View Full Version : Why is carbonic acid formed?
JonJon82
09/16/2013, 12:26 AM
Follow me, carbonic acid is formed when a CO2 molecule interacts with an H20 molecule. Both molecules are held together by polar covalent bonds. Since they are both polar the molecules attract each other, but what makes them combine into carbonic acid H2C03 when both molecules already have stable bonds?
From what I've read, not alot of carbonic acid is formed from the C02 in the water, even though CO2 is surrounded by H20 molecules because it is attracting them. Why is it only sometimes created when the two molecules meet?
Also, the oxygen ends of C02 are the electronegative ends right? I havent seen a picture of C02
Jimmy54
09/16/2013, 05:17 AM
pH ;) The ratio between H+ and OH- ions i.c.w. the ratio between CO3 and HCO3 ions and the osmotic pressure of CO2.
And that's just one explanation. Just keep in mind that "water in solution" are H+ and OH- ions flying all over the place.
(There is even H3O ;))
The equilibrium equation is;
CO2 + H2O <=> H2CO3 <=> H+ + HCO3- <=> H+ + H+ + CO32-
Jimmy
disc1
09/16/2013, 08:00 AM
CO2 is non polar. You really don't get any more non polar than CO2. So there's no structuring of water molecules around it like there is with ions and highly polar molecules.
This is hard to explain all the way out, so read this and then use google to get at any new words.
What is happening is a chemical reaction. That polarity on the water molecule leads to the oxygen having a partial negative charge. That makes water a good nucleophile.
CO2 has a carbon atom sitting between two oxygen atoms. The oxygen atoms are very electronegative just like they are in water. This pulls electron density away from the carbon atom and leaves a partial positive charge. Thus, the carbon atom is a bit of an electrophile.
Put a nucleophile and an electrophile inclose proximity and you get a reaction. I can't draw the mechanism, so I'll do my best in words. The oxygen atom on the water molecule attacks the carbon atom in CO2 forming a covalent bond. One of the double bonds to one of the oxygen atoms on CO2 is broken and becomes a single bond as the carbon atom forms the bond to the oxygen on the water molecule. Now the oxygen on CO2 with the single bond has a negative charge, and the oxygen from the water molecule, which still has 2 hydrogens attached, has a positive charge. The positive oxygen loses one hydrogen ion to become neutral and the negatively charged oxygen picks up a hydrogen ion to kill the negative charge and...
viola... You have carbonic acid.
JonJon82
09/16/2013, 03:19 PM
So this is what I understand from your posts:
Formation of Carbonic Acid:
Whenever the Carbon atom in C02 touches the Oxygen atom of H20 they stick together and form a single covalent bond. This is the first step in the formation of Carbonic acid. This weakens the bond between Carbon and one of its Oxygen atoms in C02 but all three atoms of CO2 remain attached. The Oxygen atom with the weakened bond in C02 scoops up an H+ ion from the water and the Oxygen atom from H20 releases an H+ ion into the water.
Carbonic Acid and pH:
Depending on how many Hydrogen atoms Carbonic Acid loses, Carbonate or Bicarbonate will be formed, which float around eating up H+ ions. But if they eat to many H+ ions they will turn back into an acid, which can then re-release the H+ ions in a constant cycle. However Carbonic acid is a weak acid, which means it will not release many H+ ions into the water. This also means not a lot of Carbonate and Bicarbonate will be created. This all means CO2 does create a helpful, natural buffer against pH changes. As long as C02 levels don’t become become too high. The buffer system fails if H+ ions greatly outnumber OH- or vice versa.
The covalent bonds that keep Carbonic Acid from releasing H+ ions into the water can come apart the same way some water will always split apart into Hydrogen and Hydroxide, or if Carbonic Acid comes into contact with an OH- ion?? In this case, Bicarbonate and water are formed. This means OH- ions will neutralize Carbonic acid and free H+ ions.
Does this mean the higher your KH the more basic your water becomes? It seems to indicate that it does. I definitely need to increase my pH from the neutral pH tap water to somewhere around 8.5 This requires a tenfold decrease in H+ ions.
Carbonic Acid reactions:
If one H+ ion breaks away from Carbonic Acid, HC03- is formed (Bicarbonate)
If both H+ ions break away from Carbonic Acid C032- is formed (Carbonate)
If Carbonic Acid comes into contact with an OH- ion HC03- and H20 is formed (Bicarbonate and water)
If Bicarbonate eats up an H+ ion it will turn back into Carbonic Acid
If Carbonate eats up two H+ ions it will turn back into Carbonic Acid
blanden.adam
09/16/2013, 04:29 PM
Well, almost, but the caveats are really beyond our purposes here. Suffice it to say that this equilibrium happens:
CO2 + H2O <--> H2CO3 <--> HCO3- + H+ <--> CO3(2-) + 2H+
All of these species are in equilibrium at any given pH. In the reef aquarium, we can interact with this equilibrium in a number of places. We can add hydroxide from Kalkwasser for instance, which consumes H+ ions and shifts the equilibrium to the right (via Le Chatlier's principle). Similarly, we can add carbonate, again with the potential to consume H+ ions and shift the equilibrium to the right. This also raises the pH, but temporarily. Conversely, we could add straight H+ ions to the system in the form of a mineral acid, which would push the equilibrium to the left, decreasing bicarbonate, or even alter the most important player, CO2, which we will discuss below.
Notice, there is a volitile component to this buffer system -- the CO2. The rest of the components are "trapped" in solution as they are non-volatile. CO2 however will be in constant equilibrium with the surrounding air (provided there is good flow and the surface is properly agitated). Because CO2 is essentially the equivalent of adding Carbonic Acid (from which you get equal parts protons and bicarbonates), higher CO2 means lower pH, but unaltered alkalinity. Because CO2 is free to come and go as it wishes (or at least how thermodynamics dictates it will), the CO2 is the ultimate arbiter of pH in a reef aquarium.
While it is true that at any given CO2 concentration the pH will be higher with higher alkalinity, the best approach to dealing with pH is to make sure your alkalinity is in an appropriate range (8-12 dKH give or take a dKH unit is usually recommended), and then address the CO2.
disc1
09/16/2013, 04:36 PM
Well, you've got part of it. The rest of it is hidden in a concept called chemical equilibria. Here's the readers digest version.
First: A note on terminology. I'm going to call the hydrogen ion a proton. Hydrogen has one proton in its nucleus and the cation has lost its one electron, so a naked proton is all that remains. I kept stopping myself on the last post because I didn't want to be confusing. Maybe it will be easier just to explain that bit.
None of these reactions are one way. They all go both ways. The water attacks CO2 to form carbonic acid, but at the same time carbonic acid is breaking that bond and becoming carbon dioxide and water. The same can be said for these acid base reactions (actually ALL acid base reactions) so bicarbonate is constantly being converted to carbonate by losing a proton and carbonate is being converted to bicarbonate by gaining a proton, bicarbonate is going to carbonic acid by gaining a proton and carbonic acid is going to bicarbonate by losing one.
The thing that is different about all those two way reactions is the rate of the reaction. So carbon dioxide is combining with water to make carbonic acid a LOT faster than the carbonic acid is breaking down.
When we talk about the rate of a reaction, it is dependent on the concentrations of reactants and products. If you have a higher concentration of the reactant or a lower concentration of a product, then the reaction goes faster. In reality it is way way way more complicated than that, but that is the basis of LeChatlier's principle.
Let's leave carbonic for a minute and talk about precipitation. It is a little easier to understand equilibrium from here and we'll go back to the acid base thing. For a precipitation reaction, let's use calcium carbonate since we're saltwater aquarists, the reaction looks like this. (aq means aqueous or dissolved in water)
Ca<sup>2+</sup> <sub>(aq)</sub>+ CO<sub>3</sub><sup>-</sup><sub>(aq)</sub> <-----> CaCO<sub>3</sub> <sub>Solid</sub>
The arrow is double headed because the reaction goes both ways. It is important to remember that the reaction is always going both ways at least to some extent. Now at low pH, the backwards reaction goes much faster than the forwards reaction. So calcium carbonate is dissolved. At high pH, the forward reaction dominates and calcium carbonate precipitates. Add some straight kalk water to a little tank water and spike the pH to about 11 or so and watch what happens. Calcium carbonate is precipitated at high pH. Why that happens differently at different pH is beyond the scope of this discussion, but it isn't hard to find on the internet.
