# The PH-KH-CO2 equation completely wrong? UPDATE on 22nd post



## John P. (Apr 10, 2004)

When you guys figure this out, let me know! :wink:


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## djlen (Sep 14, 2003)

You're putting us on........right?

Len


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## Rolo (Dec 10, 2003)

djlen said:


> You're putting us on........right?
> 
> Len


No, not at all! Except yeah, doesn't it sounds ridicules that the CO2 calculations we've been using might have been wrong this whole time? It took me four days of re-studying, researching, calculating, recheaking...making sure of myself that I _MIGHT_ be right. But I'm still waiting for the first person to point out my very simple/most careless mistake so you guys can laugh at me. :tongue:

I wonder how many people actually 'achieved' reading that post...


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## Laith (Jul 7, 2004)

I want to know too.

However, let's say that Chris is right  . So new charts/calcs are produced. All that would happen would be that, according to the new calculations, it is 12-15mg/l of CO2 that is the right amount for planted tanks and not 20-25. 

In other words, what we thought was 20-25mg/l of CO2 was actually 12-15mg/l. Same thing in the end vis a vis the plants no? The "optimum" would just go down to the 12-15 level.

Or is my logic off here? :icon_conf


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## KevinC (May 24, 2004)

This is good - after posting something on this in another thread, I realized I made the same error you did below:



Rolo said:


> dKH*17.8 = CaCO3 mg/L
> CaCO3 mg/L * (61.02/100.09) = HCO3 mg/L
> HCO3 mg/L * (mol/61020) = *1.78E-4M HCO3*


The problem above is that 1mole of CaCO3 neutralizes TWO moles of acid, so TWICE as much as HCO3- can neutralize. Therefore, you must also multiply by 2 when converting to HCO3 mg/L.



Rolo said:


> Therefore the only difference between my equation and the current one: 1dKH = 1.78E-4M HCO3 instead of 1dKH = 2.92E-4M HCO3. Completing the new equation...
> 
> [CO2] = [HCO3]*10^(6.37-pH)
> [CO2] = 1.78E-4*dKH*10^(6.37-pH)
> ...


Now the above "Correct" CO2 (mg/L) is actually 15.66*dkH*10^(6.37-pH)

and "correct" from your example is actually 29.10mg/L. This means the original equation gives only an 18% error (still not great - and since the chemistry is wrong anyway, who cares!). 

BUT - let me add to the confusion. There is a little thing we deal with in upper-level chemistry courses called "activity" - this represents the effect of waters of hydration (the ionic atmosphere) on the ion. It depends on both the ionic strength of the solution (total dissolved ions) and the size of the ion being considered. The bicarbonate ion has an activity coefficient between 0.82 and 0.96 depending on the ionic strength - this value is multiplied by the concentration to get the activity. Since we assumed that [X] = Activity(X) for everything above, the CO2 concentration is merely an estimate anyway.

Fun stuff!

Kevin


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## Rolo (Dec 10, 2003)

Thanks Kevin for the help!

I was thinking to myself the same exact thing when converting CaCO3 into HCO3, would it be 1 or 2 mol HCO3. But when using 2 mol of H+, aren't we assuming CaCO3 is directly taking part in the reaction. I was always under the understanding CaCO3 was merely just a representation; like for instance, that we could even use Cl2 gas as a measure of hardness if we wanted. And isn't the net balanced equation, where the product is HCO3:

CaCO3 + H+ <---> HCO3- + Ca++

So you would need 2 mol of CaCO3 to make 2 mol HCO3, meaning a 1:1 ratio, right? Or am I just completely off? lol. Yeah this is very fun stuff.

Now the original equation has only an error of 18% (actually as you go lower in CO2 concentration the error is less, but higher the error is more). However, I still believe that original equation should not be followed, if really in fact it is incorrect. It makes no sense in using something we know is wrong - and especially wrong in making such an elementary error. In light of the activity subject, I think this proves all the more we need to use the correct equation. Since we are only acquiring an estimate, compounding error on top of error just makes the whole pH-KH-CO2 relation useless.

I'm not ready yet just to take the quotes off our "correct" equation and make it official. Having only two people now agree isn't enough to completely erase all the old CO2 charts.


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## Rolo (Dec 10, 2003)

Yes laith your logic is correct.

Just a simple question. For anyone whos been into planted tanks long enough, when did the pH-KH-CO2 relationship begin to show up. Do you know of the source it probably first originated from?


