemistry: An Integrated Approach Third Edition); pKa = 6.352 at 25 °C and Ionic Strength = 0.0.(NIST CRITICAL Database)

Whereas this pKa is quoted as the dissociation constant of carbonic acid, it is ambiguous: it might better be referred to as the acidity constant of dissolved carbon dioxide, as it is particularly useful for calculating the pH of CO2-containing solutions.

pH and composition of carbonic acid solutions

At a given temperature, the composition of a pure carbonic acid solution (or of a pure CO2 solution) is completely determined by the partial pressure $\scriptstyle p_{CO_2}$ of carbon dioxide above the solution. To calculate this composition, account must be taken of the above equilibria between the three different carbonate forms (H2CO3, HCO3 and CO32−) as well as of the hydration equilibrium between dissolved CO2 and H2CO3 with constant $\scriptstyle K_h=\frac{[H_2CO_3]}{[CO_2]}$ (see above) and of the following equilibrium between the dissolved CO2 and the gaseous CO2 above the solution:

CO2(gas) CO2(dissolved) with $\scriptstyle \frac{[CO_2]}{p_{CO_2}}=\frac{1}{k_\mathrm{H}}$ where kH=29.76 atm/(mol/L) at 25°C (Henry constant)

The corresponding equilibrium equations together with the $\scriptstyle[H^+][OH^-]=10^{-14}$ relation and the charge neutrality condition $\scriptstyle[H^+]=[OH^-]+[HCO_3^-]+2[CO_3^{2-}]$ result in six equations for the six unknowns [CO2], [H2CO3], [H+], [OH], [HCO3] and [CO32−], showing that the composition of the solution is fully determined by $\scriptstyle p_{CO_2}$. The equation obtained for [H+] is a cubic whose numerical solution yields the following values for the pH and the different species concentrations:

$\scriptstyle p_{CO_2}$
(atm)
pH [CO2]
(mol/L)
[H2CO3]
(mol/L)
[HCO3]
(mol/L)
[CO32−]
(mol/L)
1.0 × 10−8 7.00 3.36 × 10−10 5.71 × 10−13 1.42 × 1009 7.90 × 10−13
1.0 × 10−7 6.94 3.36 × 1009 5.71 × 10−12 5.90 × 1009 1.90 × 10−12
1.0 × 10−6 6.81 3.36 × 1008 5.71 × 10−11 9.16 × 1008 3.30 × 10−11
1.0 × 10−5 6.42 3.36 × 1007 5.71 × 1009 3.78 × 1007 4.53 × 10−11
1.0 × 10−4 5.92 3.36 × 1006 5.71 × 1009 1.19 × 1006 5.57 × 10−11
3.5 × 10−4 5.65 1.18 × 1005 2.00 × 1008 2.23 × 1006 5.60 × 10−11
1.0 × 10−3 5.42 3.36 × 1005 5.71 × 1008 3.78 × 1006 5.61 × 10−11
1.0 × 10−2 4.92 3.36 × 1004 5.71 × 1007 1.19 × 1005 5.61 × 10−11
1.0 × 10−1 4.42 3.36 × 1003 5.71 × 1006 3.78 × 1005 5.61 × 10−11
1.0 × 10+0 3.92 3.36 × 1002 5.71 × 1005 1.20 × 1004 5.61 × 10−11
2.5 × 10+0 3.72 8.40 × 1002 1.43 × 1004 1.89 × 1004 5.61 × 10−11
1.0 × 10+1 3.42 3.36 × 1001 5.71 × 1004 3.78 × 1004 5.61 × 10−11
• We see that in the total range of pressure, the pH is always largely lower than pKa2 so that the CO32− concentration is always negligible with respect to HCO3 concentration. In fact CO32− plays no quantitative role in the present calculation (see remark below).
• For vanishing $\scriptstyle p_{CO_2}$, the pH is close to the one of pure water (pH = 7) and the dissolved carbon is essentially in the HCO3 form.
• For normal atmospheric conditions ($\scriptstyle p_{CO_2}=3.5\times 10^{-4}$ atm), we get a slightly acid solution (pH = 5.7) and the dissolved carbon is now essentially in the CO2 form. From this pressure on, [OH] becomes also negligible so that the ionized part of the solution is now an equimolar mixture of H+ and HCO3.
• For a CO2 pressure typical of the one in soda drink bottles ($\scriptstyle p_{CO_2}$ ~ 2.5 atm), we get a relatively acid medium (pH = 3.7) with a high concentration of dissolved CO2. These features contribute to the sour and sparkling taste of these drinks.
• Between 2.5 and 10 atm, the pH crosses the pKa1 value (3.60) giving a dominant H2CO3 concentration (with respect to HCO3) at high pressures.