But those rates are also dependent on the concentration of all the species. So as the calcium carbonate falls out of solution, there is less reactant and more product (actually in the case of precipitation we consider it the same amount of product, solids all count as a concentration of 1 no matter what the unit is... But work with me here) so the rate of the forward reaction slows down and conversely the rate of the reverse reaction speeds up.
Pretty soon we hit a magical point where the rates of the forward and reverse reactions meet. Where both reactions are going at the same speed. For every molecule I precipitate, I also dissolve one. And nothing changes. It's called equilibrium.
Reactions that can go both ways like that are called equilibrium reactions because they naturally tend to find an equilibrium. No matter which rate starts off faster, it will slow down and the other will speed up until they meet at some point and things will stay there until acted upon by something outside the system. It's a beautiful concept.
To Be Continued...
Twistofer
09/16/2013, 05:01 PM
I loved Chem 101/102 :cool: Takes me back 35 years to those wonderful college years as a freshman Pre-Med.
Wait until next time for Chem 201/202, Organic Chemistry and the organic behavior of Carbonic Acid :lolspin:
PS: Just throw in all those inorganic and organic molecules, like ammonia, nitrite, nitrate, lactic acid, urea/uric acid, which are found in your average aquarium's soup, and you'll really have a good time.
disc1
09/16/2013, 05:55 PM
Picking up from post 6.
So it is with these carbonic acid molecules. Everything is an equilibrium. Let's look at just the bicarbonate and carbonic acid for now.
H<sub>2</sub>CO<sub>3</sub> <------> H<sup>+</sup> + HCO<sub>3</sub><sup>-</sup>
Again this goes both ways and the relative concentrations all depend on the equilibrium constants. Each reaction has equilibrium constants which allow you to calculate the rates of reactions from the concentrations of reactants and products. Let's not get quite there just yet. The takeaway message is that H+ appears on the right side. So adding H+ makes the reverse reaction go faster and taking it away makes the forward reaction go faster. (See Blanden.Adam's post)
When all this math boils down you're left with a relationship that says at any given pH the relative ratios of the different species in the acid base reaction will be fixed. It's worth noting that we are talking about weak acids here. Strong acids like hydrochloric acid and sulfuric acid dissociate completely. In reality they just have very very very large equilibrium constants, but that has the effect of making most everything in this discussion null and void for those types of acids.
Now let's talk pH for a second. pH is a measure of the concentration of H+ ions in solution. Actually it is the negative of the base 10 logarithm of the chemical activity of hydrogen ions in solution.
Water is both an acid and a base. That's what makes it interesting. Water is constantly dissociating to form H+ and OH-.
H<sub>2</sub>O <-----> H<sup>+</sup> + OH<sup>-</sup>
This happens to an equilibrium point, which for pure water happens when the pH is 7. We can push this equilibrium around, for instance if we remove OH- from solution say by precipitating magnesium hydroxide, then the reaction shifts towards the right and more water dissociates to form more H+ and OH- until the equilibrium condition is satisfied again. In this example, it causes pH to rise because we have put more H+ into solution.
Back to carbonic acid, we can start to think about what happens in any situation with that system. Say we add base, sodium hydroxide. We've added more OH- to the situation so we push water back towards the left and the pH goes up. Had we added that much base to pure water it would have gone up by the same amount as the amount of base we added, but in this situation there is carbonic acid.
Remember that relationship that said that at any given pH the ratios of the forms of carbonic acid would be constant? Well, as the pH goes up, the amount of carbonic acid has to go down. Think of it this way, there is now less H+ in the water, so that carbonic acid reaction is pushed towards the right. So is the reaction from bicarb to carbonate. That means some of those protons on those carbonic acid and bicarbonate molecules are coming off. Those H+ ions are coming off into solution in the water. That replaces part of the H+ that we lost when we added that hydroxide, so the pH doesn't rise nearly as much as it would have without the carbonic acid in solution.
And this is what we call a buffer. Any time you have a mixture of a weak acid and its conjugate base (carbonic acid - bicarbonate) AND (bicarbonate - carbonate) then the act of keeping those ratios right tends to resist changes in pH.
I have just been challenged to a match of "The Very Hungry Caterpillar" game. I have to go for now. I will return with the rest of this story.
disc1
09/16/2013, 07:59 PM
Does this mean the higher your KH the more basic your water becomes?
Not necessarily. Let's be clear and state that since we are talking about seawater, we are mostly only going to talk about the carbonic acid species. Almost all of the alkalinity in our tank water comes from bicarbonate and carbonate. Only a very small portion from borate and others so we for the most part ignore them.
So alkalinity is like the number of ions of bicarbonate and carbonate, but pH is about the ratios between them. In the alkalinity game, bicarbonate counts once because it has only lost one proton and so only one basic group but carbonate counts twice since it has lost two protons. So if we don't change the amount actual amount of carbonate species in the water but only raise the pH, then since there are less protons around the equilibrium between bicarb and carbonate gets shifted to the right. The sum of bicarb and carbonate is the same, but now more of it has shifted into carbonate form. Since carbonate counts twice we end up having a rise in alkalinity.
Remember the bicarbonate to carbonate equilibrium:
H<sub>2</sub>CO<sub>3</sub> <-------> HCO<sub>3</sub><sup>-</sup> + H<sup>+</sup><-----> CO<sub>3</sub><sup>2-</sup> + 2 H<sup>+</sup>
Now lets consider a tank with a relatively high pH already and we add bicarbonate to the tank as baking soda. When we put the baking soda in the tank we have altered the ratio of bicarbonate to carbonate in the water. We've added more bicarb. We've added to the left of the equilibrium reaction above. So to get to equilibrium, that reaction has to go to the right and make some carbonate to get back to the right ratio. In the process of making the carbonate, we lose protons to the water. We are adding alkalinity, but we are lowering pH.
Now in practice this effect is small. In a reef tank under normal conditions, roughly 80% to 90% of the carbonic acid equivalents are in bicarbonate form. So when we add more bicarb, not so much has to shift over to carbonate to make things balance back out.
If on the other hand we add sodium carbonate, now we are adding to the right side of the bicarbonate -> carbonate equation. To get to equilibrium it has to shift to the left. That means making bicarbonate which consumes a proton. Since we are removing protons from the water, pH goes up. But the change is much more drastic with sodium carbonate. Since the natural state of things is 80% to 90% bicarbonate, most all of that sodium carbonate has to be converted to bicarb. This means we are going to consume a lot more protons than we would have liberated had we added the same amount of alkalinity with bicarbonate. So the pH rise from using carbonate is much more dramatic than the pH drop using bicarbonate.
Finally there is CO2. It has several equilibria going on at once. First of all, the amount dissolved in the water is in equilibrium with the amount in the air. And the part that is dissolved in the water is in equilibrium with carbonic acid.
The neat way to think about equilibrium reactions (and one of the secrets to oh the second or third test or so in any good Orgo class) is to see equilibrium like a balance. If you take away from one side, then you suck the reaction that way. If you have a reaction that has water as a product, then removing water from the reaction "pulls the equilibrium that way" and forces you to make more product. Same with these protons in solution. If you add them to any of these chemical equations, then you shift the equilibrium and force the reaction to go towards the other side. When you have a complex set of equilibria, you can look at the whole chain the same way. Removing any one species sucks all the other equilibria towards where that species is in the chain. Adding any one species pushes all the other equilibria outwards from it.
So when CO2 enters the water, most all of it ends up quickly as bicarbonate. So does it add to alkalinity? The answer is no, that process is alkalinity neutral. In a relative sense, you cannot create or destroy alkalinity, you just move it around. Since neither CO2 nor water have any basic groups as they came to the reaction, no set of products can have a net gain or loss in alkalinity.
How can that be? Think about it, we added carbonic acid with two protons, it lost one to go to bicarbonate which adds one unit of alkalinity, but the proton you lost consumes one unit of alkalinity. So the net process is alkalinity neutral.
