# Lowering KH via Lime & Muriatic Acid



## JoeRoun (Dec 21, 2009)

*I AM IN NO SUGGESTING OR RECOMMENDING THIS TO ANYONE!*


This is in response to a question asked in another thread, people are curious why folks have so much difficulty altering the pH of their water with acid.

Indeed it is Alkalinity of which carbonate hardness KH, is a subset that resists (buffers) acid lowering pH. 



When sufficient acid “overcomes” the KH, the pH drops.
The problem is that unless the carbonate is actually removed the “buffering” will return and the pH will return to its previous value.
 
*Poor man’s alternative to RO as well as maintaining the general hardness (GH).*

First is liming, many of us have seen this when dosing fertz containing Calcium, in fact it is the one point where KH and GH actually cross. 



I have read many times about people saying it “snows” when they dose fertz. I resist saying anything as many “experts” like more complicated answers. 



It is really Calcium bonding with the carbonate and precipitating, usually Calcium carbonate. 



Effectively reducing the KH and not adding to the general hardness.
Industry uses this method to protect equipment, the
Calcium carbonate settles or is filtered out.
It is not difficult to reduce the KH in tap water using lime and maybe some heat.
 
The second method that I really do not recommend, but use myself, :icon_biggMuriatic acid, HCl.

Follow all the safety rules, serious do this out of doors, in a well-ventilated area, breeze blowing away from you. Use gloves, safety goggles, away from kids, pets so on. If you are a kid, get parental permission and supervision. The stuff can ruin your day (and life)!

I really recommend either being able to determine carbonate hardness accurately either by titration (some idea of chemistry) or a Lamotte test kit, still titration but the stuff is worked out for you. Be careful of even Hach’s and others have end points above pH 5. Make sure you have a reasonably accurate method of measuring pH.

The problem of pH returning, actually it is the carbonates reforming as with the liming above, we removed the carbonates from solution.


In this case we are going to drive the CO2 off the carbonate, effectively removing the carbonate from solution by aeration. 

I would never dose Muriatic acid, HCl into my tanks directly, my critters are to curious and even well diluted I would fear the potential damage, damage that might not be immediately apparent.

Which brings the issue of pH swings, if your tank has critters, keep the


 pH changes to no more than pH 0.3 per day.
In fact if you are changing the pH more than 1 whole point,
remember that a full point lower is 10 times more acidic
 
allow a couple days between the 0.3 drops in pH,
take your time.
 
Fish can be harmed and you may not see the results of the harm for a couple of weeks.
Large changes in pH can irreversibly harm fish, 


the classic example are Neon tetras experiencing a full pH point swing,
even if the water is immediately stabilized and the fish keeper thinks he got away with it,
they all look fine, swimming around, eating but two-weeks later the tetras begin dying off and
within 3-days, a week, they are all dead.
 
 
For this procedure I recommend using a tub, garbage can (preferably new, at least clean:icon_wink), it is nice to have an air pump, capable of really getting the water frothing, moving. Glass air diffusers are better than stone or wood for this (my opinion anyway).

Determine how many dKH you wish to remove, frankly in my tub I usually aim for 1-dKH, figuring I will mix tap water to get the desired d-KH. Honestly there are situations to go under 4-dKH. Generally too much KH is not a problem, there are definite exceptions though.

Now, I have certainly learned that any mention of “moles” of something, “specific gravity,” “normalizing,” “molarity,” and so forth are a big turn off. So I am ripping off Dr. Randy Holmes-Farley.

Dr. Randy Holmes-Farley’s method seems awkward, but I get it. He uses Muriatic acid, HCl straight out of the bottle and refers to it as 11,000 meq/L, so he is seems to be asserting normal Muriatic acid is the 32%, 20-*Baumé variety. 
*

*Be aware *Muriatic acid is sold in varying concentrations so adjust accordingly.

Thinking of “acidity as negative Alkalinity” and since 1-meq/L will drop the KH 2.8-dKH

Figuring the water out of my tap is 12-dKH and 



I wish to drop it to say 4-dKH, means
12=dKH – 4-dKH = 8-dKH.
That is I want to reduce 8-dKH, or
8-dKH ÷ 2.8 meq/dKH = 2.9 and
I operate in 20-gallon batches (my preference)
 So: 



(1-dKH/11,000-meq) × (1-meq Muriatic acid) × (2.9-dKH) × (3.8-L/1-gal) × 20-gal × (1000-ml/1-L)
= (1 ÷ 11,000) × (2.9 Muriatic acid) × (3.8) × 20 × 1000-ml
= 20-ml Muriatic acid
 


Add about half, 10-ml Muriatic acid in this case.
Aerate for a while, check the pH to see if it looks right.
Assuming things seem correct
add the remaining 10-ml Muriatic acid
continue aerating until you have blown off the CO2
once the pH has stabilized
(remember it will be rising as the CO2 dissipates)
 
change out enough water to lower the pH 0.3
 
It is actually less complicated than it sounds.

Joe
FBTB


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## FlyingHellFish (Nov 5, 2011)

Very informative, thanks. 


So if a person wanted lower KH, they just add a mild acid? I'm guessing the end result is a lower KH with GH being the same? 

Would adding more Magnesium sulfate (Epsom salt) help as well? Doesn't Magnesium sulfate keep the calcium level from raising too much?


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## JoeRoun (Dec 21, 2009)

*Degas*



FlyingHellFish said:


> Very informative, thanks.
> So if a person wanted lower KH, they just add a mild acid? I'm guessing the end result is a lower KH with GH being the same?
> Would adding more Magnesium sulfate (Epsom salt) help as well? Doesn't Magnesium sulfate keep the calcium level from raising too much?


You are welcome.

Well there is nothing mild about Muriatic Acid, but I get your meaning.
Yes, pretty much any acid, for instance you could use sulfuric acid, then aerate to blow off (degas as the gentry would say) the CO2.


You are correct this will not change the general hardness.
Adding Magnesium sulfate really shouldn’t do anything except add to the GH. In a sense the biggest problem with adding extra Epson salt is the laxative effect on the critters.


Technically any of the metal ions can precipitate out with carbonates; Iron (II) that isn’t chelated or chelated with glucose or even EDTA at somewhat higher pH would be an example.
Joe
FBTB


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## Zorfox (Jun 24, 2012)

Nice write up Joe. 

When using muriatic acid, is the chlorine a problem? 

Liming I get. It’s been around for a long time and still used today. Btw, now you’re showing your age :icon_wink.

What I don’t understand was why dosing fertilizers causes’ calcium and carbonates to precipitate. We aren’t adding Ca so what is causing the reaction? Forgive my ignorance. I simply don’t get it.

Also, thank you for mentioning Dr. Randy Holmes-Farley. I’ve never heard of him. I searched for pertinent information about him and wow, what a wealth of knowledge!


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## JoeRoun (Dec 21, 2009)

*Anions, Cationsand Polarity... Oh My...*

Thank-you.

Actually the chlorine really has no effect and the aeration tends to blow it off as well.

The “Calcium” we add is not the elemental Calcium (Ca) but an ionized form Ca⁺⁺ or Ca⁺⁺. The whole Ca⁺⁺ thing is why pure elemental Calcium is relatively rare. We dose Calcium compounds. 

As an example when we dose Calcium chloride, which is really a compound made of three atoms, one Calcium ion, Ca⁺⁺ a “cation” with a formal charge of 2, and two chloride (ionized form of Chlorine) “anions each with a formal charge of 1 (negative), Cl⁻.

Because water is “polar” the Calcium chloride CaCl₂ is broken apart and now the Ca⁺⁺ is looking for another stronger bond and in our case manages to find (mainly due to the pH of our systems) a couple of Hydrogen carbonate anions (HCO₃⁻). 

I am thinking of trying a bit better explanation but getting into equilibria of Hydrogen, carbonates, bicarbonates, and primary, secondary and tertiary phosphate ions is a bit much for now.

This is part of the complexity we face when trying to explain separate situations, the “snow” people see could be Calcium phosphate or any number of combinations of things. 

I understand why people want simple, definitive answers, even if they are wrong and many retailers and manufacturers are more than happy to sustain the misinformation, but ultimately, facts need to win and facts do not care what we think or prefer.

I think Dr. Randy Holmes-Farley is one of the brighter bulbs and perhaps, along with the fact that reefers are always operating at the edge of the envelope, there is less tolerance for misstatement of fact in the saltwater tank world.

The Skeptical Aquarist is one of those places I always check before posting anything technical (controversial here) just to make sure I am not too far off.

Respectfully,
Joe
FBTB


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## Zorfox (Jun 24, 2012)

I agree that the BB code be enlarged here. However, you can use HTML Entity (decimal) for unicode. Here is a good site that list them all. There is a HUGE number of them so I suggest writing your own list rather than searching every time. 

Here is calcium with two superscript positive signs

Ca⁺⁺

This is what you have to type to see them (only in preview btw). Just remove the parentheses. I had to add those so you could see the code. 

Ca&#(8314);&#(8314);

Example of an equation...

H₂O + Co₂ ⇔ H₂Co₃ ⇔ HCo₃⁻ + H⁺

versus this...

H2O + Co2 <--> H2Co3 <--> HCo3- + H+ 

and the code for the equation...