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## KevinC (May 24, 2004)

Rolo:

I was going to quote some of your message, but then realized I would be quoting the whole thing - kind of pointless. Anyway, here's the deal:

1. alkalinity/buffering capacity/carbonate hardness actually is meant to be measured as the acid equivalents needed to neutralize to the carbonic acid endpoint (thus pH 4.5). Since carbonate ion requires two protons to become carbonic acid and bicarbonate ion only requires one, it is a 2:1 ratio. Expressing this as ppm CaCO3 (or dKh) is weird, but it is the convention.

2. I had sort of ASSUMED (know what that means), that the chart was the result of measurements of CO2, pH, and kH concs in a lab - this is very do-able with fairly simple (for a lab) equipment. Once this discussion came up it reminded me that I have seen a way to make your own CO2 sensor using a pH meter. The pH meter must be able to show the mV reading (not just the pH), but it is easy to construct and easy to calibrate. No, I'm not volunteering to do this experiment (I teach this for a living - don't want to do too much of it at home), just thought it might be of interest to some of you:



InfoTrac said:


> Journal of Chemical Education, Sept 1999 v76 i9 p1253(3)
> CO(sub.2) - potentiometric determination and electrode construction, a hands-on approach. (carbon dioxide)(Statistical Data Included) Santiago Kocmur; Eduardo Corton; Liliana Haim; Guillermo Locascio; Lydia Galagosky.
> Abstract: A student-used hands-on approach for potentiometric determination and electrode construction for carbon dioxide is discussed. To improve learning students build the sensitive device to detect CO(sub.2) and solve real-life problems, measuring the compound in a carbonated beverage, as do bottlers of soft drinks, mineral waters and some beers. In a post-lab analysis session, students discuss biological rationale for importance of monitoring gas content in yeast cultures in different growing conditions.



Kevin


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## Rolo (Dec 10, 2003)

KevinC said:


> Rolo:
> 
> I was going to quote some of your message, but then realized I would be quoting the whole thing - kind of pointless. Anyway, here's the deal:
> 
> 1. alkalinity/buffering capacity/carbonate hardness actually is meant to be measured as the acid equivalents needed to neutralize to the carbonic acid endpoint (thus pH 4.5). Since carbonate ion requires two protons to become carbonic acid and bicarbonate ion only requires one, it is a 2:1 ratio. Expressing this as ppm CaCO3 (or dKh) is weird, but it is the convention.


This makes perfect sense to me now. I made the corrections in the first post.



KevinC said:


> 2. I had sort of ASSUMED (know what that means), that the chart was the result of measurements of CO2, pH, and kH concs in a lab - this is very do-able with fairly simple (for a lab) equipment. Once this discussion came up it reminded me that I have seen a way to make your own CO2 sensor using a pH meter. The pH meter must be able to show the mV reading (not just the pH), but it is easy to construct and easy to calibrate. No, I'm not volunteering to do this experiment (I teach this for a living - don't want to do too much of it at home), just thought it might be of interest to some of you:
> 
> Kevin


Yes this does interest me. So where can I find the insturctions/equipment?


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## KevinC (May 24, 2004)

Rolo said:


> Yes this does interest me. So where can I find the insturctions/equipment?


What I posted is a reference to an article in the Journal of Chemical Education. Your nearby college/university library will either have hard copies of this journal or will be able to get you a copy of the article through Interlibrary Loan. I tried to find my copy here in my office, but it is not in the file folder it should be in. As I remember, you take a glass pH electrode and wrap a piece of porous material (like a nylon filter membrane) around it. There is a small amount of liquid between the membrane and the electrode. On exposure to a solution containing CO2, CO2 crosses the membrane, decreasing the pH of the internal solution. The response in mV is linearly related to the CO2 concentration. The more CO2, the higher the mV reading. I had thought about doing this for one of my courses, but haven't yet, so I don't have firsthand experience with how well it works.

As far as the equipment, you MIGHT be able to order directly from places like Fisher Scientific (fishersci.com) for a pH meter that measures mV. Be prepared - we're talking the $400 to $4000 area for a meter+pH electrode (that lets you see the mV reading). Maybe your local college/high school would let you borrow one.
Kevin


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## malkore (Nov 3, 2003)

Why not contact high school and community college chemistry professors and see if they could possibly have students do a lab over it...maybe an extra credit project.


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## bigpow (May 24, 2004)

Would be nice to see the CO2 table of original calculation & new calculation side-by-side...