If we remove CO2 from the aquarium we do that by lowering the concentration of CO2 in the air around it. Thus CO2 leaves the system as the equilibrium between dissolved CO2 and atmospheric CO2 gets pulled towards gaseous CO2. So more carbonic acid goes to CO2 and water. The CO2 leaves the game. That loss of carbonic acid pulls bicarbonate to carbonic acid to get that back to equilibrium. That takes a proton and that raises pH. This further pulls a bit of carbonate over to bicarbonate taking another proton. That whole cycle continue until the CO2 is back in equilibrium between dissolved and gaseous with the net product of raising the pH without affecting the alkalinity at all.
disc1
09/16/2013, 08:07 PM
It seems to indicate that it does. I definitely need to increase my pH from the neutral pH tap water to somewhere around 8.5 This requires a tenfold decrease in H+ ions.
You see that all the time. pH doesn't add and subtract like other concentrations. We're talking about buffer solutions here.
The take home message is that it doesn't matter one bit what the pH of the freshwater you start with is. When you add it to the saltwater that buffer system is going to set the pH.
Let's say we have some RODI water at pH 5.5. The pH is low, likely due to a little dissolved CO2. When CO2 dissolves in pure fresh water it has to make bicarbonate and carbonate to satisfy those equilibria and that makes protons which lower the pH dramatically. The difference is that there is no buffer there to soak it up. It all stay in solution so the pH changes by a lot for just a little bit of CO2.
So we add that to some seawater at pH 8.2 and what happens? Have we changed the ratios of carbonic acid to bicarbonate to carbonate? By a tiny tiny amount, the tiny amount of CO2 that was dissolved in that freshwater. So not a lot of shifting is going to happen. The pH will now be 8.2 or really really close to it. The buffer system "soaks up" the pH change by shifting things around from carbonic to bicarbonate to carbonate or back the other way. Whatever has to happen to resist the pH change. There's more than enough of the buffer in seawater to more than soak up what you are adding when you top off or make up your salt mix.
disc1
09/16/2013, 08:14 PM
Ooops Caught a mistake on post 8. Last sentence of the first paragraph after the second chemical equation.
This happens to an equilibrium point, which for pure water happens when the pH is 7. We can push this equilibrium around, for instance if we remove OH- from solution say by precipitating magnesium hydroxide, then the reaction shifts towards the right and more water dissociates to form more H+ and OH- until the equilibrium condition is satisfied again. In this example, it causes pH to rise because we have put more H+ into solution.
That should say causes the pH to drop because we have put more H+ into the solution.
Sorry. I type fast sometimes and it is too late to edit.
Jimmy54
09/17/2013, 01:06 AM
Wow, you must be a teacher, and if not, you should be one ;)
The best and easy to understand reading/lecture about the relationship between CO2, H2O and pH I have ever seen :thumbsup:
Jimmy
JonJon82
09/17/2013, 07:20 PM
You guys are great. I was going absolutely INSANE waiting for my tank to cycle. It’s not much fun growing things that take an electron microscope to see. At least I know I'm figuring stuff out now instead of staring at an empty tank.
This is the framework my mind has set up to understand, manipulate the pH of a solution of an aquarium.
Follow me, if pH is a measurement of H+ ion concentration, then pH is a measurement of proton concentration. So, different fish live at different pH, means different fish live in different proton concentrations.
If you apply this definition of pH, then decreasing my pH simply means adding protons via bicarbonates, and increasing my pH simply means removing protons by neutralizing them carbonates.
Do you think this conceptualization of pH being a "concentration of protons" will work in the aquarium world?
I've also been thinking about the last post I read from you guys about equilibrium. The equilibrium equation for carbonic acid, its salts, and water seem to hold true at any pH. How do you tell your buff to hold this pH instead of the last one?
In theory, I know how to raise and lower my pH. What I don't know in theory is how to maintain that pH.
Cant wait to read your guys posts from today! If you answered any of these questions obviously then I'll find them
We got a long way to go this is only pH!
disc1
09/17/2013, 08:22 PM
You guys are great. I was going absolutely INSANE waiting for my tank to cycle. It’s not much fun growing things that take an electron microscope to see. At least I know I'm figuring stuff out now instead of staring at an empty tank.
This is the framework my mind has set up to understand, manipulate the pH of a solution of an aquarium.
Follow me, if pH is a measurement of H+ ion concentration, then pH is a measurement of proton concentration. So, different fish live at different pH, means different fish live in different proton concentrations.
That's about the size of it. pH is nothing more than a measure of H+ concentration.
If you apply this definition of pH, then decreasing my pH simply means adding protons via bicarbonates, and increasing my pH simply means removing protons by neutralizing them carbonates.
Assuming the pH is above about 8 or so. Below that, when there are more protons around, the bicarbonate gets pushed towards carbonic acid and that takes up a proton.
Do you think this conceptualization of pH being a "concentration of protons" will work in the aquarium world?
I cartainly hope so, it is the only conceptualization that there is.
I've also been thinking about the last post I read from you guys about equilibrium. The equilibrium equation for carbonic acid, its salts, and water seem to hold true at any pH. How do you tell your buff to hold this pH instead of the last one?
So you tell it what pH to hold by setting the ratio of carbonic to bicarbonate. For instance, when I want a carbonate buffer in the lab I go to a table called appropraitely "The Buffer Tables." Without the tables I could use the Henderson-Hasselbalch equation (google that) to calculate it myself, but it is just so handy to have a table.
I look up what pH I need in the table and it tells me to use a certain ratio of carbonate to bicarbonate. I mix up those two at the prescribed ratio and I get a solution buffered to my desired pH.
In the aquarium we have another player in the game and that is CO2. If we add or take away CO2 we push the ratios of the other species around and create a situation with a lower or higher pH. If I add CO2, then everything pushes towards carbonate and pH goes down. If I take away CO2 then everything gets pulled towards carbonic acid and the pH goes up.
In theory, I know how to raise and lower my pH. What I don't know in theory is how to maintain that pH.
The simplest thing to understand about pH in seawater is that the primary player is CO2. pH, alkalinity, and CO2 are caught in this equilibrium dance and the three are constantly moving to maintain equilibrium. So if you know any two of those three, you can calculate the third.
So with alkalinity the same, increasing CO2 decreases pH and vice versa.
With CO2 held constant, increasing alkalinity increases pH and vice versa.
So to maintain a certain pH, keep alkalinity and CO2 constant and pH will stay the same.
Cant wait to read your guys posts from today! If you answered any of these questions obviously then I'll find them
We got a long way to go this is only pH!
The real story is don't worry too much about it. In freshwater systems pH can range from the 5's to the 9's. That's a huge range and most fish only like a narrow window in there. Tetras like the 6's and soft water. Cichlids like up to pH 9 and really hard water.
And without much buffering capacity to "soak up" changes in pH, the pH of a freashwater system can get way way out of whack.
In seawater the situation is much different. Here we have a strong carbonate buffer system holding our pH stable for us. We are pretty well locked into that high 7's to low 8's range. You have to get pretty drastic to get pH to go outside of that range in seawater. So pretty much whatever you do animals in seawater are going to have a pH that they are comfortable with.
JonJon82
09/18/2013, 04:12 PM
While I'm learning this...My first real experiment is going to be how effective a calcium carbonate, magnesium carbonate solution is holding a pH level. These two compounds sound intriguing since "in theory" it would establish a pH equilibrium and help keep it there by making soft water harder. Anyway, its going to be interesting to see how well it performs. If its working correctly, neutral water should increase in pH without wild pH shifts with the added Calcium and Magnesium.
Also thanks for finally hammering it in about how big a role C02 plays. Since it controls the rate at which carbonic acid is made, C02 is the biggest donator of protons in your tank
I'm still reading your last posts until we move on haha. I cant get this stuff out of my head
You know what would be really interesting is knowing the maximum rate at which C02 can be dispersed through the opening in your tank. Then you could measure the C02 of your water. If you do this initially and take the same measurements in a week and use subtraction you could call that the
"Weekly C02 Bioload" of your tank. Theres actually alot of interesting things you can do with a C02 measurement. You could even predict what day in the week your pH is likely to go up and down.
.
blanden.adam
09/18/2013, 05:13 PM
You know what would be really interesting is knowing the maximum rate at which C02 can be dispersed through the opening in your tank. Then you could measure the C02 of your water. If you do this initially and take the same measurements in a week and use subtraction you could call that the
"Weekly C02 Bioload" of your tank. Theres actually alot of interesting things you can do with a C02 measurement. You could even predict what day in the week your pH is likely to go up and down.
.