H&#(8322);O + Co&#(8322); &#(8660); H&#(8322);Co&#(8323); &#(8660); HCo&#(8323);&#(8315); + H&#(8314);

Edit: Here are a few to start your list. Remember to remove the parentheses 

*Subscript numbers...*

*C₀* C&#(8320);
*C₁* C&#(8321);
*C₂* C&#(8322);
*C₃* C&#(8323);
*C₄* C&#(8324);
*C₅* C&#(8325);
*C₆* C&#(8326);
*C₇* C&#(8327); 
*C₈* C&#(8328);
*C₉* C&#(8329);

*Superscript numbers...*

*C⁰* C⁰
The HTML entity does not work for 1-3 here. You can use a keyboard shortcut though.
C¹ <-- Alt + 0185
C² <-- Alt + 253
C³ <-- Alt + 0179
*C⁴* C&#(8308);
*C⁵* C&#(8309);
*C⁶* C&#(8310);
*C⁷* C&#(8311);
*C⁸* C&#(8312);
*C⁹* C&#(8313);

*Superscript charges...*
C⁺ C&#(8314);
C⁻ C&#(8315);


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## Zapins (Jan 7, 2006)

I'm not so sure that the Cl in HCl (muriatic acid) on its own will blow off with aeration. Cl2 (gas) can blow off, but as far as I know ionic Cl- will not because it is in the aqueous form. To blow off you'd need to remove an electron from one Cl- allowing the two to bind together to form a neutral gas.

If Cl- blew off then you'd expect table salt (NaCl) to blow off when mixed with water, but we all know that when you let a cup of salt water dry out you get the same amount of NaCl salt back afterwards (because the Cl- never left the water).

Will the Cl- convert to Cl2 by interacting with the carbonates in the water (the buffering system)?

Cl- can cause harm to plants and animals if it gets too high. I'd estimate a few hundred ppm, maybe a few thousand ppm Cl- is harmful depending on the species, so I suppose you have quite a bit of wiggle room.


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## JoeRoun (Dec 21, 2009)

*I know the answer is in here, somewhere...*

Hi Zapins, All,

Good catch, when answering Zorfox, what passes for my brain was thinking the chlorine disinfectant in the water, completely missed the point.:redface:

I know I have an answer to the point about the chloride that however will require a dangerous journey to the inner reaches of my dilapidated brain.

Though just on the net amounts of chloride the dosing levels to remove or reduce 8-dKH will add about 82-ppm Cl⁻.

Respectfully,
Joe
FBTB

Bump:


Zorfox said:


> I agree that the BB code be enlarged here. However, you can use HTML Entity (decimal) for unicode. Here is a good site that list them all. There is a HUGE number of them so I suggest writing your own list rather than searching every time.
> 
> Here is calcium with two superscript positive signs
> 
> Ca⁺⁺


 Hi Zorfox, 

Very cool, I actually think I know how to do that via MS Word, I will do some messing about.

I am sure if the swells that run this joint would just give up some of their jet-setting ways the BB and the environment would be much better off.:smile:

Respectfully,
Joe
FBTB


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## Zapins (Jan 7, 2006)

JoeRoun said:


> Though just on the net amounts of chloride the dosing levels to remove or reduce 8-dKH will add about 82-ppm


Very cool. I'll keep it in mind. This is the first chemical method to remove KH that I've read about. 

Have you tested it out yet?

82 ppm Cl shouldn't cause problems.

Bump: 82 ppm Cl shouldn't cause problems with plants, so we are safe on that front.


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## JoeRoun (Dec 21, 2009)

*Poor Mans RO... Sort-of...*

Hi Zapins, All, 

Yes, I have actually been doing this for 6 or 7 years, I rather hate to recommend it to general populations due to the potential for misuse or rather people who are not careful.

I think it is a great alternative to RO, especially when I wish to keep the other minerals.

I will say I have never considered anything under 140-ppm Cl⁻ maybe even higher particularly problematic. In fact I like to keep 40-ppm Cl⁻ minimum primarily for fish health, not sure that is anything I can prove, but it is what I do.

Respectfully,
Joe
FBTB


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## plantbrain (Dec 15, 2003)

Zapins said:


> I'm not so sure that the Cl in HCl (muriatic acid) on its own will blow off with aeration. Cl2 (gas) can blow off, but as far as I know ionic Cl- will not because it is in the aqueous form. To blow off you'd need to remove an electron from one Cl- allowing the two to bind together to form a neutral gas.
> 
> If Cl- blew off then you'd expect table salt (NaCl) to blow off when mixed with water, but we all know that when you let a cup of salt water dry out you get the same amount of NaCl salt back afterwards (because the Cl- never left the water).
> 
> ...


This is correct, you cannot evaporate/degas Chloride ions in water. Chlorine gas, yes. I had about 100 ppm or thereabouts, some folks have been using KCL at very high levels, a client of mine uses KCL to soften the Ca/Mg out of their entire house water, so the water is very briny coming in, most plants are find though as long as I add the Ca and Mg back. Chloride is about 150ppm.


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## JoeRoun (Dec 21, 2009)

*Aeration is the Key*



Zapins said:


> Very cool. I'll keep it in mind. This is the first chemical method to remove KH that I've read about.
> Have you tested it out yet?
> 82 ppm Cl shouldn't cause problems.
> Bump: 82 ppm Cl shouldn't cause problems with plants, so we are safe on that front.


 Hi Zapins, All, 

I might add one of the reason Dr. Randy Holmes-Farley’s formula confused me a bit is the assertion that Muriatic acid is 11,000 meq/L out of the bottle. Years ago I thought I recalled the Muriatic acid I had was 20%. The stuff I use now is a 37% I get through a lab supply outfit. I tend to calculate the stuff based on its v/v molarity about 12M (12.06M actual) for my stuff as opposed to about 10.2M (I originally wrote 9.6M) for the 32% Muriatic acid referenced. 

Everyone needs to remember the numbers I gave are based on the “normal” 32% (10.2M (I originally wrote 9.6M)) Muriatic acid.

I am doing it now with sulfuric acid, I am using some drain cleaner that appears to be about 96% sulfuric acid (18M), buying it as drain cleaner at the big box store is a heck of lot less expensive than buying from my lab supply guys, it will be tomorrow at the earliest before I have the results but there is no reason any strong acid shouldn’t work to reduce KH as long as it is well aerated.

Respectfully,
Joe


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## plantbrain (Dec 15, 2003)

JoeRoun said:


> Hi Zapins, All,
> 
> Yes, I have actually been doing this for 6 or 7 years, I rather hate to recommend it to general populations due to the potential for misuse or rather people who are not careful.
> 
> ...


Joe, you recall the old toys that used vinegar and baking soda to cause them to move or the volcano experiments kids do/did? 

Acetic acid is cheap and widely available and safe at the 5% level(white distilled). Baking soda is basically the KH in water.

So adding Vinegar would be a simple and safer method, and not that much difference in cost.

CH3COOH

Breaks down the KH in tap water also like HCL.

High M HCL and H2SO4 are both nasty stuff. Vinegar, something you eat and put on salad dressing. Both will do the same job. 
I use to use H2SO4 at high M for cleaning diatoms, hated it, still do.


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## Zorfox (Jun 24, 2012)

Sorry for the confusion. I actually did mean chloride. Never the less question answered. Thanks.

I personally use white vinegar for mixing fertilizers. It's not only cheap but I usually have it handy. I honestly don't bother trying to change my PH anymore. It's too much of a hassle. If I can't grow something in my tap water it just isn't in my tank. Makes life much easier.

I still don't understand the "snow" issue. I understand the various forms of calcium. My question is how does adding fertilizers (KNO3, KH2PO4 and Plantex) cause this reaction. We're not adding calcium (any form) or carbonates. It's my understanding that it's the iron and phosphates reacting which causes this. If I'm wrong I really want to understand why.


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## JoeRoun (Dec 21, 2009)

*Worth It Or Not, Not for Me to Say*

Hi Tom, 

Certainly acetic acid will work, if I recall the objection to vinegar was not the acetic acid but the other organic materials that came along. 

I think my post was inspired by the fact that a fair number of folks seem confused by the reason Muriatic acid and sulfuric acid in particular did not accomplish their goals. 

My sense is that .89M acetic acid is more trouble than I wish to take. The final answer is without regard to which acid the answer is still aeration. 

Worth the effort or not that is an individual decision, to me it is, for the tanks I desire a lower KH, since I moved from a place that had 12-ppm elk pee & 3-ppm bear pee, I have just had certain plants and fish that do not get along with what passes for water here.
************
Tom on a separate note, you ought to advise your clients to get a new water softener. Under ideal conditions they should be getting 76-ppm K⁺ for every 100-ppm of Ca⁺⁺ and Mg⁺⁺ removed and 0-ppm Cl⁻. All of the chloride should be out with the backwash.

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Hmmm...*

Hi Zorfox, 

I will do a bit somewhere on ionic bonds and double bonds in a little bit, I have it mainly written I am just dealing with a little exhaustion and lawyers, which are doubly exhausting.

For now let us assume Tom’s clients water softener were working properly.

Most of what goes on in a water softener is like, your “snow” as a result of dosing fertilizers. An exchange happens a precipitate is formed, one bond is broken, and another is made. 