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## jart (Jan 17, 2003)

rolo i will throw in $0.02. although i really don't know if this is correct. i'm wondering if there may be some sort of solubility co-effecient missing from one of the equations. when you are talking about converting caco3 to hco3, i'm wondering if there should be a correction factor since caco3 is not 100% soluble.

please don't be too hard on me if i'm way off base! i haven't done chemistry in a few years. 

consider for example the normal ph = 7.4, normal pco2 = 40 and normal hco3 = 24. pka of carbonic acid is 6.1. plug in the #'s and it won't work unless you multiply the pco2 by 0.03.

you are certainly to be commended for your efforts to work this out!


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## Rolo (Dec 10, 2003)

Jart-

I have no reason to be hard on you. We are all learning this together! When we speak about are hardness in CaCO3, our test kits really didn't measure CaCO3 and is not actually what we need. The only unit of hardess that is important is [HCO3]. CaCO3 by convention is the common standard that we express our hardness in. For instance, say something cost $10USD (HCO3), but we simply express our answer into $15 Canadian (CaCO3). If you can understand the poor example, I think you got it. Kevin, I might have explained this pretty bad, so mabey you can fill in the holes?

Bigpow - 

I'm going try to contact Chuck (the one with the fert calculators) and see if he could put in the new equation. I'll see about a side-side comparison too.

Kevin, Malkore - 

I didn't know I was asking for at least $400 to do this experiment, so I'm can't say just yet I'll voulneteer for the experiment. But I'm on really good grounds with my highschool chemistry teacher, so as soon as I back in september I'll see what I can do. An experiment however shouldn't be essential, as the math alone from the henderson-hasselbach equation provides only provides a good estimation of CO2.


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## Rolo (Dec 10, 2003)

In another forum, a user said he's had the same suspicions, and found this link here - http://fins.actwin.com/aquatic-plants/month.200207/msg00173.html If you look at the first equation, the constant is 15.7, like the 15.65 that was calculated in the 1st post. The constant of the current equation is 12.838. Someone else has also calculated the same constant!! :icon_bigg .

Currently the new equation: CO2 = 15.65*dKH*10^(6.35-pH)
I say "currently" since in another forum a member has brought up an issue of salinity, that also affects CO2 conc. It might be made more accurate. If you care to follow, this (pdf) is a link to the topic of salinity and its affect on pKa of H2CO3. 

Kevin - does this have anything to do with what you were saying about 'activity'?

I also corrected the pKa to 6.35 instead of 6.37 on the tip that 6.35 is the correct one for H2CO3. This mean at a KH of 5 and pH of 6.8 the CO2 is 27.8 mg/L. The original equation gives us 23.8. Therefore the current %error is 14%.


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## KevinC (May 24, 2004)

Rolo said:


> Currently the new equation: CO2 = 15.65*dKH*10^(6.35-pH)
> I say "currently" since in another forum a member has brought up an issue of salinity, that also affects CO2 conc. It might be made more accurate. If you care to follow, this (pdf) is a link to the topic of salinity and its affect on pKa of H2CO3.
> 
> Kevin - does this have anything to do with what you were saying about 'activity'?


Chris:

Indirectly, yes I think it does. They are changing the K while I (and other chemists) advocate changing the apparent concentration based on the "salinity" (ionic strength) in the solution.

FYI - I teach college chemistry (including Analytical chem) at a small 4-year college (The Citadel). For the past several years I have had students in my Quantitative Analysis course measure various water quality parameters (pH, salinity, O2, Ca, Mg, NH3, NO2, NO3, SO4, PO4, etc) in local bodies of water (brackish and fresh). This year I'm shifting gears - we are going to study various brands of bottled water (using many of the same tests). Maybe I can get a sophomore or freshman chem major interested in our experiment above. I usually have one or two I'm trying to break into research.

My webpage:  Crawford
Kevin


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## Rex Grigg (Dec 10, 2002)

The Citadel is not a small college. It's a great college.


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## John P. (Apr 10, 2004)

Rex Grigg said:


> The Citadel is not a small college. It's a great college.


One of my family friends (& high school classmates) attended the Citadel in 1991 ... now a major in the US Army Rangers. roud:


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## KevinC (May 24, 2004)

Rex Grigg said:


> The Citadel is not a small college. It's a great college.


Rex-

I do like it here, though I've never been in the military (faculty wear uniforms - we are members of the Unorganized Militia of South Carolina - but that is mainly to fit in the chain of command). 

