[Squawkbert]

My reading indicates that there's ~300ppm atmospheric CO2, ~3ppm of which is willing to dissolve into aquarium water. In streams, lakes etc, you get more than 3ppm CO2 in solution (underground, minerals, other organisms contribute to CO2 more in nature than in aquariums).

Note to MTechnik - I bet that checker didn't change color ove the course of Borat - the test tube appears to be too skinny. I'd suggest a fatter one and a lot less liquid in it (enough to get liquid column height=width should be good).

 

[banderbe]

My reading indicates that there's ~300ppm atmospheric CO2, ~3ppm of which is willing to dissolve into aquarium water. In streams, lakes etc, you get more than 3ppm CO2 in solution (underground, minerals, other organisms contribute to CO2 more in nature than in aquariums).

Note to MTechnik - I bet that checker didn't change color ove the course of Borat - the test tube appears to be too skinny. I'd suggest a fatter one and a lot less liquid in it (enough to get liquid column height=width should be good).

Where do you get that information?

I thought a liquid would always reach equilibrium with its surrounding atmosphere.. so there should be the same amount in a standing body of water as there is in the surrounding air, up to the point of saturation of the liquid of course.

I would love to know why CO2 magically stays out of equilibrium.

 

[dennis]

The information is everywhere. Atmospheric CO2 levels are ~350ppm or .03% of the total volume of the atmosphere. I'll stop now lest another thread becomes locked

As for why drop checkers don't show atmospheric levels... A great many google searches, chemistry books and teachers, etc will all explain that the reaction between CO2 and H2O it is a complicated and intricate, multi-step process, compounded when one has extra ions already present in the water. CO2 is actually very soluable in water, about 1.45g/L at 25 degrees C. This means that water is capable of holding, in theory, 1450ppm CO2. This would never happen though as solubility and readiness to dissolute are two very different things. Because CO2 dissolves into water about 10,000 times slower than into air, the equilibrium with pure water is only 0.5-5mg/l depending on temp, pressure and water chemistry. This is fortunate because otherwise, all the water in the world would have a pH of like 2 or something.

The reason your drop checkers turn blue very quickly once out of the water is two fold. First, when the checker is in the water, any CO2 that makes it into the drop checker will find its way into the drop checker water because of pressure and equilibrium. For example, if you have 30mg/l CO2 in your water, the water in the drop checker, since it is in a sealed system with the aquarium, will very quickly reach equilibrium due to a difference in partial pressures between the two waters. If the CO2 in the aquarium lowers, or if you remove the checker from the tank and expose it to air, the pressure/concentration of the CO2 in the checker is suddenly higher than its surroundings and the CO2 exits, causing the color to switch back to blue. This process is hastened by the extreemly small volume in the drop checker and by the chemical fact that the CO2<=>H2O equilibrium is very low even though the solubility is high.

The term equilibrium is a very confounding word actually. In almost all chemical reactions and processes, the concentrations of reactants and products is generally very squewed, even though the reaction is at equilibrium. A great many chemical reactions have equilibrium points at which the concentration of the products may be thousands or millions (sometimes even more) less that the concentration of what you started with. For example, if you look at the equilibrium point where H2CO3 and HCO3 are at equilibrium for one part of the CO2+H2) process, a concentration of H2CO3 in the water of 0.033M (hypothetical fully saturated level of the gas in water) the concentration of HCO3, at equilibrium, will be only 1.2x10^-6M!!!! Equilibrium doe snot mean the same.

Sorry for the lengthy post.

 

[MTechnik]

I replaced the velcro with fishing line and removed a lot of the reagent. I also added a piece of white solo cup clipped to the outside of the test tube, so that I have a white background for reading it.

You are very right - it wasn't green after Borat. Turns out I wasn't doing my CO2 dosing right, and now it's all fixed. This was just what the doctor ordered. Things have begun greening nicely, and soon I'll be able to trim away all the BBA covered growth.

-MT

 

[k-maub]

Interesting. So Dennis, is the implication that the small amount of air in the drop checker has an extremely high level of CO2 (if this scales linearly, something on the order of 3,000 ppm CO2), in order to maintain equilibrium?

Thank you for your explanation.

 

[dennis]

No. The air and solution in the drop checker, when its in the aquarium, has the same amount of CO2 as the aquarium water. In this case, the high equilibrium (high amount of CO2 in drop checker) is the result of the CO2 gas having no where to go but the water in the drop checker. In this case, if the pressure of CO2 in the aquarium water (called partial pressure and in this case is the same as concentration)) is higher than the pressure of CO2 in the drop checker, the concentration gradient will force CO2 out of the aquarium water in to the air gap and into the drop checker solution. It has to go there. The direction of change reverses as the CO2 in the aquarium drops and the drop checker has more CO2 than the aquarium... the CO2 will have a higher pressure/concentration in the check and will move to the aquarium. In this case, equilibrium is equal and it has no physical or chemical choice but to be as the two liquids are essentially in a closed system. Imagine pinching a balloon in the middle then inflating it so only half inflate... let the pinched part go and the whole balloon inflates. The balloon itself is the system and the air you push in must fill the whole thing.