No you couldn't. With good flow, this you will establish an equilibrium with the surrounding air on the order of tens of minutes to hours depending on the size of your tank, flow, and surface area available for equilibration (opening a window and watching your pH rise will tell you this).
I'm also not quite sure I know what you mean when you want to see how well calcium carbonate and magnesium carbonate will hold pH. They are really quite insoluble until you get down in the sub 7.6 range, and are the basis for most substrate used in marine tanks and virtually all live rock.
Jimmy54
09/18/2013, 05:36 PM
While I'm learning this...My first real experiment is going to be how effective a calcium carbonate, magnesium carbonate solution is holding a pH level.
To get a Calcium "carbonate" solution, you're going to need a calciumreactor which infact produces calcium and bicarbonates.
If you don't have one, the only other way left to get calcium, magnesium and (bi)carbonates into the water
is by using CaCl2, MgCl2, NaHCO3 and/or Na2CO3.
Or a lot of patience, waiting until the CaCO3 and MgCO3 is dissolved :)
JonJon82
09/18/2013, 05:45 PM
I just poured bleach into tap water and watched the pH go up. My first experiment. And I knew what was happening! Which made watching the reaction satisfying. That's crazy. The water went from 7.6 to higher than I can measure.
This solution is really designed to keep a pH between 7.8 and 8.5 so I'm going to start at a neutral fresh water pH. The added hardness from the magnesium and calcium wouldn't benefit salt water since salt water is already hard enough and keeps pH well.
Just found C02 measuring table
Jimmy54
09/18/2013, 06:20 PM
OK... I guess you mean hardness as in dGH (total hardness)? instead of alk in dKH or meq/l ?
JonJon82
09/19/2013, 04:28 PM
I couldn't find an alkalinity tester or any good buffering compounds to test at the stores my area. It's a shame that the chemists at API, who's products are sold in almost all fish stores, have really let us down. A) They don't sell anything that tells you what IS actually in the bottle. It's like they don't want you to understand what you're doing. Most bottles say they include "patented" things. No, just no. A chemist does not want to drop a bunch of molecules into his tank, until he knows what they are.
I'm going to need another day of conceptualizing acid base reactions and how they apply to my tank.
But yeah I think magnesium carbonate is a good way to add carbonates into fresh water, thus raising the pH of ordinary tap water until it stabilizes itself at pH 8.5. I dont know how the magnesium and calcium ions affect pH yet. I'm not there yet. But from what I have read, tap water is usually low in K and Ca when its pH is low.
Also learned measuring alkalinity is usually done as a titration measurement, which is impossible without that equipment, and are usually found as part of a swimming pool test kit.
Your alkalinity is going to stop your water from becoming too acidic, so that means its going to be neutralizing a growing number of acidic particles in your tank as C02 levels increase. If you could measure your C02 levels, you could conceptualize how saturated your water is with C02. Based on that you can determine how fast carbonic is being made. Based on your alkalinity reading you could also conceptualize how quickly your pH will rise. If your alkalinity is low it means your carbonate buffer is being strained, and pH will start to go down. I'm also hoping other ions in the water like Ca and K help slow down pH shifts when the carbonic buffer is strained.
This is hard but fun I can see why you guys like it. I dont want to be afraid of the water in my tank.
Also interesting fact, through thermal energy fluxuations at the subatomic level, H20 molecules are able to split a part, which makes H30 and OH, but almost instantaneously (1 picosecond) they come back to form water. Water neutralizing itself to stay neutral. How quickly water breaks apart and reforms, its almost as if it hadnt happened at all
Jonathan
bertoni
09/19/2013, 04:54 PM
Potassium and calcium have no effect on pH. A bit of baking soda will increase the buffering of tap water and will raise a low pH. I'm not sure what you're trying to do, though. Magnesium carbonate isn't very soluble in water.
Jimmy54
09/19/2013, 05:08 PM
0.106 g/l to be exact.
disc1
09/19/2013, 05:10 PM
I couldn't find an alkalinity tester or any good buffering compounds to test at the stores my area. It's a shame that the chemists at API, who's products are sold in almost all fish stores, have really let us down. A) They don't sell anything that tells you what IS actually in the bottle. It's like they don't want you to understand what you're doing. Most bottles say they include "patented" things. No, just no. A chemist does not want to drop a bunch of molecules into his tank, until he knows what they are.
I'm going to need another day of conceptualizing acid base reactions and how they apply to my tank.
But yeah I think magnesium carbonate is a good way to add carbonates into fresh water, thus raising the pH of ordinary tap water until it stabilizes itself at pH 8.5.
Won't happen. You can't dissolve any more magnesium carbonate above a pH in the low 7's. Don't remember exactly the number. In an acidic solution it will dissolve and pH will rise but that decreases the solubility. Eventually you hit a balance point where you have as much in solution as is soluble at whatever pH you've reached and it all stops. That pH where it stops is going to be acidic or just barely above neutral.
I dont know how the magnesium and calcium ions affect pH yet. I'm not there yet. But from what I have read, tap water is usually low in K and Ca when its pH is low.
They don't. They are both INCREDIBLY weak acids. I mean they are to weak acids like a newborn baby is weak compared to weak football players. Technically they can play the acid role in a reaction, but they are not going to affect pH in any noticeable way.
Also learned measuring alkalinity is usually done as a titration measurement, which is impossible without that equipment, and are usually found as part of a swimming pool test kit.
Or one of our aquarium test kits. I like Salifert myself, although I hear lots of good things about the newer RedSea kits.
If you don't have an alkalinity kit, then stop RIGHT NOW and go get one before you play with any chemistry in your aquarium. You don't want to go at this blind.
Get calcium and magnesium while you're at it.
Your alkalinity is going to stop your water from becoming too acidic, so that means its going to be neutralizing a growing number of acidic particles in your tank as C02 levels increase. If you could measure your C02 levels, you could conceptualize how saturated your water is with C02. Based on that you can determine how fast carbonic is being made. Based on your alkalinity reading you could also conceptualize how quickly your pH will rise. If your alkalinity is low it means your carbonate buffer is being strained, and pH will start to go down. I'm also hoping other ions in the water like Ca and K help slow down pH shifts when the carbonic buffer is strained.
That buffering action works both ways. People always seem to get it into their head that high alkalinity goes with high pH. That's not it. High alkalinity goes with STABLE pH. Higher alkalinity makes it harder to change pH. If you've got high pH and high alk, then it's hard to bring that high pH down. If you've got low pH and high alk then it is just about as hard to bring that low pH up.
The buffering action resists changes in pH in any direction.
This is hard but fun I can see why you guys like it. I dont want to be afraid of the water in my tank.
Also interesting fact, through thermal energy fluxuations at the subatomic level, H20 molecules are able to split a part, which makes H30 and OH, but almost instantaneously (1 picosecond) they come back to form water. Water neutralizing itself to stay neutral. How quickly water breaks apart and reforms, its almost as if it hadnt happened at all
Jonathan
Not just an interesting fact, but the basis of what we are talking about. Water splits into H+ (which goes to another water molecule to make the H3O we usually just say H+ or a proton) and OH- at a rate just barely faster than they get together. So there is always a little bit of the water broken up at any given time. The equilibrium constant is 10^-14. That means that:
[H+] * [OH-] / [H2O] == 10^-14
That's the equilibrium expression. The [] brackets mean the concentration of whatever is inside them. The concentration of the water is always taken as 1 so the equation reduces to
[H+] * [OH-] = 10^-14
So if we have absolutely pure water. The purest of the pure. Then we know that all of the H+ came from water molecules and so we know that for every H+ there is also an OH-. So the concentrations of the two MUST be equal.
If the two are equal then let's call that concentration X.
X * X = 10^-14
or
X^2 = 10^-14
So X = the square root of 10^-14 which is 10^-7
So the H+ concentration in pure pure water is 10^-7 M.
The pH of a solution is defined as the negative of the base ten logarithm.
log (10^-7) = -7
Take the negative and we get 7
Which is the pH of pure pure water.
disc1
09/19/2013, 05:12 PM
0.106 g/l to be exact.
Depends on the pH.
And the temperature and pressure. But mostly on the pH
JonJon82
09/20/2013, 05:17 PM
OK I ordered the kits. Went into my fish budget a little bit but what goods a dead fish. Now I can measure the carbonate, bicarbonate ions. Dispelling the myth that high alkalinity doesnt mean high pH helped A LOT. You have no idea how much. And also about the role of the Ca and K ions.