One thing I find surprising though - outside of military people, alumni, and people from the Southeast, many people had not heard of The Citadel before the coeducation fight (myself included) and many have not heard of it since. This year we got a very nice ranking from US News in their College Rankings issue (#2 for public universities in the South offering up to Masters degrees). I think The Citadel is a very fine comprehensive liberal arts college. Another thing many people don't realize - only about 33% of our graduates actually go into the military. The rest are here for the disciplined environment, the alumni network, because a relative went here, etc.

Kevin


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## Rolo (Dec 10, 2003)

While I never heard of the Citadel, I am considering a military college. All depends though since the one I'm considering binds you the military service.


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## Rolo (Dec 10, 2003)

By the looks of things now, a new CO2 chart does not have to be made. Currently the maxium amount of error of the current charts are 14%. When using the correct equation, and factoring in the affects of temperature and ionic strength (salinity) in calculating the conc. of CO2, the %error drops.

_We have been fortunate that the large mistakes in finding the original equation luckily balanced themsleves out!_

So if you have a reading of 20ppm CO2, really the most you have is 22.8ppm. If you have 30ppm, thats actually 34.2ppm. This amount of error can easily be made by our test kits, and we arn't talking about perfect analytical chemistry. _IN GENERAL_, if your water is soft & has a low amount of ionic solutes, then the error is closer to 14%. The harder water gets, including many other ionic solutes (NO3, SO4, PO4, Na, K, Cl...) the % error drops. 

Here are both equations in their raw form. Niether take into account the temp. or ionic strength.

Correct CO2 = 15.65*dKH*10^(6.35-pH)
Current CO2 = 12.84*dKH*10^(6.37-pH)


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## ninoboy (Jan 26, 2004)

Rolo, have you seen this : http://www.aquatic-plants.org/khph_table.html
Do I miss something or that chart is totally screw up :icon_roll ?


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## malkore (Nov 3, 2003)

Nino, you beat me...I was gonna post that too cuz I saw it on plantgeek.net this morning.

It has to be screwed up, as I'd have 69ppm of CO2 in my tank, yet my fish are super healthy and never gasping for air first thing in the morning.
That or we underestimated the amount of CO2 fish can handle...but since the rest of Chuck's site is so 'correct', and the chemistry guys are doing the calculations...i think that other sites just wrong.

Well I sent their webmaster an email to see how they arrived at the figures in their chart, and sent them a link to this thread in case they just wanna explain it themselves instead of going through me.


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## dennx (Aug 11, 2004)

I think the above chart is correct. Most likely what happened was when they built the tables for the chart it was done incorrectly. Just shift the pH values to the right two columns and everything matches up as they should.


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## Rolo (Dec 10, 2003)

Thanks ninoboy and malkore for bring this up and e-mailing the site. I too just saw it this morning. Kinda worried me cause I thought we had this KH-pH-CO2 thing corned. Dennx is right though. The pH values just got shifted on the chart so they have to be realigned.

Anyhow, it still is using the equation that has a 14% error to it. Keep in mind that for the values between 20 - 30 ppm CO2, we are underestimating by 3-4ppm.


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## malkore (Nov 3, 2003)

Duh! I totally should've noticed that. Having a new puppy depriving you of sleep will screw your brain up.

I emailed them back to let them know they still need to fix the table by shifting pH 2 cells to the right.


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## Georgiadawgger (Apr 23, 2004)

KevinC said:


> Rex-
> 
> I do like it here, though I've never been in the military (faculty wear uniforms - we are members of the Unorganized Militia of South Carolina - but that is mainly to fit in the chain of command).
> 
> ...


Go Furman!!! :icon_mrgr


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## oldfarmhouse (May 18, 2004)

Maybe the old figures are a built in mechanism to keep all us non scientist from killing our fish.. Robert Parrish went to the Citadel The great center from the best front three of all time in the NBA (Larry Bird, Kevin Mc Hale, Robert Parrish) Boston Celtics!!


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## SCMurphy (Oct 21, 2003)

You guys have confused me.

CaCO3 is calcium carbonate, normally a precipitate
Ca(CHO3)2 is Calcium bicarbonate, which is solulable. 

Which one are you working with? 

Do I have these right?
CO2 Carbon dioxide
CH2O3 carbonic acid
CHO3- bicarbonate ion
CO3-- carbonate ion


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## Rolo (Dec 10, 2003)

SCMurphy said:


> You guys have confused me.
> 
> CaCO3 is calcium carbonate, normally a precipitate
> Ca(CHO3)2 is Calcium bicarbonate, which is solulable.
> ...