If you take the drop checker out of the tank and put it in the air, even though the concentration of CO2 is higher in the atmosphere than in the checker, the CO2 will exit to the atmosphere anyway because there is nothing forcing a high equilibrium level. Its like opening the end of an inflated balloon so it flies around the room. True, pressure is involved more with the balloon, but the analogy should make it easier to understand. In the aquarium, CO2 goes into the drop checker because it has to go somewhere...in the atmosphere, CO2 leaves (or does not enter) because there are easier places for it to go.

Make sense?

 

[k-maub]

Hm. I see where you are going with that, but I don't see how the drop checker solution's interface has any "knowledge" of the condition outside of the drop cheker. That is, if the concentration of the air at the solution/air interface is 300 ppm, it should not matter what is going on at the exit of the drop checker. It is the relative values of CO2 in the solution and air that will dictate how much CO2 is diffused, I would think.

If 300 ppm CO2 in the atmosphere is only capable of raising water's CO2 ppm to 3, it is hard to believe that, even in a "closed system", air with ten times less CO2 (30 ppm) could possibly raise the drop checker solution to ten times more (30 ppm) that it would otherwise.

If CO2 is that much more easily dissolved in air than water, then by my reasoning, the 30 ppm water would reach equilibrium with 3,000 ppm CO2 air in the drop checker. Being ten times higher than atmosphere, this will bring the drop checker solution 10x higher than atmosphere, from 3 ppm to 30 ppm. Otherwise, if 30 ppm air can raise water to 30 ppm CO2, a sealed bag of half water/half air would eventually reach 300 ppm CO2 in the water (or 150 ppm or so once you subtract from the air's supply of CO2), possibly causing distress to fish being brought home from the fish store.

Am I making any sense?

 

[epicfish]

I enjoyed dennis' explanation, but I think it lacked a discussion of partial pressures, although he alluded to it with the phrase "nothing forcing a high equilibrium level".

Here's a list of terms and concepts I'll use to try to explain this...I don't want to clutter the text with this, so everything will be listed first.

http://en.wikipedia.org/wiki/Partial_pressure

Gases will dissolve in liquids to an extent that is determined by the equilibrium between the undissolved gas and the gas that has dissolved in the liquid (called the solvent).

The form of the equilibrium constant shows that the concentration of a solute gas in a solution is directly proportional to the partial pressure of that gas above the solution.

http://en.wikipedia.org/wiki/Co2

It is present in the Earth's atmosphere at a low concentration of approximately 0.04% and is an important greenhouse gas.

http://en.wikipedia.org/wiki/Atmospheric_pressure

In 1982, the International Union of Pure and Applied Chemistry (IUPAC) recommended that for the purposes of specifying the physical properties of substances, "the standard pressure" should be defined as precisely 100 kPa (≈750.062 torr) or 29.9230 inHg

http://www.google.com/search?hl=en&q=29.9230+inches+to+mm

29.9230 inches = 760.0442 millimeters

http://www.madsci.org/posts/archives/2001-09/1001605307.Es.r.html

Oxygen makes up about 21% or 210,000 ppm of the atmosphere.

CO2 constitutes 0.04% of atmospheric gases. At 760 mm Hg (sea level pressures), this amounts to 0.304 mm Hg. We're dealing on mm Hg (milimeters of mercury) since solvation of gases in liquids is based upon the concept of partial pressures.

It's true that there is about 300-400 ppm of CO2 in atmospheric air, but how does that compare with 210,000 ppm of oxygen gas? Or roughly 790,000 ppm of nitrogen gas? The partial pressure of CO2 is comparatively very very low in atmospheric air.

That's why even though there's 300 ppm of CO2 in the air, only a few ppm will actually dissolve in our aquariums.

As the CO2 in the aquarium comes out of solution into the air space within the drop checker, the partial pressure of CO2 in that air space rises dramatically because of the confined space. This increased partial pressure of CO2 in that space allows for a speed (within a few hours) equilibration between the drop checker liquid and the aquarium water.

When the drop checker is green, the water contains 30 ppm of CO2, which equates to a very high partial pressure of CO2 within that liquid. When you take it out of the water, this high partial pressure of CO2 within the liquid rapidly forces it out of solution to equilibrate with atmospheric CO2 which contains 300 ppm of CO2 but a very low CO2 partial pressure.

...hope that helps.

 

[k-maub]

That makes pretty good sense, but it also implies that in order to create the "very high partial pressure of CO2 within" the drop checker solvent, the drop checker's air would similarly need a very high partial pressure. Which to me would mean a much higher concentration than the normal 300 ppm atmospheric concentration, much less higher than 30 ppm...

What am I missing?

 

[epicfish]

Less oxygen and possibly less nitrogen in the drop checker air space would lead to an overall higher partial pressure of CO2 even though the CO2 amount is still 30 ppm.

 

[banderbe]

Now I remember why I got a degree in computer science instead of chemistry!