In the case of the water softener, the resins allow the Sodium chloride or Potassium chloride to be held until the strength of the water combined with the greater attractiveness of the Magnesium or the Calcium releases a Potassium or Sodium (depending on KCl or NaCl) cation and bonds with the now free chloride anion, which still remains effectively attached to the resin. The resin in this case is acting as the filter to remove the undesirable hardness.

Then in a despicable act of environmental cruelty, the resin is backwashed, the brine released to be dealt with by sewage treatment folk, and the process repeated.

Respectfully,
Joe
FBTB


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## Okedokey (Sep 2, 2014)

We have to factor in the dissociation properties of acids not just the pH (e.g. pK). Carbonic acid in coke makes it have a pH of 1, but its a weak acid (pKa of 6.7). HCL is much stronger and acetic acid is weaker again. Does anyone have an opinion of what factor dissociation has here for the objective? I don't think its an issue with acetic acid, as it will be over 90% dissociated at a pH over 5.

e.g. nitric acid is a weak acid (pKa of -1.3), and the conjugate base is nitrate. Anyone considered using this instead?


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## JoeRoun (Dec 21, 2009)

*Nice Try... Do You Want Me to Give the Answer?*

Hi, 

Actually we do not have to factor a dissociation properties. As a practical matter they are irrelevant, particularly the convenient method Dr. Randy Holmes-Farley’s formula provide. I love his “negative Alkalinity.”

Given that in no example are we attempting to overcome the acid-neutralizing capacity of solutes, in this case we are calculating an approximate amount of acid, using a molar-equivalent (meq) per liter to reduce a definite acid-neutralizing capacity (ANC). Then as a matter of procedure we use half or less of the acid to determine if we are on track.


For the heck of it you should not have too much trouble calculating it; all the relevant information is available.


Respectfully,
Joe
FBTB




Okedokey said:


> We have to factor in the dissociation properties of acids not just the pH (e.g. pK). Carbonic acid in coke makes it have a pH of 1, but its a weak acid (pKa of 6.7). HCL is much stronger and acetic acid is weaker again. Does anyone have an opinion of what factor dissociation has here for the objective? I don't think its an issue with acetic acid, as it will be over 90% dissociated at a pH over 5.
> 
> e.g. nitric acid is a weak acid (pKa of -1.3), and the conjugate base is nitrate. Anyone considered using this instead?


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## Okedokey (Sep 2, 2014)

His work is focused on saltwater environments where hardness is usually a lot higher correct? Because of this, we need to factor in pK to our considerations in a freshwater environ, as the pH changes due to adding acids will be temporary and in the case of some acids reduce the dissolved oxygen dangerously.


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## JoeRoun (Dec 21, 2009)

*Phosphoric acids maybe*

Hi, 

In fact depleting oxygen would be one of the reasons for the aerating and never adding the acid directly to the tank, saltwater or freshwater. The greater reason is the chance of harming critters.

We can factor the pK, but it does not change the outcome when the reactions involve 20-ml of 9.6M HCl into 75-liters of water that will then be added to at least another 75-liters (assuming 50% water change).

The case of nitric or phosphoric acids, reducing certain buffering materials by less than an amount to significantly decrease pH could in a marginal system cause lowering, perhaps even significant lowering of oxygen, but given that most aquarium systems provide for transportation of water and introduction of air, gas exchange, so for an operating aquarium supporting life, I would see this as a minor concern.

This would be analogous to a lake or body of water undergoing eutrophication, low flow stagnation introduction of farm wastes or perhaps even “acid rain,” such an equivalent aquarium would in my opinion benefit from any water change as certainly the act of changing water would remove the organic materials that tend to add to greatly to the acid-neutralizing capacity, among the reason in the field we need to filter the water prior to testing for carbonate hardness or Alkalinity.

Among the problems I have seen folks having problems with their aquariums is having to get them in good enough shape that hobbyist test kits can be useful.

Respectfully,
Joe
FBTB


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## Okedokey (Sep 2, 2014)

Hi Joe

As I understand it, for what you're saying to be true regarding pK, the log of the conjugate base and acid need to be zero, and this is very very rare, and certainly not the case with HCL.

That is, for the same molar value of different acids you will have different pH (i.e. H3O+) values due to different pK (e.g. ionisation of the H) .

To be fair its been many years since I last studied this...


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## JoeRoun (Dec 21, 2009)

*Full Disclousure: I Own No Mason Jar Stock*

Hi, 

I am not trying to be contentious, I think I understand your point of view and I think it is a correct view for “natural” settings. I seriously doubt we would find much disagreement on protection of our natural environments, protection of resources. 


Our tanks are not natural,
even so called natural systems are anything but.

 
At the same time I am aware of lakes and ponds, even a swamp of my childhood, brought back to life by pumps and aeration, stunning and dramatic what a little cleanup and a little water movement can accomplish. 

These are the very things that we do in our tanks and really there is nothing wrong with deciding to alter the chemistry, any more than it is to landscape or add mechanical or biological filtration, though it is best to think these things through. 

I am shocked by the things people will do with a gorgeous display tank, they have invested stunning amounts of time and money, then with no particular thought make a significant change in something they do not understand and they never bothered to test.

Rather than listening to “Guru’s;” think, study, experiment before doing, why does anybody think “they” sell mason jars anyway…

Respectfully,
Joe
FBTB


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## Okedokey (Sep 2, 2014)

Thanks for the reply Joe, but I wasn't under any idea that our tanks are representative of natural systems, nor under any illusion that pK has much to do with conservation. I simply am referring to the fact that pK and pH are linked beyond the meq of an acid and important in soft, freshwater remediation considerations. They're directly linked via the logarithmic ratio of the conjugates (salt and acid); therefore pH and pK are almost never the same, regardless of the molar value.

Anyway... I disagree that pK doesn't matter, but for the purposes of this post, we can agree to disagree.


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## JoeRoun (Dec 21, 2009)

*Conjugate them Acids and Bases*

Hi, 

Full confession it is not like I remember all this stuff from the 8th grade, I went to the local community college and took a couple of chemistry courses. Now I have taken a couple more university courses, because I am able too and realized I did not understand this stuff. I try not to plagiarize, but obviously I didn't think any of this stuff up on my own. For this I looked it up to make sure I was correct.

I hope my explanations make sense. 

Let us start with a rule from Bronsted acid theory that an acid-base reaction is the transfer of Hydrogen cation (or proton) H⁺. Water reacts with itself transferring a Hydrogen cation (or proton) from one to another to form H3O⁺ and OH⁻. Accordingly an acid is a “proton donor” and of course a base is a “proton acceptor.”

We tend to divide acids into weak and strong: Ka = [H3O⁺][A⁻] / [HA]
The Ka of HCl is 1 X 10³ as opposed to a weak acid such as acetic acid Ka = 1.8 X 10⁻⁵, compared to water a very weak acid Ka=1.8 X 10⁻¹⁶.

Remember that “p” is an operator that means “-log” and “[]” is a quantity, an absolute value.
pH = -log [H₃O⁺] and POH = -log[OH⁻] Now here comes the problem with pKa, it makes the strong look weak. 
pKa = -log Ka, therefore HCl: pKa= -3, CH₃CO₂H: pKa = 4.7 and H₂O: pKa = 15.7

Since we are into Bronsted we _must_ recognize each acid has a conjugate base and vice-versa. 

Since magnitude Ka, measures strength of a conjugate acid; Kb is the strength of a conjugate base.
If we multiply the Ka for an acid (HA) by the Kb for a conjugate base (A⁻):
([H₃O⁺][A⁻] / [HA]) X ([HA][OH⁻] / [A⁻]) = [H₃O⁺][OH⁻], we can replace this equation with equilibrium constants; KaKb=Kw=1 X 10⁻¹⁴
The product KaKb is a small number, either the acid or the conjugate base can be strong, but if one is strong the other is weak.

Now that all of that is clear as mud, we throw in water; water has a limiting effect on acids and bases, but that is for another day.

Note here that HCl, Ka is 1 X 10⁻⁶, its pKa is -6, while Cl⁻, Kb is 1 X 10⁻²⁰ and pKb is 20.

Where if anywhere do we wish to go with this?:icon_eek:
:smile:
Respectfully,
Joe
FBTB


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## Okedokey (Sep 2, 2014)

You've just regurgetated some text. The application to this thread is the type of acid, ph and pK matter for the purposes of a stable KH/PH remediation. I understand the theory well...


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## JoeRoun (Dec 21, 2009)

*I Am Glad You Do!*



Okedokey said:


> You've just regurgetated some text. The application to this thread is the type of acid, ph and pK matter for the purposes of a stable KH/PH remediation. I understand the theory well...


 Hi,

Yeah I had to really look this one up, funny how confident I was about knowing it when I started… Oh well… :redface:



What bothers me the most is it almost makes sense.


Respectfully,
Joe
FBTB


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## Okedokey (Sep 2, 2014)

Thanks for the discussion Joe


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## 141130 (Jan 25, 2014)

*jumping over from another post (as you said)*

Okay. Thanks for your patience. I definitely have some issues resolved. But I still don't understand other things.



> The problem of pH returning, actually it is the carbonates reforming as with the liming above, we removed the carbonates from solution.
> 
> In this case we are going to drive the CO2 off the carbonate, effectively removing the carbonate from solution by aeration.