So that we don't get more confused, the names of the compound/ions are:

CO2 Carbon dioxide
H2CO3 Carbonic acid*
HCO3- Bicarbonate ion*
CO3-- Carbonate ion
CaCO3 Calcium carbonate
Ca(HCO3)2 Calcium bicarbonate*
H+ Hydronium ion
*You just switched some of the H and C around. :tongue: 

For the pH-KH-CO2 charts, all we are working with are CO2, HCO3, and H+ (which pH measures). Ca(HCO3)2, H2CO3, and CO3 are not important. So forget these. The CaCO3 is not at all existent in our tanks - this only has to do with are hardness readings. 

pH = 6.36 + log([HCO3]/[HCO3])
pH = 6.35 + log([HCO3]/[CO2])
[*CO2*] = [*HCO3*]*10^(6.35-*pH*)
ppm CO2 = 15.65*dKH*10^(6.35-pH)

That is the breakdown of the equation in its raw form to our useful one. I assume you read the first post which would explain all the calculations. As you can see, the only three players in the relationship are CO2, HCO3, and pH (H+)

So where did H2CO3 come in? The Henderson-Hasslebach equation finds pH from buffered solutions. The carbonate buffer we have is between H2CO3 and HCO3 - so it is correctly written with H2CO3. I only somewhat understand why, but we are allowed to replace H2CO3 with CO2. 

So where did CO3 come from? KH kits measure HCO3, CO3, and other pH - resisting buffers. At a pH lower than 8.32, ~100% of the carbonate buffer system is HCO3 and little CO3 exists. So CO3 is eliminated, which is a good thing. This means at a pH higher then 8.32, the chart is invalid since our KH kits are affected by both CO3 and HCO3 - we only want HCO3. The same logic applies if we have a lot of other non-carbonate buffers, which influence KH and make it appear we have more HCO3 than actual.

The CaCO3 is simply the units of our hardness. The person doing the calculations for our current charts thought when the KH kit reads, "150ppm" that it was in HCO3. _This is the reason why the current charts have a 14% error._ By convention it is really expressed as CaCO3, not HCO3; the amount of calcium carbonate we must dissolve to get that particular KH level. A few calculations take place to convert ppm CaCO3 into ppm HCO3. Thus CaCO3 is not important, it's just in our hardness unit.

And as you well know, CO2 and pH are involved. I don't know where Ca(HCO3)2 is, but that doesn't take any place here.


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## KevinC (May 24, 2004)

Chris:

Couldn't have said it better myself :icon_bigg 

FWIW, here are the relevant reactions/equilibria (as best I can write them in this format):

1. CO2(gas) <----> CO2(aqueous)
2. CO2(aqueous) + H2O(liquid) <----> H2CO3(aqueous)
3. H2CO3(aqueous) + H2O(liquid) <----> H3O+(aqueous) + HCO3-(aqueous)
4. HCO3-(aqueous) + H2O(liquid) <----> H3O+(aqueous) + CO3-2(aqueous)
5. CO3-2(aqueous) + Ca+2(aqueous) <----> CaCO3(s)

As Chris mentioned - we have the Henderson-Hasselbach equation (which results from #4). In addition, each reaction above has its own equilibrium constant (if you don't know what that is, don't worry - it's not important at this point).

This set of equations really does explain a lot. For instance: adding CO2 decreases the pH (increases the amount of H3O+) - equation 1, 2, 3.

Adding limestone (CaCO3) increases buffering capacity: equation 5, 4 (reversed so the product is HCO3-).

Notice that HCO3- can react WITH acid (reverse of equation 3) and can react with water to PRODUCE acid (equation 4). This is what makes it a buffer - it can remove small amounts of acid OR it can remove small amounts of base (by reaction with the acid it produces).

Fun Fun Fun! Now I've got to go teach some freshmen what an atom is!

Kevin


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## SCMurphy (Oct 21, 2003)

In the name of helping make sure the underlying chemistry is correct since you are finding such an interesting conclusion....and I find the chemistry of CO2 dissolution interesting. :icon_bigg 

CO2 in water is only partially dissolved, it starts the same way other gasses dissolve, what ever partial pressure is. Then the fun starts as CO2 associates with H2O and becomes H2CO3. So while I can see H2CO3 being considered CO2 it doesn't work saying CO2 is H2CO3. 