I don’t quite understand equilibrium yet, because there are quite a few players involved, but I’m glad its there. It means if you’re increasing H+ ions you’re decreasing OH- ions. I’m beginning to understand this as a reactant to product ratio. Where if you load the reactant, like hydrogen ions, you’re going to get a lot more product, carbolic acid. Not a lot of carbolic acid is going to be turned back into water, because there are less OH molecules for it to react with.
See I’m finally getting the concepts
Sodium bicarbonate. Theoretically it can be used to stop waters ability to resist the pH shift you intend to make. In the right proportion it won’t be harmful to fish. I also realize that when your alkalinity is low, water becomes very vulnerable to wild pH shifts. Which is why if you don’t know what you’re doing you’re pretty much dead meat.
I cant wait to see waters buffer getting strained by the change in alkalinity. So thats how I'm going to observe alkalinity first, water and sodium bicarbonate
This is going to be fun
blanden.adam
09/20/2013, 05:23 PM
"Not a lot of carbolic acid is going to be turned back into water, because there are less OH molecules for it to react with."
Well, not quite. You are right that when you add protons, you shift the equilibrium to more carbonic acid, but then that in turn equilibrates with dissolved CO2 and H20, and then that equilibrates with the CO2 concentration in the air. So much of the formed carbonic acid will turn back into water and CO2 because CO2 isn't forced to build up in the aquarium, it's free to equilibrate with the air, which has a huge volume relative to your aquarium.
JonJon82
09/20/2013, 06:31 PM
Since the water is trying to reach equilibrium with the air above it, does that mean the water will expel C02 and bring in C02 at different rates depending on how much C02 is in the tank? Thats how I understood part of what you said
I was just looking back and you're right water molecules don't stick to C02 its free to come and go as it pleases as long as it doesnt turn into carbonic acid on the way up out of your tank
JonJon82
09/20/2013, 07:42 PM
That would also mean most of the gases in your tank are 02 and N2 and very little C02. Which means your tank will be getting rid of C02 at the surface faster than it takes up 02 in order to get back to equilibrium.
Then respiration turns 02 into C02 which can then diffuse into the atmosphere OR react with water to form carbonic acid.
The proteins in your fish cells react to the ion concentration in the water. They move around. That's why any of this is important to the aquarium world or people who keep fish.
I dont plan on messing with the pH of tank water once the fish are in it. That's something you cant recommend anyone do, especially when lives are at stake. I just have this thing where I believe every type of fish should live in the pH of the water they evolved in. My reason is, if you know evolution, adaptations are actually beneficial genetic mutations. The fish haven't found a way to live happily in a new pH. That would take a genetic mutation. No fish you buy has adapted to the new pH it just IS adapting, which means its cells are acting differently than they would in their natural waters.
And that is why I want to know how to set a pH.
Jonathan
JonJon82
09/20/2013, 08:52 PM
Slowly introducing a fish into the pH of its natural habitat is not going to harm the fish. In fact they are going to start behaving more like themselves. Im not a hippy but that sounds really hippy doesnt it
I going to learn how to count the concentration if ions in moles next. That's how I'll know how much acid or base to add. Anyway I got a week until the kits arrive.
Can you ever have alkalinity that's too high? It seems like carbonic acid determines how high your alkalinity can go
disc1
09/20/2013, 09:21 PM
Slowly introducing a fish into the pH of its natural habitat is not going to harm the fish. In fact they are going to start behaving more like themselves. Im not a hippy but that sounds really hippy doesnt it
I going to learn how to count the concentration if ions in moles next. That's how I'll know how much acid or base to add. Anyway I got a week until the kits arrive.
Can you ever have alkalinity that's too high? It seems like carbonic acid determines how high your alkalinity can go
You don't want to be adding acid and base to try to push pH around. You'll knock everything out of balance in the end. Really, the natural buffering action of seawater is going to hold you in a good range. As long as you've got good alkalinity you should be able to ignore your pH unless you see some specific problem.
You can definitely have alkalinity too high. You'll crash the calcium out of solution.
Carbonic acid doesn't define how high the alk can go, those two are independent. Carbonic acid decides what the pH will be at whatever alkalinity level.
bertoni
09/20/2013, 09:31 PM
Well, I think actually carbonic acid can allow the alkalinity to go higher without a precipitation event, as in a calcium reactor, although I am rather tired right now. I'm not quite sure what you meant by the penultimate sentence.
dkeller_nc
09/21/2013, 08:54 AM
Guys, you might find this example useful to visualize what David is saying about the independence of alkalinity/acidity and pH.
Note that this is a very simplistic example compared to the extraordinarily complex ionic interactions in the chemical soup of seawater, but it might help.
As David has noted, there's a class of chemicals termed "strong acids". The one most commonly encountered by the hobbyist are hydrochloric (often sold as "muriatic acid"). Also as David noted, the equilibrium between the molecule HCl and the dissociated ions H+ and Cl- is very strongly shifted to the right. So strongly, in fact, that the concentration of the parent molecule HCl is effectively zero in aqueous (water) solution:
HCl <===> H+ + Cl-
So here's the point about the independence of pH and acidity/alkalinity. If you were to put a pH meter probe in a 0.1 Normal HCl solution (fairly dilute), you will get a reading below pH 2. If you were to do the same thing with a 10 Normal solution (100 times more concentrated), you will also get a pH reading below 2.
Yet, if you needed to add the correct amount of sodium hydroxide to these 2 acid solutions to completely neutralize the acid and bring the pH to 7.0 (essentially making salty water), you would need to add 100 times as much sodium hydroxide to the 10 N HCl solution as you would the 0.1N HCl solution.
By convention, the amount of neutralizing substance that you must add to the HCl solutions to bring it to pH 7 is the "acidity" of the solution. So the 10N HCl solution is way more acidic than the 0.1N HCl solution, even though their measured pH would be almost identical.
The situation is a bit more complicated with weak acids like acetic (commonly called vinegar), but the visualization is the same - pH tells you whether the solution is acidic or basic (below or above pH 7), but not the degree to which the solution is acidic or basic.
blanden.adam
09/21/2013, 09:10 AM
Guys, you might find this example useful to visualize what David is saying about the independence of alkalinity/acidity and pH.
Note that this is a very simplistic example compared to the extraordinarily complex ionic interactions in the chemical soup of seawater, but it might help.
As David has noted, there's a class of chemicals termed "strong acids". The one most commonly encountered by the hobbyist are hydrochloric (often sold as "muriatic acid"). Also as David noted, the equilibrium between the molecule HCl and the dissociated ions H+ and Cl- is very strongly shifted to the right. So strongly, in fact, that the concentration of the parent molecule HCl is effectively zero in aqueous (water) solution:
HCl <===> H+ + Cl-
So here's the point about the independence of pH and acidity/alkalinity. If you were to put a pH meter probe in a 0.1 Normal HCl solution (fairly dilute), you will get a reading below pH 2. If you were to do the same thing with a 10 Normal solution (100 times more concentrated), you will also get a pH reading below 2.
Yet, if you needed to add the correct amount of sodium hydroxide to these 2 acid solutions to completely neutralize the acid and bring the pH to 7.0 (essentially making salty water), you would need to add 100 times as much sodium hydroxide to the 10 N HCl solution as you would the 0.1N HCl solution.
By convention, the amount of neutralizing substance that you must add to the HCl solutions to bring it to pH 7 is the "acidity" of the solution. So the 10N HCl solution is way more acidic than the 0.1N HCl solution, even though their measured pH would be almost identical.
The situation is a bit more complicated with weak acids like acetic (commonly called vinegar), but the visualization is the same - pH tells you whether the solution is acidic or basic (below or above pH 7), but not the degree to which the solution is acidic or basic.
This is incorrect. pH absolutely tells you how acidic or basic a solution is -- in fact that's what it is precisely what it's meant to tell you. What you are describing is simply that pH is a log scale, so every 1 pH unit you move up or down represents a decrease or increase in H+ concentration by 10. So a 0.01 N HCl solution will have a pH of 2, a 0.1 N a pH of about 1, a 1 N a pH approaching 0, and once you go beyond that things get more complicated because water starts to play a more active role.
