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Re: [APD] Lowering KH, Which Is Better: RO Unit or Strong Acid?

> Date: Sat, 25 Feb 2006 17:38:58 -0500
> From: cS
>>The simplest arguement for RO vs ACID boils down to mere ease- of-
> use.
> Is using the acid THAT complicated and involved?...

Oh, I'm _quite_ sure that the eventual quantity- per- volume result could be 
reduced to a "plug and play" punch- the- numbers game.

Dave Wilson describes using mineral acids to reduce alkalinity in tanks of 
400,000 liters in his response. I must say that, were I dealing with 150,000 
gallons at a time, I might have difficulty coming up with 100,000 gallons of 
pure water with which to dilute the 50,000 gallons at hand. In a case such 
as this, the risks of using an acid to achieve the goal might not outweigh 
the cost effectiveness of the approach (economies- of- scale and all that).

> ... After all, do we not use such calculations to determine how much
> KNO3/CaCl2 to add?...
> ...Is it not simply a relationship of 17 ppm H+ neutralize 17 ppm
> HCO3- or am I simplifying too much?...

This is the basic tenet of titration as applied to, say, an alkalinity test 
kit - so many drops of acid at a known concentration to produce a color 
shift in the indicator means there was a predictable amount of carbonates 
present when the test was conducted. But a titration test will also give you 
a visible clue as to what's happening in the long run if you sit and watch 
it a while.

In order to produce the color shift, the pH of the solution must be at or 
below a specific point to effect the change. That means enough hydrogen ions 
must be consumed by the carbonates / bicarbonates to drive both of them to a 
solution of carbonic acid so that the addition of more hydrogen ions will 
shift the pH to the desired endpoint. At that point, you have driven the 
amount of carbonic acid to a level unsustainable by the water into which it 
is dissolved. When you stop adding the acid, the larger amount of carbonic 
acid will cause two things to happen - some will dissipate out into the 
atmosphere, and some will create "pressure" on the equilibrium point between 
H2CO3 and HCO3-, causing a percentage of the H2CO3 to convert back to HCO3-.

The part that's hardest to predict is how much will outgas and how much will 
convert back. One of the things that will affect this is the temperature of 
the water - warmer water will hold less gas. Another is the partial pressure 
of carbon dioxide in the atmosphere with which the water is in contact - the 
higher the partial pressure, the more will remain in the water. Yet another 
involves the surface area of the water that is exposed to the atmosphere - 
and turbulence creates more surface area just as would a larger (or wider- 
mouthed) container. If all of these combine to cause quick outgassing, then 
less of the H2CO3 will be converted back to HCO3-.

The indicator that makes the pH shift visible in one direction works just as 
well in the other direction. Once you've added just enough acid to your 
titration test kit to turn the blue indicator red, STOP ADDING ACID and 
watch the container. In a vast majority of the cases, the red will 
eventually return back to blue, indicating a return back to a more alkaline, 
less acidic condition.

Try it with warm water and with cold water. Try it in different- sized 
containers. Try it with less and less air in contact with the water in 
sealed containers. You can see a difference under controlled conditions.

This is the "rebounding" effect mentioned almost invariably in using acid to 
affect alkalinity.

Then add acid again to the kit - you'll notice it takes less to cause an 
indicator change the second time around. Depending on the original starting 
point of your alkalinity and the combined conditions of the first titration, 
you may or may not see a change back to blue the second (or third, or 
fourth) time around.

Simple dilution does away with all of these concerns and is far less time 
consuming. Figure how much you need to mix, do it once and use the water 
immediately. Without all of the associated dangers in handling, using and 
storing a strong mineral acid.

> Maybe my mistake in asking the question is that I referenced some a
> strongly concentrated product like muriatic acid (HCl) when I should
> be using something milder and less avian-flu-like such as Seachem's
> Acid Buffer (H2SO4, actually a bisulfate salt) as an example.  How
> would you feel about using Seachem's Acid Buffer to lower KH?  How
> much danger does that pose?  The concept is the essentially the same
> right: acid decreases KH?

Ummm... first off, H2SO4 is not a bisulfate salt, it is sulphuric acid - an 
even stronger acid than hydrochloric. Be careful of your chem shorthand.

But the point of Seachem's Acid Buffer is not the removal of the buffering, 
just a replacement of the type of buffer involved. Here you're replacing 
carbonate buffers with phosphate, which hold the pH value of the water 
almost as stable but at a lower pH value. This may, if you start with water 
containing carbonate buffers, simply involve overwhelming the carbonates 
with phosphates. Yes, the lower pH will decompose a lot of the carbonates, 
but you've achieved this at the expense of adding more compounds to the 
water and does nothing to reduce the total amount of compounds in the water 
the way dilution would. Probably OK for most plants, but creates 
"artificial" conditions for the fish. Which, if memory serves, is one reason 
Seachem suggests using the buffer to reconstitute RO or distilled water 
rather than say, convert Lake Tanganyika's "liquid rock".

Not to mention what it does for anyone's attempts to predict CO2 content 
during supplementation...


David A. Youngker
jaafaman at comcast_net

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