CaCO3 is the white precipitate you get when the pH rises above 8 (also it is the product of biogenic decalcification--when the water is short of CO2 and the plants, that can, strip CO2 from bicarbonate). In acidic conditions isn't Ca++ a free ion not attached to HCO3-? 

Ca is usually measured by GH (permanent hardness along with Mg and a few other ions), KH is usually HCO3 (temporary hardness). You loose me when you say that KH is CaCO3.


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## KevinC (May 24, 2004)

SCMurphy said:


> In the name of helping make sure the underlying chemistry is correct since you are finding such an interesting conclusion....and I find the chemistry of CO2 dissolution interesting. :icon_bigg
> 
> CO2 in water is only partially dissolved, it starts the same way other gasses dissolve, what ever partial pressure is. Then the fun starts as CO2 associates with H2O and becomes H2CO3. So while I can see H2CO3 being considered CO2 it doesn't work saying CO2 is H2CO3.


Yes, that's what I meant by saying that each reaction has its own equilibrium constant (they do not go 100% to products). It is actually a pretty complicated INTERCONNECTED set of reactions. If you influence one (say injecting CO2, thus increasing its partial pressure), you change all the others - though the further you are from the original, the less dramatic the change.



SCMurphy said:


> CaCO3 is the white precipitate you get when the pH rises above 8 (also it is the product of biogenic decalcification--when the water is short of CO2 and the plants, that can, strip CO2 from bicarbonate). In acidic conditions isn't Ca++ a free ion not attached to HCO3-?


Yes, the great majority of calcium in a tank is dissolved Ca++, but remember - most people have snails - the shell is composed of CaCO3 (mostly). In equilibrium reactions we treat EVERY concentration as a NON-zero number. It may be extremely small, but not actually zero. We also assume excess solid is present somewhere in the system. Without any CaCO3 and at lower pH's, the last reaction is just not a factor.



SCMurphy said:


> Ca is usually measured by GH (permanent hardness along with Mg and a few other ions), KH is usually HCO3 (temporary hardness). You loose me when you say that KH is CaCO3.


The kH kits REPORT the amount of CaCO3 (in ppm or dkH) you would need to dissolve to produce the same resistance to pH change - so every chemical that resists pH chage is treated like it is calcium carbonate - thus the name.

Kevin


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## Rolo (Dec 10, 2003)

Kevin hit everything right on. roud: 

Just to add to the KH confusion, when our readings are in dKH, the units are CaO (calcium oxide).?? Ok ok, to make it very simple - you call up a US book store and ask for the price of a book. The merchant looks it up, see the actual amount as "20 dollars", but then tell you "500 pesos." HCO3 is the dollars, and pesos is CaCO3. We don't pay for things in the US with pesos, so the 500 pesos is just a representation of 20 dollars. :icon_roll . Same logic applies to GH kits. Why the charts have an error is b/c they didn't do the conversions.



SCMurphy said:


> CO2 in water is only partially dissolved, it starts the same way other gasses dissolve, what ever partial pressure is. Then the fun starts as CO2 associates with H2O and becomes H2CO3. So while I can see H2CO3 being considered CO2 it doesn't work saying CO2 is H2CO3.


I'm very glad this point came up. Kevin explained everything correctly as all the various carbonate species are interconnected to each other. But I too do not _fully_ understand why we can replace H2CO3 with CO2. I've been doing a ton of research and replacing H2CO3 and CO2 is a VERY common procedure. They don't explain why though. Especially all the medical papers I read through do this to find the conc. of CO2 in blood...except they use "0.03 pCO2."

Anyhow, what I don't get is how [H2CO3] = [CO2]. The K (constant) between CO2 and H2CO3 is I believe 2*10^-3. The equation is K = [CO2]/[H2CO3]. Basically meaning very little CO2 becomes H2CO3. I'm sure that they are equal at a pH of around 4.5. Bla bla, ok maybe Kevin can help clear this. roud:


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## Rolo (Dec 10, 2003)

Nevermind.

This site cleared it up for me very well. http://www.chemistry.wustl.edu/~edudev/LabTutorials/Buffer/Buffer.html The yellow box (13-18) shows mathmatically how we go from H2CO3 to CO2 in the H-H equatoin. The site also starts from the very basics of acid-base reactions and buffers, so if your really lost it should help.


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## SCMurphy (Oct 21, 2003)

That's an excellent reference, thanks for finding it. I hope you guys don't mind me digging through the chemistry like this. I don't disagree with any of the conclusions. I'm just trying to help make sure the supporting discussion is complete.


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