If we do not aerate off the CO2 then the carbonates reform. Okay. But presumably different ones from the original otherwise the original sulfuric acid would just lie around at the bottom of the tank waiting for a chance to react at a later point when the wind blows or something. How does it work in reality? Thanks.

And the final question I asked I'll just quote for other readers.



> By adding sulfuric acid are we just substituting the carbonic acid that would normally be produced by natural CO2 (keeping the pH down)? If so what are the consequences? I mean, if you had high pH as a problem, why couldn't you just set up an automated pH regulator to drip feed a very dilute concentration of sulfuric acid (or other acid) into a filtering system to lower the pH and just do regular measured water changes, more or less as normal, from the tap to replace the conjugate bases converted to acids over a given week? Then you wouldn't need a big old CO2 canister in your sitting room.


Regards,


----------



## JoeRoun (Dec 21, 2009)

*No Carbonic Acid For Now*

Hi David, All,

For the moment I need you to accept the idea that in chemistry, physics and mathematics usually any action or event can be run in both directions and we observe this everywhere with everything except time.

For now I wish to set the carbonic acid (H₂CO₃), thing aside that is a hot button issue and for the time being we will simply accept that enough CO₂ can produce enough carbonic acid, HCO₃⁻ to alter pH, but the relationship is different than reaction of carbonates with metals.​ For the moment at least accept that all strong acids act the same in water and will act the same as equal solutions of H₃O⁻ (for the purists, assume 1M or more concentration, we will also stipulate this implies very weak acids cannot act as acids in water). _Here we are not interested in the full range of action._

I will use words since I do not know how to do the HTML stuff Zorfox so graciously attempted to teach me. 
​
We have talked about hydrochloric, sulfuric and nitric acids so I will use these and kind of covers the range of interests. 



Also in freshwater and in our aquariums the “big 7” ions are Ca⁺⁺, Mg⁺⁺, K⁺, Na⁺ (the cations),HCO₃⁻, Cl⁻ and SO4⁻⁻ (the anions). A bit different proportions in our tank than in the “real world,” but close. Since we are in water we will have H⁺ and OH-, due to the dissociation of water. 



During the following reactions the Hydrogen atoms are replaced by metal ions forming the salt:


Metal carbonate + hydrochloric acid, gives us; metal chloride, plus water, plus carbon dioxide.
 

Metal carbonate + sulfuric acid gives us; metal sulfate, plus water, plus carbon dioxide.
 

Metal carbonate + nitric acid gives us; metal nitrate, plus water, plus carbon dioxide.
 
In each case the resultant metal carbonate is ‘aqueous’, the subscript ‘(aq)' should appear that is to say in solution. Figuring our nitric acid found a Calcium carbonate, we now have CaNO₃ for our plants or for the eutrophication of our lakes from run off. 



As a note for those learning chemistry all nitrates are soluble. The water is of course liquid, subscript “(l)” behind it and the carbon dioxide is a gas, subscript “(g)” behind it.

This is the point in the process where driving off the carbon dioxide, will permanently reduce the carbonate hardness.​ Now a post or three ago _Okeydokey insisted_ on talking (wisely perhaps) about the acid dissociation constant, I added the base dissociation constant (just cuz), these are special cases of the equilibrium constant of a chemical reaction (_reaction quotient when the reaction has reached equilibrium_).

Over time if all the parties are available the reactions will reverse themselves. This can be complex since in a situation where a number of reactions have taken place the swaps ultimately are governed by availability and the various base disassociation and equilibrium constants and just as acid disassociation constants governed the reaction based on strength, the base and equilibrium constants are going to govern the move back.

I think I will stop here and see what confusion this caused.:hihi:

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Let It Snow Let It Snow*

Hi Zorfox, All,

I do not know if you noticed in the “No Carbonic Acid for Now” post replying to David’s question about where the acid goes, was actually a bit of a preview answer to your “snow” question.

That is indeed the answer to how adding KNO₃, KH₂PO₄ and Plantex can make it snow.

The big 7 ions are already there along with waters little 2.

It is the equilibrium constants that provide the rules, 

in your example the most likely precipitate (or snow) is going to be Calcium phosphate (Ca(H₂(PO₄)₂), it could be as Monocalcium phosphate somewhat soluble. 

If you are seeing it, it is most likely Dicalcium phosphate, generally considered insoluble. If you work at it you can get about 20-mg into a 100-grams of water. 

Dicalcium phosphate ends up on the substrate where it tends to eventually dissolve, though depending on a number of factors break up and bond with your substrate material especially if it is clay. 

Somewhere I read that someone say they had Tricalcium phosphate, seriously insoluble, maybe 2-mg per 100-grams of water. My understanding though is the only way to get this, is to burn bone.

The Plantex has a number of things that can precipitate, though I think Magnesium would be the main one. Most of the other interesting stuff is chelated. 

I suppose at higher pH the iron (II) might. 

Gluconated Iron (II) at even mildly base water might also be subject to precipitating out of solution.

Good old Potassium nitrate is the least likely to participate in such rude behavior.

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

> Over time if all the parties are available the reactions will reverse themselves. This can be complex since in a situation where a number of reactions have taken place the swaps ultimately are governed by availability and the various base disassociation and equilibrium constants and just as acid disassociation constants governed the reaction based on strength, the base and equilibrium constants are going to govern the move back.


acid + metal carbonate → a salt + water + carbon dioxide

So the bonds of the newly produced salt break because the CO2 would react with the water to produce an acid and metal carbonate. I understand the theory but I'm confused because it doesn't match up to what I have experienced.

When I stick acid in the pH drops. I mix but don't de-gas. It bounces back. I add more the next day. I mix but don't de gas. It bounces back a little but not all the way. I add some more. I mix but don't de-gas. It stabalises. From what you're telling me the pH should never drop because I'm not de-gasing. But it does drop and it does stabalise. And if you're right, then where has all that HCl gone? Is it just sitting, un-reacted, at the bottom of the tank? Surely it would react just as it did when I added the stuff for the first time?

I can remember doing this stuff with acid in high school: stirring in an acid and blowing over the beaker to complete the reaction. So when i make my solution there must be some CO2 blowing away naturally. Over time all of the acid will react (having the CO2 blown away naturally) and complete. Presumably, in the meantime, the acid exists in an un-reacted dilute form in the water (occasionally bonding with a metal carbonate every now and then and hoping for some wind to blow its way and complete the reaction...but why wouldn't the pH go up and down and all over the place for that, say, three-day period, where the reaction is taking place)?

Is the blowing a kind of a "catalyst" (in the non-technical sense)?


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## JoeRoun (Dec 21, 2009)

*No Catalyst Here*

Hi David,

First I want to say to everyone we tend to get hung up on CO₂ and or carbonic acid, because it is something we focus on for our plants. In this case CO₂ just happens to be the “_thing_” we can most easily remove to lower the KH.

I must issue a Guru alert, if your Guru tells you otherwise and you are true believer that is okay, I will not argue any point.​ Sticking acid into the pH drops is a great example of seeing for yourself.
As a side note what you are doing is called titration, this is exactly what those KH test kits are, an acid (usually dilute sulfuric acid), and a pH indicator (I think most use at least two; to give a clear end-point), you then add the acid a drop at a time, shaking or upending the container (test tube) a couple of times to give the acid, base reaction time to occur and as it happens, help get rid of the gas (CO₂) in this case.​ What you have done in your experiment is allow much of the CO₂ remain, also it is important to remember that as water come to equilibrium, depending on the density altitude (partial pressure, issue) water, exposed to the air will end up with about 3-ppm CO₂.

Now this is the part that seems to get everyone, the CO₂ is released from its carbonate bonds as a *gas*, this means it is going to start obeying a set of rules that apply to gases as well as of course all the other laws of chemistry and physics, but this is a transition, a _phase_ change. 

Now comes a problem and opportunity (kind of go hand-in-hand), while the reaction rates (haven’t talked about this yet) for the “acid + metal carbonate → a salt + water + carbon dioxide” part was relatively quick, CO₂ is a slow actor.

Now back to where has all the “acid” gone, let’s go back to a basic definition from Bronsted acid theory that an “_acid-base reaction is the transfer of Hydrogen cation (or proton) H⁺_," in the case of sulfuric acid it gave up a proton (_protonated_ is a term you may hear) and the sulfate ended up as part of the salt, the acid exists no more (kind of sad, isn’t it:icon_sad. ​ That continues until such time as the “acid” no longer has a base resisting, or “taking up” the protons, or if prefer the H⁺ cation. Remember it is _*Hydrogen*_, there is isn’t much there to start. 

When they had you “blow” across the top of the test tube that indeed helped remove the gas, technically blowing across the top of the test tube lowered the pressure over the liquid, thereby reducing the “partial pressure,” squeezing the gases out and moved excess gas along.

Shaking, stirring, aerating, pouring, all add kinetic energy, this is the energy of motion, or possessed by an object _due to its motion_, and quickens the rate of reaction while quickening the gas exchanges, while providing uniformity in the solution.

Heating can add to this and be kinetic, in our situation we really are not interested in heating’s addition to the kinetic energy. 


In our situation adding heat does two things for us
first it reduces the amount of gas that can stay in solution.
the second point is an oddity in some of our planted tank favorites and that is that contrary to most compounds;
temporary hardness compounds become less soluble as the temperature rises.