But seriously, instead of all the unnecessary rhetoric and quasi-scientific jargon, let's get to the point:

your drop checker should turn blue if exposed to air for an hour or so.

If it doesn't, you're doing something wrong. End of story.

 

[k-maub]

I suppose that much we can all agree on.

On the other hand, while I won't push the issue here, or at least start another thread to continue the conversation, I've always considered that one of the great things about this hobby is all the inadvertant knowledge from seemingly disparate fields one is exposed to.

That being said, I won't lead this thread further down this tangent; I very much appreciate the replies, especially the one to my original question.

Thank you.

 

[dennis]

Barry- YES

Epic, excellent and thorough job describing that, thank you!

One last thing then I will stop and let you get back to the fun part of making these things.... In response to k-maub's last post: Again, its partial pressure of the gas in liquid vs. the gas in air. At the surface of your tank, the CO2 is leaving the water and escaping into the air because the partial pressure of the air is less than in the aquarium water. Regardless of what ppm is in each, it is the "percentage" that epic described that leads to the difference. Your body works the same way by converting CO2 produced by muscles into carbonic acid for easy transport in the blood, then back to CO2 gas as it exits the capillaries in the lungs. The ppm concentration of CO2 (or carbonic acid HCO3) in the blood is much lower than the atmosphere but the pressure in the blood is higher, so CO2 readily transfers across the capillary walls and exits the body. With the drop checker, you have a similar situation. The little air space essentially is like the top of the aquarium and forces the dissolution of CO2 because of that. In addition, because it is a closed space, the CO2 concentration in that little gap is the same in ppm as in the tank, a much lower amount than in the atmosphere and a much, much lower pressure . In theory, this would make the drop checker work even faster underwater, of course it has to redissolve into the water of the checker which slows things back down again.

Sorry.

 

[Squawkbert]

OK - you've all covered this quite well and have made me ashamed at how much of this type of chem. I've forgotten.

The only little correction I can offer now is with respect to pH of carbonic acid solutions. Carbonic acid is a weak acid, so it will bottom out w/ respect to pH somewhere around 5.

Good discussion!

 

[gforster]

OK, I just got through reading the last 28 pages - whew! here's something maybe you can put into plain english for me. I would think most of us have scales for our ferts to measure things out. Many, like myself, are dosing PPS-pro where we have a nice solution in 1 liter bottles. Can you give me a "recipe" for 3dkh, 4dkh, and 5dkh solutions of distilled water in a 1 liter bottle?

For example -


desired dkh- 3dkh 4dkh 5dkh
distilled water- 1 liter 1 liter 1 liter
baking soda - ??? ??? ???

This way there shouldn't be so many people saying, "I didn't realize how little baking soda it took." Instead they can get it pretty accurate from the get-go. You could then double check from there and test to make sure. I've seen something that resembles math in some earlier posts, but I can't figure it out. Me and numbers just don't go together. You really have to spell it out for me. Anyway, thanks for the help.

 

[hoppycalif]

Baking soda is not an exact crystal form of a chemical compound. It can have lots of water incorporated into the structure or very, very little. And, that affects the weight of a given amount of bicarbonate. Also, sodium bicarbonate can contain some sodium carbonate, which changes the percentage of carbonate per gram. (If you heat baking soda to dry it, you also convert part of it to sodium carbonate.) So, it isn't possible to specify with great accuracy what weight of baking soda to mix with a liter of water to arrive at 4 dKH. But, if you buy a certified KH solution, whatever the KH is, you can accurately dilute it with distilled water to arrive at 4 dKH with good accuracy. That is how the 4 dKH solutions we can now buy are made.

 

[mhoy]

One teaspoon of baking soda will increase the hardness of 50 liters (13 gallons) of water by about 4 dKH.

See The Krib CO2 FAQ

 

[Squawkbert]

Assuming your baking soda is anhydrous (don't bake it out, undesirable chemistry will occur) - best to use a new box, dig down a bit...

I think these are the corrected values...
60.06mg per liter of RO/DI water = 4dKH
74.60mg per liter of RO/DI water = 5dKH

Now let's assume you don't have a balance that can accurately weigh out mg.
So the alternative is to make a more concentrated solution, then dilute.

6.0g per 5L RO/DI water = 80dKH (we'll call this the stock)
(1.20g in 1L will also work, if your balance is accurate enough)

1 part stock diluted to 20 parts total volume = 80dKH/20=4dKH - even multiples will also work. For example, 10mL stock + 190mL RO/DI=200mL total at 4dKH

25 parts stock diluted to 400 total parts = 80dKH*25/400=5dKH that would be 25mL stock + 375 mL RO/DI.

 

[mhoy]

Not sure if this would work. (Not sure which of the two holds onto water better though).
Save up some of those drying gels that come in the tiny packets. Dry them in a oven at low heat. Toss this into a sealed bottle, allow to cool then add some baking soda, wait a while and perhaps you have dry baking soda.

 

[styderman]

i got a scale for my guns and it goes to the .000 i think i put something like .029g in a gallon of distilled water. Correct me if this is wrong. It looks as if its working.