Alkalinity on the other hand, is a measure of buffering capacity. In other words, 2 solutions with a pH of 8.0 can have alkalinities of 6 dKH and 12 dKH respectively. The 12 dKH solution requires twice as much acid to drop it's pH to 4.5 (the convention used for alkalinity measurement), even though their starting pH is the same because the concentration of buffers (mostly carbonate/bicarbonate) is twice as high as the 6 dKH sample. They have the same concentration of protons and are thus equally acidic/basic, but the 12 dKH sample requires much more acid or base to change the pH.
The bottom line:
pH = how acidic or basic a solution is
alkalinity = how concentrated your buffer is/what's the buffering capacity
dkeller_nc
09/21/2013, 11:25 AM
Yep, not nearly enough coffee this morning, and a strong acid or base may not have been the best example since all of it is essentially ionized and the hydronium ion concentration is therefore a direct function of the overall concentration.
But the point is as you put it - pH is insufficient information to describe the amount of an acid or base required for neutralization.
You're probably going to hate this ;), but in the chemical industry the term "acidity" is used rather loosely to describe the amount of base necessary to neutralize it, regardless of whether that usage is precisely accurate in scientific terms.
blanden.adam
09/21/2013, 01:03 PM
Hah, you're right. I do hate it ;).
disc1
09/21/2013, 01:29 PM
Acidity in that way is the inverse of alkalinity. You're still measuring buffer capacity, just measuring it going the other way.
JonJon82
09/21/2013, 02:19 PM
You guys are such great teachers. What I learned just this week is staggering. Btw, this is also career training for me (If I play the right hand I can get a job at a store that sells fish) while I go to college.
JonJon82
09/21/2013, 02:33 PM
Anyway my goal now is to fill a container with 29 gallons of water. Let it stabilize with the atmosphere. Perform a chemical reaction that moves the pH out of the range of 7.4 to 7.8 to a new range of 6.8 to 7.2 I'm putting everything we learn into figuring this out. I'll probably have questions while I design the experiment. Finally assigning values is what I'll be working on first. "Mole Money, Mole Problems" was the name of my chemistry project. LOL
blanden.adam
09/21/2013, 03:03 PM
Anyway my goal now is to fill a container with 29 gallons of water. Let it stabilize with the atmosphere. Perform a chemical reaction that moves the pH out of the range of 7.4 to 7.8 to a new range of 6.8 to 7.2 I'm putting everything we learn into figuring this out. I'll probably have questions while I design the experiment. Finally assigning values is what I'll be working on first. "Mole Money, Mole Problems" was the name of my chemistry project. LOL
But why? You have a buffered solution, you add acid the pH goes down. What precisely are you trying to learn?
JonJon82
09/22/2013, 12:25 AM
If sodium bicarbonate can be used to strengthen the alkalinity of water in cases when it is low. pH changes can be swift and catastrophic when dealing with fresh water species, which I also hope to care for. That would be the only reason an aquarist would place NaHC03 into their tank. Which would be the point to any experiment with that molecule.
Addition:
Strengthen alkalinity without changing pH in a fresh water system with NaHC03.
Jimmy54
09/22/2013, 02:13 AM
Well, there is a difference between freshwater and freshwater tanks.
e.g. For a rift-valley fish tank like a Malawi or Tanganyika tank, I would think you can use small quantities of chemicals like NaHCO3 - Na2CO3 - MgCl2/SO4 6H2O/7H2O - CaCl2 2H2O and a few others to achieve pH 8.4 and dKH/dGH 12 -14.
For an amazone-like fishtank it's a different story; pH around 6.5 and dKH 5/6 requires a different approach to stabilize these parameters.
dkeller_nc
09/22/2013, 06:58 AM
Acidity in that way is the inverse of alkalinity. You're still measuring buffer capacity, just measuring it going the other way.
Yep, that's exactly what's meant by the less-than-precise common use of the "acidity" term in my industry.
dkeller_nc
09/22/2013, 07:12 AM
If sodium bicarbonate can be used to strengthen the alkalinity of water in cases when it is low. pH changes can be swift and catastrophic when dealing with fresh water species, which I also hope to care for. That would be the only reason an aquarist would place NaHC03 into their tank. Which would be the point to any experiment with that molecule.
Addition:
Strengthen alkalinity without changing pH in a fresh water system with NaHC03.
Hmm - Jimmy54's answer is the correct one. If you want hard, alkaline water for African Cichlids, one typically adds carbonate/bicarbonate/sulfate salts of magnesium, calcium, sodium and potassium (so-called "African cichlid buffer"). Depending on the manufacturer, borate salts may also be included. That gets you a stable, buffered pH around 8.5.
But one thing that David/Adam hasn't discussed in this thread yet is that different buffer salts only buffer in a specific range of pH's. For example, borate is a suitable buffer salt to use when you want a buffered solution in the pH 8-10 range. But it won't have much use if you need the pH to be stable in the 5-7 range.
JonJon82
09/22/2013, 02:11 PM
Bicarbonate is a salt, so basically when you add baking soda you’re adding a salt. Some of the salt is going to be cancelled out by free hydrogen atoms. You have to add enough baking soda to cancel out the free protons plus some extra if you just want to increase alkalinity/maximize your alkalinity for your pH. But if you put too much, you’ll of added to much salt and the pH will change. At any one time, there would be more salt available than free protons.
Molecules that can raise alkalinity and pH . (pH 8.4 and dKH/dGH 12 -14)
NaHCO3
NaH2C03
Na2C03
MgCl2/SO4 6H2O/7H2O
CaCl2 2H2O
You guys helped a lot. I thought adding bicarbonate would turn a solution acidic instead of basic. Making water more acidic seems much more dangerous because you’re decreasing alkalinity to let the pH go down. Since bicarbonates make your pH go up, I don’t see any other way to make water more acidic than adding an acid. The salts you guys are talking about are commonly known as “cichlid salts” and you can mix your own by mixing magnesium sulfate - MgS04 7H20 – with sodium bicarbonate NaHC03. I HAVE to test this because some people say it works some people say it doesn’t.
I’ve always tried to get the alkalinity to follow the pH in my head, but its easier to let the pH follow the alkalinity. This would stop the pH from swinging while you’re adjusting it.
JonJon82
09/22/2013, 02:35 PM
I'm just having the hardest time figuring out exactly how the ratio of carbonic acid, carbonate, and bicarbonate change the pH. Its in the notes of the thread somewhere. Gotta keep reading until it jumps out at you
Jonathan
Addition:
I can see equilibrium as a predictable ratio of ions that are present at any one time. It is difficult to imagine that at any one moment there are literally trillions and trillions of ions that exist at any one time. And one picosecond later they disappear. I found the equation for hydrogen concentration, there are alot of constants you dont need to measure because they remain the same at standard temperature and pressure. One day I hope to be able to solve that equation
disc1
09/22/2013, 03:44 PM
Molecules that can raise alkalinity and pH . (pH 8.4 and dKH/dGH 12 -14)
NaHCO3
NaH2C03
Na2C03
MgCl2/SO4 6H2O/7H2O
CaCl2 2H2O
NaHCO3 and Na2CO3 can raise alkalinity. Their effect on pH depends on where the pH was to begin with. For example, if you ad bicarb to a solution that has a pH of 9 it will drop the pH but if you add it to a solution with pH 6 it will raise the pH.
NaH2CO3 --- There's no such molecule.
The magnesium and calcium salts you listed have no effect on alkalinity or pH.
blanden.adam
09/22/2013, 03:49 PM
JonJon,
Of the molecules you listed:
NaHCO3/Na2CO3 -- the basis for the marine aquarium buffer system
NaH2CO3 -- Not a real molecule
MgCl2/SO4 6H2O/7H2O and CaCl2 2H2O -- Only affect pH in that they can alter the stability of their conjugate anions, they do not directly effect pH by themselves and their effects are typically very small.
To solve for the concentrations of the buffering species you need to buffer at a certain pH given the pKa of that transition (which you can look up) is given by the Henderson-Hasselbalch equation:
pH = pKa + log([A-]/[HA])
where:
pH is the pH of the solution
pKa is the pKa of the transition from HA <--> H+ + A-
[A-] is the molar concentration of the basic form of your buffer
[HA] is the molar concentration of the acidic form of your buffer
This is, of course, for simple buffers and not complex mixtures, but it will give you a place to start.