 
The reason liming hot water is more efficient, the reason that scale builds up in your hot water heater.

The answer is that nothing we have done has used any “catalyst” in any sense of the word.:smile:

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

*thanks for your reply*

I need to see some equations for this to make sense. It's not as simple as I thought. :icon_smil There's a whole soup of reactions taking place. :confused1:

acid + metal carbonate → a salt + water + carbon dioxide

And if I don't aerate, does this happen?

a salt + water + carbon dioxide → acid + metal carbonate

If so what acids are formed? The original acid? This will be the complicated bit, right. I'd love some equations. It makes more sense than talking. I hope. Actually, not so sure. I was just trying to work it out but for myself but how on earth do chemists remember how many electrons are available for exchange in a given reaction with a given element (do they just keep looking at the periodic table all day?).


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## JoeRoun (Dec 21, 2009)

*A=[hco3-]+2[co3--]+[b(oh)4]+[oh-]+2[po4---]+*

Hi David,

No, the acid is not going to reform, essentially the “acid”…:confused1:

Let me try it this way something is only an acid if it has a proton (H⁺ cation) to donate and a base only exists if it can accept a proton (H⁺ cation), this is by definition.

In this solution you are seeing the result of “buffering,” more base absorbing the protons and the carbonates reform, at our planted tank pH range it will be overwhelmingly bicarbonates. 

Actually keeping track of charges is why we use the superscripts and subscripts.


A=[HCO₃⁻]+2[CO₃²⁻]+[B(OH)₄]+[OH⁻]+2[PO₄-³⁻]+[HPO₄²⁻]+[SiO(OH)₃⁻]-[H⁺]-[HSÒ⁻]


Subscript "T" s have been omitted.


Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

JoeRoun said:


> Were you to be using sulfuric, hydrochloric or any other strong acid to lower KH (buffering) to lower the pH, you would need to strongly aerate, to blow-off (degas, the gentry would say) the C02 to keep the carbonates/bicarbonates (at our pH’s, nearly all bicarbonates), since the chemical reactions move in both directions.


So if I add acid and aerate my pH drops but my KH remains the same? I don't understand what you mean when you say "move in both directions" because you just said that the acid is gone and cannot return -- so what takes away the ion from the salt to create the original carbonates/bicarbonates (and then what happens to that "other thing"?)?


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## JoeRoun (Dec 21, 2009)

*The Ups Are Down*

Hi David,

If you are aerating the solution after the acid is added the carbonate hardness is reduced, it does not remain the same.

A base assumption we are making is that the Alkalinity (the buffering) is overwhelmingly bicarbonate/carbonate. Since the strong acid followed by aeration is designed to remove (change) the bicarbonate/carbonate portion by driving of the CO₂ component.

As long as KH is the buffer adding acid will reduce KH even if you are not to measure the pH change. Avoiding being too technical, the pH has changed just not enough for you to measure, remember adding the hydrogen ion is a reduction in pH; pH is the number of Hydrogen ions (H⁺). 

Part of our confusion in all of this is the _more_ H⁺ the _lower_ the pH.

My suspicion is that you are using water that has something other than (in addition too) KH as the buffer. Whether it is phosphate, acetates, borates or take-your-choice.

Too confirm this for yourself; do the same tests you have been doing using baking soda and distilled water, make up a 10-dKH solution and try the same thing.
:icon_smil
Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

*super thanks*

And if I *don't* aerate then the carbonates/bicarbonates will reform (and so the KH reading will return to what it was originally??) -- and so the original acid (sulfuric acid) has (in this case, *eventually*) turned into sulfur (minus an electron), which would be phosphorus (is that right?)?

If I add sulfuric acid and don't aerate I'm basically dosing phosphorus into my tank!? Which would surely produce an algae bloom (and they sell it as algae free stuff...and don't mention aeration).

Can you show me the equation for the carbonates reforming?

acid + metal carbonate → a salt + water + carbon dioxide

And then...


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## JoeRoun (Dec 21, 2009)

*No PP for You...*



david765 said:


> And if I *don't* aerate then the carbonates/bicarbonates will reform (and so the KH reading will return to what it was originally??) -- and so the original acid (sulfuric acid) has (in this case, *eventually*) turned into sulfur (minus an electron), which would be phosphorus (is that right?)?
> 
> If I add sulfuric acid and don't aerate I'm basically dosing phosphorus into my tank!? Which would surely produce an algae bloom (and they sell it as algae free stuff...and don't mention aeration).
> 
> ...


 Hi David,

No there is no phosphorus.

You end up with a metal sulfate, carbon dioxide and water. In your case it is probably going to be Calcium sulfate, Carbon dioxide and water, those two Hydrogen ions are going to look for and find a lonely Oxygen. 

H₂SO₄ + CaCO₃ ⇒ CaSO₄ + CO₂ + _H₂_O

We always must keep in mind this is taking place in an aqueous solution.

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Sorry, I Missed What You Were Clearly Saying*

Hi David,

My apologies David, you had asked the question clearly enough earlier, what-passes-for-my-brain, simply did not process it.

Generally when there is not an excess of H⁺ that is when there are still bicarbonates/carbonates available, the H⁺ end up as water. One of the tricks is in general usage to remember the big 7 and little 2 ions.

My background is in Physics (you know, real science:hihi, the problem with Chemistry is there are so many things which must simply be memorized, or in my case looked up each time.:confused1:

Respectfully
Joe
FBTB


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## 141130 (Jan 25, 2014)

Thanks for your help. It's been a trip down memory lane for me.

It's funny you mention that physics is your thing -- I nearly studied physics at university. I always remember chemistry being harder simply because, as you said, there was so much stuff to memorise (it expanded on the logic of physics to make that mess we call the picnic table...sorry periodic! table).

In physics it was all narrowed down to a few formulas and then it was just maths all the way down. I can remember studying imaginary numbers and the core of quantum mechanics at sixteen or seventeen but now I can't even understand how stupid I've become. :hihi: And if I don't like chemistry let's just agree not to even talk about biology! Yuck. :hihi:


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## JoeRoun (Dec 21, 2009)

*Acetatea*

Hi,

I figured I would post to report on my trials with sulfuric acid as Zapins and David765 had asked about and tried with vinegar as Tom Barr had recommended.

I think the sulfuric acid is actually much better than the Muriatic acid, HCl, in fact I am rather shocked at how much more pleasant the task is. Muriatic acid, HCl is really kind of nasty to deal with. 

The big thing I noticed beyond the lack of fumes was that it is better to use dilute the sulfuric acid, I started with an 18M (call it 96%), but it is much easier to work with 6M or 3M acid since it is best to work the KH levels down in stages, not to mention not nearly as dangerous (not safe, but…).

I suppose the knock on sulfuric acid will be the added sulfate, not a big deal from my perspective.

When Tom Barr recommended vinegar for the process, I could not think of a reason why that was not a good idea, other than “organic material” and the fact it would take a lot of vinegar to reduce 8-dKH in 20-gallons.

So I set my targets a little lower, like 4-dKH from a quart. Turned out to be a good idea. Among other things HCl and H2SO4 are strong acids, they fully disassociate, Acetic acid is a weak acid, meaning it does not fully disassociate.

The real problem though is the things that happen when you turn an acetic acid loose acetates that form. So though you may have reduced the buffering of bicarbonate/carbonates but depending on the acetate you may have a greater buffer, but certainly makes it impossible to determine KH with hobbyist test kits.

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Liming*

Hi,

A question or two about liming have come to me.

The answer is yes, boiling or aggressively heating water can and will lower KH to 50-90-ppm (3-5-dKH), good enough for most purposes, but it is a pain and a questionable use of resources. Adding a quarter-teaspoon of gypsum per gallon of water can help.

Dilution is an obvious answer as well and perhaps the simplest, especially for smaller quantities.

The water I use is high KH and varies by the season, from about 12 to 17dKH. Copper is the problem child, _usually_ “they” are pretty good at keeping it down. 



The water into the house is filtered, first stage is a 5-micron sediment then a 5-micron charcoal filter followed by a 2-micron filter. No water softener as the GH is usually around 5-8-dGH and the brine into the wastewater system would be unconscionable. 

Before I get too self-righteous I will admit I have a system that a local company services, a cation/anion-uv-system that produces type 1, low organic lab water, but I only use that for “legitimate stuff.” And still have to by TOC water! No matter what “they” say.​ Liming, the use of Calcium oxide (CaO), known as quicklime is slaked, water added to form Calcium hydroxide (Ca(OH)₂), slaked lime. To remove carbonates (KH) from water. 



This process can also be a method of determining if there are other buffers at work in your water.
 I happen to use a 32-gallon trash (“they” say multi-purpose) plastic can, set on foot high platform. I set the thing near the water heater as hot water is an advantage. 



I do my KH reduction in 20-gallon batches, but I use 22-gallons, the extra 2-gallons means I do not have to be quite as carful at the syphon/filter part.
 
About the 4-dKH I like to use unless I am dealing with some plants that really balk at approaching 1-dKH. 



I like buffering each time I mention this one of the holy-people jumps in and says 1-dKH is all anyone ever needs,
fine that may well be true and if that is a matter of your religious belief, great.
I like 4-dKH because it gives me room to have a thing-or-two go wrong. No big deal.