Jimmy54
09/22/2013, 04:18 PM
Only if necessary MgCl2/SO4 and CaCl2 is added to raise the dGH (general or total hardness) in fresh water.
JonJon82
09/23/2013, 10:49 AM
Calculations:
About .001 tsp/gallon NaHC03 to increase water's Alkalinity by 10ppm
About 1 mL/gallon C2H402 - Acetic acid - to lower water's pH by .3
These calculations would be EXTREMELY helpful to anyone with basic tap water who like to keep FW species like angelfish, or anyone with low alkalinity in their water. Only problem is, theres no way to measure .001 teaspoons. Im still working on that one.
Bicarbs act differently depending on pH that was helpful.
Thank you for the equation it was a great starting point.
JonJon82
09/23/2013, 03:15 PM
Found a way to measure the concentrations of carbonate, bicarbonate, and C02 given the pH.
Hydronium Constant for pH 7.6: 2.5x10-8 moles/liter
Carbonic Acid Equilibrium: 1.7x10-3 moles/liter
Bicarbonate Dissociation Constant: 1.70x10-3 moles/liter
Carbonate Dissociation Constant: 4.69x10-11 moles/liter
C02 Constant: 3.4x10-2 mol/liter x atm)
Total Concentration of Dissolved Inorganic Carbon (DIC): .290
I already found the equations I need to calculate the relative concentrations of molecules at equilibrium, but I need a little algebra help (I failed Algebra 3 Trig twice)
The equation has a numerator and denominator, but to the right of the fraction is "x DIC"
What do I do with this "times DIC" in the equation.
Addition:
I'm plugging these values into the Bjerrum plot equations which I can't write here. I saw how solving these equations can be used to make the graphs. But I dont plan on making a graph
Jonathan
JonJon82
09/23/2013, 06:07 PM
Alrighty guys I want to thank you for your posts. I know what I'm doing now. Once I figure out this Henderson-Hasselbalch equation, I'll be able to make my own buffer solutions out of sodium carbonate and sodium bicarbonate solutions. I'll just experiment until I find a solution that sets the correct pH. I think I'll make a solution for 6.8 to 7.2 and one for 7.8 to 8.3 I'll find it just by experimenting!
Thanks again omg I cant thank you enough. You made me a pH and carbonic master! I'm going to feel much more comfortable working with tanks. It really warms my heart to see you guys come together like this to help someone who wanted to learn more.
Jonathan
If I start asking algebra questions in here this thread will never end believe me
JonJon82
09/23/2013, 06:58 PM
The testing should go fast I have lots of little 1 gallon jugs to test in. When I started this thread I was literally on the water molecule.
Thats the great thing about science, you can either know the the answer or test UNTIL you find the answer. I would kill for your guys knowledge in chemistry.
I have 13 open word documents with notes from you guys
JonJon82
09/23/2013, 07:41 PM
Does anyone else want to know what David is constructing?
disc1
09/23/2013, 10:42 PM
There's really no reason you have to set the ratio right off the bat. It's not that you can't change the pH of a buffered solution, it's just that it takes more acid or base to do it than it would without the buffer.
So what you do in the lab is this. You hit a buffer table to get close. There are a million out there that give you the ratios for different pH. Or you plug into the HH equation and calculate it yourself, it isn't hard at all with a good scientific calculator.
Once you have something close, you simply titrate it with acid or base until you get the pH you want. With strongly buffered solutions it might take a surprising amount of acid or base, but eventually you get the pH to where you want it and as long as that is within the buffering range for that buffer (pKa + or - about 1 pH units) then it will hold that pH.
That works fine in a bottle or a test tube, but it doesn't work that way in the aquarium unfortunately. The reason comes back to the CO2. Remember that one end of our equilibrium set was CO2 gas going to CO2 dissolved. That's an open ended equilibrium there. The atmosphere is so huge and the tank is so small that you can't even begin to try to think you are going to make any change in the atmospheric levels of CO2 based on the fish tank taking on or giving up CO2.
So in the end no matter what you do in the tank carbonate / bicarbonate wise, as long as it is open to the atmosphere the CO2 is going to drive the equilibrium in the carbonate buffer system. You can set it wherever you want, and the atmospheric level of CO2 will set it back for you.
There is a nice equation in one of Randy's articles that lets you calculate pH from CO2 and total alkalinity. You don't need to know what forms the alkalinity is in, just the total carbonate alkalinity. But from the HH equation you can calculate the ratio from pH.
Since the total alkalinity in the tank doesn't change (on the really short timescale anyway I'm not talking about coral consumption) then it is the pH and CO2 that are locked into this firm relationship. The CO2 sets the pH, the pH sets the ratios of carbonate to bicarbonate and then the whole system tend to try to hold that state and resist any changes.
Where the alkalinity fits into that relationship it only shows up as total alkalinity. What it means is that for higher alkalinity, it takes more CO2 to get the same pH drop.
disc1
09/23/2013, 10:47 PM
Does anyone else want to know what David is constructing?
Yeah, mine got wiped out in a two week power outage last winter. It's been a busy year and I'm having to rebuild things kid-proof this time. I redid some plumbing and rebuilt my sump. I've got to build a little box behind the stand to hold all the equipment and electricity and then I can filler up and start cooking this rock and get ready to cure it.
The old lady expects it to be as beautiful as it once was right off the bat. She's gonna be sore when it just looks like a glass box full of rocks for six months. :hmm2:
JonJon82
09/23/2013, 11:21 PM
You're right, there's no way around the Ideal Gas Law is there. Pressure and volume inside will always equal the pressure and volume outside. I'm willing to guess that it doesnt take very long for water to equalize with the atmosphere. If thats true than alkalinity will always be set by the atmosphere. That's pretty interesting. Carbonic acid is always going to go back to the same rate of C02, which will make the concentration of carbonate/bicarbonate drop. You can see that in a way your alkalinity evaporates. But even if the concentration of carbonates to bicarbonates falls (your alkalinity falls) to align with the atmosphere, the ratio of carbonates to bicarbonates should be stable if your alkalinity is within the right range. That would mean you can have a stable pH so long as you monitored your alkalinity. Thank you for telling me how you do it in the lab. My room is about to be turned into a lab.
Thats horrible about your tank I would be devasted. 2 weeks without power, thats almost a life changing experience. That was smart finding an "old lady" that liked fish!
disc1
09/24/2013, 08:54 AM
You're right, there's no way around the Ideal Gas Law is there. Pressure and volume inside will always equal the pressure and volume outside. I'm willing to guess that it doesnt take very long for water to equalize with the atmosphere. If thats true than alkalinity will always be set by the atmosphere. That's pretty interesting. Carbonic acid is always going to go back to the same rate of C02, which will make the concentration of carbonate/bicarbonate drop. You can see that in a way your alkalinity evaporates. But even if the concentration of carbonates to bicarbonates falls (your alkalinity falls) to align with the atmosphere, the ratio of carbonates to bicarbonates should be stable if your alkalinity is within the right range. That would mean you can have a stable pH so long as you monitored your alkalinity. Thank you for telling me how you do it in the lab. My room is about to be turned into a lab.
Thats horrible about your tank I would be devasted. 2 weeks without power, thats almost a life changing experience. That was smart finding an "old lady" that liked fish!
You're close. Look back a few posts we talked about this one. The CO2 doesn't affect alkalinity at all. When it does make bicarbonate, it also leaves a proton. So there's no net change. You make a unit of alkalinity and you consume one. So the whole process doesn't change alkalinity at all.
The only thing that it changes is the pH. You set the alkalinity, and atmospheric CO2 sets the pH.
JonJon82
09/24/2013, 01:08 PM
Ok I’m going to take a closer look. There’s still a few things about carb/bicarb I need to go back and look at.
It’s just so odd that through the chaos of interactions of atoms a balance is reached. You said that in one of your first posts, equilibria is a thing of beauty. It really is.
JonJon82
09/24/2013, 02:55 PM
So its your pH that determines if bicarbonates tend toward acid or carbonate. If your solution is basic, bicarbonate will go toward acid to balance the pH and go back to neutral. If your pH is acidic, the bicarbonates go wards carbonates for the same reason, balancing the ions in order to return to neutral.
Is this right so far
blanden.adam
09/24/2013, 03:59 PM
Nope, it's the opposite.