First determine how much KH to remove, in my case I am removing 8-dKH, 



which is 8-dKH X (17.9-ppm CaCO₃/ dKH) = 143-ppm CaCO₃ (remembering we measure KH as CaCO₃).
 

Since CaCO₃ is 100.1-g/mol and Ca(OH)₂ is 74.1-g/mol
(I recommend Molar Mass Calculator).

 

(Ca(OH)₂ 74.1-g/mol) ÷ (CaCO₃ 100.1-g/mol) = 0.74 Ca(OH)₂ / CaCO₃
 We know that it will require 74-hundereths Ca(OH)₂ to remove each unit of KH.


Therefore 143-ppm KH X (0.74 X (Ca(OH)₂ / KH)= 106-ppm Ca(OH)₂
 Since I have 22-gallons of water


22-gal X (3.78-L / gal) = 83-L
 

83-l X 106-ppm Ca(OH)₂ X (mg / ppm) / g/1000-mg) = (~9-g Ca(OH)₂)
 So I will add 9-grams of CaO (quicklime) to a liter of water swirl for a bit to form a slurry.


Reality sets in here, since I am not absolutely sure of my KH level I will add another 2-grams or so of CaO to the water, distilled water is best and technically the CaO should be added to less than a full liter and topped-off with water to make a liter.

Then add about 400-ml or so of the slurry to the tap water, actually I take it from my water heater via my “clean” hose. 



I have my trash can marked for approximately 22-gallons. I stir the water for a bit.
 
Check the pH at this point, it will rise to around pH 10, pH 9.5 is plenty


 then continue to add slurry in small increments, the pH will decease rapidly, after you do this a few times, you kind-of know.
 Since pH = log 4KH + 7 = ~pH 7.6 (Or you could look at the pH, KH, CO₂ chart.)


Once in the range of pH 8, I work carefully. This is a good time to test KH.

I usually leave it overnight, the next day I satisfy myself the pH and KH is right. 

I siphon the water without disturbing the precipitate (debris on the bottom), I filter the water through a coffee filter in a strainer.

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Always Show Your Work!*

Hi,

I was asked about the manipulation of the formula that I used to calculate KH when CO₂ and ph are known.

Anytime we know two variables, the third can be calculated, basic algebra. 

I think I caused the confusion in the post above when I calculated the KH, based on atmospheric equilibrium and did not show (just assumed) the 3/3, which is of course 1 and multiplying something by 1, changes nothing.



CO₂ = 3 X dKH X 10^(7-pH)
 

dKH = 1 ÷ ((3 ÷ CO₂) X 10^(7-pH)
 

pH = log ((3 X dKH) ÷ CO₂) + 7
 
Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

H2SO4 + CaCO3 ---> CaSO4 + CO2 + H2O

If I don't aerate then the CO2 will react to reform a carbonate, but which carbonate is it most likely to be?


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## mattinmd (Aug 16, 2014)

The CO2 won't re-form a carbonate... It would need to pick up an extra oxygen to do that. 

Even without aeration, your tank water still has a surface exposed to air. The CO2 will eventually interact with the air and escape as gas (and possibly later re-enter after more interaction with air as individually CO2 molecules move back-and-forth).

Lack of forced aeration slows this process down as there's less surface, but aeration still happens unless you completely seal the container with no exposure to air (ie: sealed glass jar completely filled with water)


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## JoeRoun (Dec 21, 2009)

*Hmmm...*

Hi David,

H₂SO₄ + CaCO₃ → CaSO₄ + H₂O + CO₂ 

I may be misunderstanding but if the example you have given has taken place with nothing the reaction can only move to the left. The Calcium sulfate has precipitated, so a very small amount of carbonic acid and eventually the CO₂ gas will normalize based on partial pressure. There is obviously not enough water produced to dissolve the essentially insoluble CaSO₄.

I suspect I know where you wish to go… But, I’ll wait for your reply…:hihi:

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

mattinmd said:


> _*The CO2 won't re-form a carbonate*_... It would need to pick up an extra oxygen to do that.
> 
> Even without aeration, your tank water still has a surface exposed to air. The CO2 will eventually interact with the air and escape as gas (and possibly later re-enter after more interaction with air as individually CO2 molecules move back-and-forth).
> 
> Lack of forced aeration slows this process down as there's less surface, but aeration still happens unless you completely seal the container with no exposure to air (ie: sealed glass jar completely filled with water)


Really? But JoeRoun said on the thread "Water Chemistry"(http://www.plantedtank.net/forums/showthread.php?t=729473&highlight=):



> If you added sulfuric acid, the net result without strong aeration would be zero, nada, nothing, equilibrium will return


And he said 



> Any acid (sulfuric or otherwise) you add will “use up” (actually it is changing some bonds) buffering capacity. When you add enough acid to overcome the buffering the pH drops until a new “equilibrium” is reached. If you leave it alone at this point and nothing is done to remove the bicarbonates, _*the process will begin to reverse and those bicarbonate bonds will reestablish*_.


And he also said this:



> Avoiding getting technical, I need to ask you to accept that this is an “equilibrium reaction,” a fancy way of saying _*this reaction happens in both directions*_.



And on this thread he also said:



> Were you to be using sulfuric, hydrochloric or any other strong acid to lower KH (buffering) to lower the pH, you would need to strongly aerate, to blow-off (degas, the gentry would say) the C02 to keep the carbonates/bicarbonates (at our pH’s, nearly all bicarbonates), *since the chemical reactions move in both directions*.


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## JoeRoun (Dec 21, 2009)

*In Water Maybe*

Hi David,

Don’t be angry with Matt, be mad at me if anything…

This is the problem with multiple threads…

Matt correctly answered your formula, the only water is that produced from the reaction, if I dump a Stoichiometrically correct amount of Calcium carbonate to sulfuric acid, you will end up with some Calcium sulfate, a little bit of water and some liberated Carbon dioxide gas.

I think everything I said is reasonably accurate as was Matt’s answer.

Pick one thread do not assume everyone has read every thread and please tell us what you wish to accomplish, I think the chemistry is getting in the way.

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

I'm not angry with Matt, I just find you confusing. Is it possible if Matt answers my question, please?


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## mattinmd (Aug 16, 2014)

Sure, sorry for the delays.. I had several family things, chores, child, etc to attend to.

I cannot comment entirely on JoeRoun's point of view, as I don't believe all of it applies in context, but I suspect a lot of the difference amounts to scale of time.

If you add sulfuric acid, it will react with carbonate to produce CO2. 

If the amount of acid/carbonate is large relative to the water carrying it, well, you're going to get wild fizzing just as when you add a tablespoon of vinegar to a cup of water with a tablespoon of baking soda dissolved in it. The CO2 will be quite visibly driven out. I hope you never react this much acid and carbonate in a tank with fish in it, as it would gas them.

When the quantities of acid/carbonate reaction are small compared to the water volume a relatively small amount of CO2 is produced. In the absence of aeration that CO2 will be dissolved in the water. In the short term, some of that CO2 in solution will react with water to produce carbonic acid H2O+CO2 -> H2CO3. That is a carbonate, see the CO3?

However, assuming the water was already exposed to air, and at equilibrium with the atmosphere, the newly created CO2 in solution, and any carbonic acid resulting, is raising the level of CO2 in the water beyond equilibrium. It will find its way out of the water and escape into the air eventually, assuming there is any air exposure.

Vigorous aeration will speed the rate at which the water equilibrates with air. Nice, calm still water in an open container is still aerated by the exposure of its surface to air. This is a very slow interchange, but it is interchange none the less. The CO2 will eventually work its way out, even without forced aeration. 

All of this is somewhat artificial, assuming distilled water in laboratory beaker. It gets a lot more complicated in the case of an aquarium, as there's CO2 being introduced by the respiration of fish, inverts, and bacteria going on. If the plants are in light, they're removing CO2, if in dark they are adding to it. There's traces of a few dozen elements in the water, adding complexity to the reactions. 

That said, it is all generally moving towards an equilibrium.


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## JoeRoun (Dec 21, 2009)

*Reactions Over Reactions*

Hi David,

Matt can certainly answer or not as he pleases, I am certainly also free to post as I see fit.

I guess I am curious why you are throwing my posts up against him in a situation that isn’t relevant to the question you asked.

Unless it is that he said “but aeration still happens unless you completely seal the container with no exposure to air (ie: sealed glass jar completely filled with water)” well in that he is incorrect in a dynamic situation such as an aquarium or lake. That however, is not the reaction, you presented.

I think I get it… :confused1:

Respectfully,
Joe
FBTB


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## mattinmd (Aug 16, 2014)

Joe knows his stuff, he really does. He's taught me a few things....

A lot of this stuff can be really hard to explain clearly as you're always walking a fine line between simplifying things too much, and getting bogged down in every small detail that happens in a very complex system. 

As for the true question... 

The extra sulfur compounds are harmless in sane quantities. A mix of potassium sulfate, magnesium sulfate and calcium sulfate is often dosed as a hardness booster. Although the calcium sulfate is hard to dissolve, particularly if your water is already hard, it isn't really insoluble.

That said, I think it is a bit of a misnomer to try and think of the sulfur as being specifically bonded to calcium and/or magnesium when you are talking about things dissolved in a lot of water. 

If you add salt (NaCl) to water, you don't really end up with NaCl molecules floating around in water. You end up with sodium ions (Na+) and chloride ions (Cl-) floating around. They are attracted to one another a bit, but there's not any kind of tight bond anymore. I tend to think of it as being a lot like people dancing on a dance floor, or disco to fast songs. There's some pairing off, and the couples are mostly dancing near each other, but not constantly touching. So it is with most ionic compounds.

Thus if you add sulfate to water containing calcium and magnesium, that sulfate may find itself pairing off with calcium and/or magnesium, or anything else, but this pairing off doesn't really matter unless:

1) the ions end up interacting with something and get reduced to sulfide.
2) they end up strongly attracted to something and become insoluble, forming a precipitate that settles out. (in which case whatever it got attracted to falls out).

I don't think either are likely in the aquarium, perhaps some calcium sulfate will precipitate out, reducing your hardness.. but that involves a bit more chemistry (electronegativity) than I'm willing to do at 3 am while awake in the middle of the night


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## Seattle_Aquarist (Jun 15, 2008)

Hi JoeRoun,

It has been decades since I majored in Chemistry, but doesn't the Muriatic Acid break down the calcium carbonate into calcium chloride + other inerts? I dose calcium chloride to increase my dGH so doesn't this process increase general hardness?


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## mattinmd (Aug 16, 2014)

Seattle,

The dGH was already raised when the calcium carbonate was added. 

It is the calcium ions that raises dGH, and adding muriatic acid doesn't change the amount of calcium, unless for some reason it causes the calcium to precipitate (which would lower dGH).


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## somewhatshocked (Aug 8, 2011)

There is ZERO reason to insult other members.

The next time this thread has to be cleaned up: there'll be 30-day suspensions and perma-bans handed out.


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## JoeRoun (Dec 21, 2009)

*No harm, no disrespect intended anyone*

Hi All,

I will withdraw from The Planted Tank, I am not sure why my posts or this thread is so repulsive, or I should so anger people…:confused1:

With the exception of some grammar, minor misstatements and typos, what I wrote, I stand by. 

I have been wrong before and I am sure I will be again. :redface:

Perhaps I am not educated by your standards and English may not be my first language, but posts on forums as this are not academic treatise, they are written on the fly, in the midst of living our lives.:icon_surp

For the record, in answer to Roy’s question:



Metal carbonate + hydrochloric acid, gives us; 
metal chloride, plus water, plus carbon dioxide. 
Yes, you end up with, 

2 water(l) and 2Carbon dioxide(g), Calcium chloride(aq).

 At this point you have reduced carbonate hardness and if it has overcome the buffering, the pH will drop. 


Since this took place in an aquarium (volume of impure water) the
 Calcium chloride is now disassociated by the water into 

an ionic Calcium and 2 ionic chlorides. 



The general hardness is increased 2.5 times the Ca++. 

My most humble apologies the above answer to Roy's question is in error.

The Calcium does not increase the general hardness, I was not thinking. Or if I was I was thinking of the CaCl₂ being added.

From this point, things can get confusing as contrary to Matt’s assertions the CO₂ does not just automatically disappear, the reactions involving CO₂ are quite slow. 

This is why so many are frustrated when adding a strong acid (>1M) and find the water tends to return the original KH and pH. 


Hence the recommendation for aeration that seems to so infuriate many.
No harm, no disrespect intended anyone.:confused1:


Respectfully,
Joe
FBTB


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## mattinmd (Aug 16, 2014)

JoeRoun said:


> Hi All,
> 
> From this point, things can get confusing as contrary to Matt’s assertions the CO2 does not just automatically disappear, the reactions involving CO2 are quite slow.


I think we are on the same page Joe... to quote myself:



> Nice, calm still water in an open container is still aerated by the exposure of its surface to air. This is a very slow interchange, but it is interchange none the less


I really did mean that "very slow" part.. in calm water it will take a long time, but it will not hold that CO2 permanently. It will, eventually, work its way towards an equilibrium.

How fast it happens is a matter of temperature, water circulation, surface area exposed to air, airflow above the water, and a dozen other factors. You can really slow it down quite a lot if you try hard.

Remember a while back I thought the difference revolved around scale of time? I think my time scale here is quite large.. weeks, months, maybe years in extreme cases. 

Also, tanks complicate this. Air equilibrium with plain water is around 3ppm. However, tanks have a near constant influx of CO2 from life-forms in them. They will eventually establish their own equilibrium, but it will be above 3ppm because of the constant influx. 

This much is just like your home, which tries to equilibrate its temperature with the air outside. However, we constantly keep adding or removing heat to create a comfortable temperature indoors. We add insulation to slow down the rate of heat exchange. Never the less, it is still equilibrating, we're just changing where that equilibrium is.


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## JoeRoun (Dec 21, 2009)

*Well One More Detail*

Hi Matt,

I really did not mean any offense and I know I am no one’s idea of an expert, nor do I claim to be. :hihi:

Pretty much everything I have given here comes from industry and experience, without regard to how we think it should be, people in business to make money do not spend the kind of money they do on aeration for no reason, bicarbonate/carbonate (temporary) hardness is big business.

Nothing I have written here is that difficult to try on a small scale for yourself, this is physics meets chemistry. This is real world stuff.

I would be willing to state that in a high school, junior high, middle school, somewhere along the way someone had you do these reactions in a controlled (greater than a home aquarium anyway) environment and I just about sure at some point you were instructed to shake/twirl, stir or blow across the top of the test tube or container after the reaction took place. One way or another you aerated the solution to drive off the CO₂ gas, it is the tricky part about CO₂ and carbonates.

Anyway, have fun. TPT seems a hostile environment to me.

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

Is it not the case that...bicarbonates do not reform, but rather it only seems that way because the reaction is not given enough time to complete?

if i have a kh of say 20 dKH and i add enough acid to reduce that to 4dKH (and my co2 were to remain the same, which it won't) then my pH will drop from say pH8 to about pH 7.3 AND it will be stable at that level, despite the fact that it still has a "buffer". It will not bounce back. When we add into the equation the increased co2 released from the reaction between the metal carbonate and the acid, the pH will be even lower (and that, too, will eventually stabalise (aeration, as i understand things, only reduces the time for the reaction to complete, because it mixes the acid -- not because it removes co2 from the solution to stop it from reforming bonds).


(see diagram below)










I do do not know enough chemistry nor have the industrial experience that you have but i am confused as to your comments below



> If you added sulfuric acid, the net result without strong aeration would be zero, nada, nothing, equilibrium will return





> When you add enough acid to overcome the buffering the pH drops until a new “equilibrium” is reached. If you leave it alone at this point and nothing is done to remove the bicarbonates, the process will begin to reverse and those bicarbonate bonds will reestablish.





> I need to ask you to accept that this is an “equilibrium reaction,” a fancy way of saying this reaction happens in both directions.


The reason people experience "bounce back", as i understand things, is not because bicarbonates reform (and maybe they do?? I'm certainly confused.) but because the reaction was not completed in the first place: people stick acid in, take a quick measure a few seconds later (from a high concentration sample of acid-water mix) and then go back later on in the day (once the acid has dispersed more equally around the whole volume of water...then they notice that it has "bounced back" say 90%, and so they add some more, etc. It may be true that co2 may take a long time (in certain circumstances) to leave water, but i don't think this co2 reforms bonds with bicarbonates. 

matt:


> The CO2 won't re-form a carbonate... It would need to pick up an extra oxygen to do that.


Either way, I'm still confused.


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## mattinmd (Aug 16, 2014)

JoeRoun said:


> Hi Matt,
> 
> I really did not mean any offense and I know I am no one’s idea of an expert, nor do I claim to be. :hihi:


None taken at all. I'm certainly no expert myself. I am an engineer, but not of the chemical sort. Regardless, my degree did include basic chemistry.

Being an engineer, I'm used to vigorous technical debates, and may be attacking the ideas at times, but never the person.



> Either way, I'm still confused.


Personally, I still think a lot of the difference in Joe's viewpoint and mine amounts to time scale.

I don't know what kind of industry Joe is in, but it seems fairly obvious water treatment is involved. I'm going to guess at drinking water treatment, rather than industrial wastewater post-treatment, but the ideas are applicable to both. Both need the job done and done quickly.

In water treatment, you would really need to vigorously aerate the water to drive off the CO2 very quickly. If you fail to do this, once the water leaves the plant, it is inside a pipe and under pressure. There's no way that CO2/carbonic acid is going to escape at that point. There's no more exposure to air inside those pipes, so you're stuck with it all the way through the system until it reaches a tap. 

My point of view is of an open container of water that you can leave sitting around for as long as needed.. months, years, whatever. Most forms of industry do not have time to let water sit in an open container for periods of time this long, so processes taking this long are completely useless in that environment.

It really is all the same reactions.. Leaving water in an open container exposed to air aerates it... However, it does so at an absurdly slow rate. From an industrial perspective, this rate is so small it is indistinguishable from zero.

It is a bit like comparing the water flow caused by capillary action if you let a piece of coffee filter wick water up (maybe 1ml/minute?) with the amount of flow caused by a massive industrial water pump pushing 500,000 gallons a minute.


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## 141130 (Jan 25, 2014)

So do bicarbonates reform?


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## JoeRoun (Dec 21, 2009)

*Nothing To Add*

Hi,

Before I leave just some more lunatic ramblings…

I sometimes wonder if anyone pays attention to the crazed ramblings about acids being nothing but the definition of something with an H⁺ to give, hmmm, what do we have when HCl gives up its H⁺ or H₂SO₄ gives up its two H⁺’s, hmmm, chloride, sulfate, hmmm wonder where salts come from. 

Hmmm, Matt wants an O²⁻, someone wonders where to get H⁺’s, darn and dadgummit, I can’t think what water is made of, hmmm…

Well maybe someone smart will come along who isn’t well, a troll or a lunatic, hmmm

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

Autism isn't a mental illness: it's a neurodevelopmental disorder. It wasn't an insult. It would be great if you stayed.


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## mattinmd (Aug 16, 2014)

david765 said:


> So do bicarbonates reform?


Yes they will reform... but they cannot last indefinitely unless there is absolutely no aeration at all.

Most of our tanks, even without airstones, have quite a bit of aeration from the surface agitation. We need aeration to provide enough oxygen influx and CO2 outflux for our fish at night. Our fish would suffocate overnight, or at least be gasping at the surface in the morning, without it.

In a tank I don't think those carbonates are going to last more than a few hours.. There's too much air interchange... If it could last a long time, folks doing CO2 injection would find their CO2 levels wouldn't fall overnight.. but they do. Quite a lot.


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## 141130 (Jan 25, 2014)

mattinmd said:


> Yes they will reform... but they cannot last indefinitely unless there is absolutely no aeration at all.


I see the light. In a rigorous scientific experiment, where there is no aeration, the bicarbonates will reform, but in the practical life of an aquarist there is plenty of aeration to quickly stabilise things within a few hours. Great. Thanks.

So in serious-science situations (where the biocarbonates reform), the newly created salt reacts with the released co2 from the non-degased water (I presume deionised), and also the oxygen from the (deionised) water...but does this mean that the original acid "returns" in the testube (and goes in a cycle of re-action and re-re-reaction??)? I guess not, but why?

Thanks.


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## 141130 (Jan 25, 2014)

Given that adding acid sulfuric acid creates calcium sulfate and magnesium sulfate why does this not raise the GH? ...after all people specifically add calcium sulfate and magnesium sulfate to raise their GH?


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## JoeRoun (Dec 21, 2009)

*Just Rambling Mindlessly On*

Hi,

Just rambling on with random thoughts…:confused1:

Well I suppose carbonate hardness is temporary, but only a few hours, hmmm…

What was that about water? What was that formula again? Dadgummit must be the Thorazine drip…
https://www.youtube.com/watch?v=hnzHtm1jhL4 

How would knowing what water is made of, help? Darn that shock therapy…

Now think, even if I knew the water formula, how would I take it apart? Electricity, maybe during shock therapy, no wouldn’t happen all the time. 

Now I know I am crazy, could water disassociate? :icon_eek:Worse yet on its own…:eek5:

And just who or what is this Bronsted? :confused1:

Respectfully,
Joe
FBTB


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## JoeRoun (Dec 21, 2009)

*Thanks Zorfox, Finally Got It*

Hi,

Vision or hallucination?

H₂O + CO₂ ⇔ H₂CO₃ ⇔ HCO₃⁻ + H⁺

Respectfully,
Joe
FBTB


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## 141130 (Jan 25, 2014)

So the calcium sulfate and magnesium sulfate get their oxygen from the water and the co2 from the non-aerated co2 from the bicarbonates. okay. but given that adding sulfuric acid creates calcium sulfate and magnesium sulfate why does this (if we aerate the water) not raise the GH? ...after all people specifically add calcium sulfate and magnesium sulfate to raise their GH?

also, if the bicarbonates reform what do they reform from? the calcium sulfate and magnesium sulfate? if so do we end up with the creation of an acid (sulfuric acid)?


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## JoeRoun (Dec 21, 2009)

*More Ramblings*

Hi,

The hallucination continues…

CO₃²⁻+ 2H₂O ⇔ H₂CO₃ + OH⁻ ⇔ H₂CO₃ + 2 OH⁻

H₂CO₃ + 2H₂O ⇔ HCO₃⁻ + H₃O⁺ + H₂O ⇔ CO₃²⁻+ 2H₃O⁺



Originally the line below appeared as the second line Anthony noted, that it did not seem balanced, of course he is correct. I humbly apologize, I am expermenting with HTML code to try to make the equations readable, I messed up somehow.

H₂CO₃ + 2H₂O ⇔ H⁻ + H₂O ⇔ CO₃²⁻ + 2H₃O⁺
Respectfully,
Joe
FBTB


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## Darkblade48 (Jan 4, 2008)

david765 said:


> but given that adding sulfuric acid creates calcium sulfate and magnesium sulfate why does this (if we aerate the water) not raise the GH? ...after all people specifically add calcium sulfate and magnesium sulfate to raise their GH?


I am not sure where this confusion is arising from. 

Assume that your aquarium has calcium carbonate in it only. Calcium carbonate is relatively insoluble, but exists as an equilibrium between the solid form, and its dissociated form, that being Ca+2 and the CO3-- ions.

Adding an acid will cause the carbonate to react, forming H2CO3; however, this equilibrium is strongly favored in the forward direction, so the H2CO3 is driven towards CO2 (which will off gas) and water.

As a result of the latter reaction, the equilibrium for the calcium carbonate is perturbed, and more (solid) calcium carbonate moves towards the dissociation reaction (and solvates).

However, during this entire process, the *total number* of calcium ions does not change, so hardness will not go up. 



david765 said:


> also, if the bicarbonates reform what do they reform from? the calcium sulfate and magnesium sulfate? if so do we end up with the creation of an acid (sulfuric acid)?


The reaction for carbonates forming from CO2 and water is much more unfavorable compared to the reverse reaction. They will not reform from calcium sulfate nor magnesium sulfate. Sulfuric acid is not reformed.



JoeRoun said:


> CO₃⁻⁻ + 2H₂O ⇔ H₂CO₃ + OH⁻ ⇔ H₂CO₃ + 2 OH⁻
> 
> H₂CO₃ + 2H₂O ⇔ H⁻ + H₂O ⇔ CO₃⁻⁻ + 2H₃O⁺


I am not sure how you are deriving these equations; they do not seem to be balanced.


Also, as a general warning, the overall tone of this thread is devolving into one of petty argument and bickering. As mentioned by the moderation team, if the tone does not improve, this thread will be closed.


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## 141130 (Jan 25, 2014)

Thanks for your help Darkblade48.



> Adding an acid will cause the carbonate to react, forming H2CO3; however, this equilibrium is strongly favored in the forward direction, so the H2CO3 is driven towards CO2 (which will off gas) and water.


I thought adding an acid created a salt, H2O, and CO2?

acid + metal carbonate → a salt + water + carbon dioxide

Why is carbonic acid created?


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## JoeRoun (Dec 21, 2009)

*Trust Me I Understand*

Hi Anthony,

The above equations were an error on my part trying to use HTML code I deeply and sincerely apologize.:redface:

For my part I am not sure why I am targeted, an individual asked I not answer any of his questions I am honoring that request. Another, who may answer, is close; my ramblings were, I hoped of some benefit. If not, oh well. Apparently, you have judged me in the wrong. No bad intentions on my part. As with my deleted post there was nothing untoward, simply my, I believe, justified outrage. 

I am and was offended by the post, perhaps due in part to work I have done with disadvantaged people, I get most do not care, Perhaps I shouldn't. If someone does not wish my aid that is fine. 

Don’t worry it has been made clear I am not welcome on TPT, I actually am not as the poster described. I will shuffle along in a bit

Now for more ramblings to no one in particular:


Conjugate pairs may be an answer floating around
though the answers have been judged incorrect for TPT purposes,
what I wrote is correct in most of the rest of the world…

 Aqueous solutions are, well, the solution to the other part.
There are further reactions, sometimes line after line…
Wondering why pH and underlying KH change in rhythm with photosynthesis
particularly in non-CO₂ injected environments,
if bicarbonates/carbonates do not come and go who the heck is dumping in the washing/baking soda?

 
Respectfully,
Joe
FBTB


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## Darkblade48 (Jan 4, 2008)

david765 said:


> Thanks for your help Darkblade48.
> 
> I thought adding an acid created a salt, H2O, and CO2?
> 
> ...


Your reaction schema above is simplified.

Ultimately, the carbonic acid that is created (mostly) becomes water and carbon dioxide.


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## JoeRoun (Dec 21, 2009)

*Excellent*



Darkblade48 said:


> Your reaction schema above is simplified.
> 
> Ultimately, the carbonic acid that is created (mostly) becomes water and carbon dioxide.


Hi,

Most excellent.

Respectfully,
Joe
FBTB


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## DaveFish (Oct 7, 2011)

Seachem has all of this on their website. Is it incorrect or something? Why don't you go there and read it all. I have about a dozen times or more.


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## JoeRoun (Dec 21, 2009)

*Apologies All*

Hi,

My bad, I have not seen Seachem’s information on *Lowering KH via Lime & Muriatic Acid.*

Respectfully,
Joe
FBTB


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