Lets say you have 10 mM bicarbonate in solution in water at it's 2nd pKa (NaCOOOH <--> Na2COO, pKa = 10.3). At the pKa, 50% of the molecule is protonated and 50% is deprotonated, so at a pH of 10.3, I'd have 5mM bicarbonate, and 5mM carbonate. Now, let's say I add 4 mM of a strong acid to this solution, say HCl. The acid will immediately ionize and protonate the Na2COOO molecules, leaving us with 9mM bicarbonate and 1mM carbonate with a pH of about 9.3 (you can confirm this with the henderson-hasselbalch equation. Also note that in the absence of the bicarbonate, the pH would have dropped to between 3-4, so it buffered significantly.)
Now lets say I do the converse and add 4mM of a strong base, say sodium hydroxide. Now the free hydroxides will deprotonate the available bicarbonates. Now you have 1mM bicarbonate, 9mM carbonate, and a pH of ~11.3.
So, raising the pH of a solution shifts the balance toward carbonate, dropping the pH of the solution toward bicarbonate/carbonic acid, and there is no tendency to go back to neutral.
Now, if you take a solution of any pH and then add a buffering compound, what pH it tends to depends on the species of buffer already present in the solution and the mixture of species you add. Can get quite complicated, but these also are predicted by more complicated iterations of the same henderson-hasselbalch equation.
JonJon82
09/24/2013, 04:00 PM
Thanks for that explanation
If your solution is basic, bicarbonate will go toward carbonate. If your solution is acidic, bicarbonate will go toward acid. Let me think about that for a day
JonJon82
09/25/2013, 11:01 AM
On test #4 I got the results I wanted. I added 100mL of a bicarbonate solution (1/2tspn pure baking soda to 1000mL H20) to one gallon of normal tap water. The pH did not change, but the alkalinity must've gone up at least a little since alot of that bicarb was turned into carbonate. Bicarb deprotonated one ion but picked up one. In small doses the deprotonation of the bicarb must be absorbed by the carbonate, but for every bicarbonate that turned into carbonate, a carbonate turned into bicarbonate. Only the concentration of the two increased. If a concentrated sodium bicarbonate solution is added, the pH skyrockets. This is a VERY powerful compound to the aquarist and should only be used by experts, like me.
JonJon82
09/25/2013, 02:55 PM
I think I understand now that your carbonates and bicarbonates will move toward a one to one ratio in equilibrium. If you want to visualize it, it's almost like carbonate and bicarbonate are playing catch with a proton. Nothing really changes. I'm not sure if I'm right about this part, but a weak solution of sodium bicarbonate should move toward carbonate or toward C02 and H20. Why it depends on pH is the confusing part and I'm going to stay away from that for awhile. It was helpful to look back at your posts and realize that of the equilibrium reaction, whatever part of it you effect, the interactions move away from bicarbonate to balance the change. If all the bicarbonate deprotonates, then it becomes carbonate, when carbonate picks up the proton it becomes bicarbonate. They switch places but nothing changes.
Does anyone know what happens to the concentration of your alkalinity as it moves toward equilibrium? It would seem like the concentration would go down. Some of the protons will form carbonic acid, and this carbonic acid will turn back into C02 and water. Which would indicate concentration is reduced.
Jonathan
disc1
09/25/2013, 03:05 PM
Alkalinity stays the same. You can kind of think of alkalinity units in the same way you think about conservation of mass energy. You can change forms, but you can't make any new ones or destroy any old ones. I'm talking chemically of course, your corals actually consume them and precipitate them into skeleton. But if two molecules of bicarbonate (2 units of alkalinity) get together and for a while exist as a carbonic and a carbonate, then there's still 2 units of alkalinity (the carbonate).
They're not going to move to a 1:1 ratio though. In your tank it is somewhere between 80 and 90% bicarbonate. The ratio depends on the pH. The pH gets set by the seemingly limitless and unchanging atmospheric CO2. So that's what sets the buffer ratio.
JonJon82
09/25/2013, 05:35 PM
If you add bicarbonate, like I wanted to do, its going to either pick up a proton and turn back into C02 and water, or give up a proton and turn into carbonate. This action is going to happen until bicarbonate is back to the ratio with carbonate set by the atmosphere above the water. What this means is, if you try to add bicarbonate without try to effect the pH it isnt going to do much. The equilibrium is going to bring it back.
I'm going to do it =( made a sad face..I really wanted, what I considered to be, an unmovable pH when I get a job taking care of salt and fresh water tanks. If the alkalinity in the fresh water tanks gets low, I'll just do a water change...how boring.
You all reawakened in me an insatiable appetite to learn. Even my daily routines I approach differently now. I printed out this entire thread. They are my eternal study materials. I don't want to lose you guys by ending this thread. But time to move on to ammonia nitrite and nitrate.
=*(
Thank you for changing me.
JonJon82
09/25/2013, 05:40 PM
Oh and I'm very sorry David must wait 6 months for his rock to bake. It's going to be tough
JonJon82
09/25/2013, 06:27 PM
Funny joke. I went up to my dad and said, "Do you know what's in our water? Hydrogen, hydroxide, carbonic acid, carbonate, bicarbonate" and he asked "Should we be drinking it?"
true story
blanden.adam
09/25/2013, 07:31 PM
If you add bicarbonate, like I wanted to do, its going to either pick up a proton and turn back into C02 and water, or give up a proton and turn into carbonate. This action is going to happen until bicarbonate is back to the ratio with carbonate set by the atmosphere above the water. What this means is, if you try to add bicarbonate without try to effect the pH it isnt going to do much. The equilibrium is going to bring it back.
I'm going to do it =( made a sad face..I really wanted, what I considered to be, an unmovable pH when I get a job taking care of salt and fresh water tanks. If the alkalinity in the fresh water tanks gets low, I'll just do a water change...how boring.
You all reawakened in me an insatiable appetite to learn. Even my daily routines I approach differently now. I printed out this entire thread. They are my eternal study materials. I don't want to lose you guys by ending this thread. But time to move on to ammonia nitrite and nitrate.
=*(
Thank you for changing me.
I remember you saying earlier that you are trying to get the fish job while going to school. If you are really interested, why not just take some chemistry? Then you'd have more, better study materials, and learn more than you ever could on a fish forum.
disc1
09/25/2013, 09:24 PM
I remember you saying earlier that you are trying to get the fish job while going to school. If you are really interested, why not just take some chemistry? Then you'd have more, better study materials, and learn more than you ever could on a fish forum.
For sure, equilibrium is one of the first concepts in freshman chem right after electron structure. Even if you don't like it and don't want to take any more you still knock out a science elective.
JonJon82
09/26/2013, 06:46 PM
I'm sure the biology has some nice chem classes. Im house sitting next week and I'll have 3 whole quiet days alone. I'm just going to open my chemistry and start learning how to do the math. The math is the part that depresses me the most. I went to bed depressed woke up depressed but now I'm right back it. It really comes down to a simple thing really, you want to know. Anway, it turns out the human body has a way to deal with ammonia, but fish dont have that, because they excrete the whole molecule. And this molecule is really hard to get out of water because of its hydrogen bonds. Its the first pyramid shaped molecule Ive seen. Going to read some of this new aquarium book I got. Take care!
blanden.adam
09/26/2013, 07:51 PM
Well, not really.
Humans and fish handle ammonia in much the same way internally, the difference is just in the excreted form. Both humans and fish conjugate ammonia to an amino acid called glutamate to form glutamine. In fish, the glutamine travels to the gills where it is converted back into glutamate and ammonia, and the ammonia is released. In humans, the glutamine travels to the liver (and to a lesser extent the kidneys), where it is converted back to glutamate and ammonia, the ammonia is then converted to urea, which is then filtered by the kidneys and excreted in the urine.
This allows us to conserve our carbon (we don't excrete the glutamine), and gives us a self-regulating, pH neutral osmotic dieresis that allows us to eliminate nitrogen at times that are convenient for us without suffering the effects of ammonia building up between bathroom breaks -- afterall, we aren't constantly bathed in a medium that can carry the ammonia away for us like fish are.
And yes, trigonal pyrimidal molecular geometry is pretty sweet. That protonatable lone-pair on the nitrogen is the crux of why nitrogen chemistry is so cool and fundamental to life. I would make mention though, as a former biochemistry major who lived in both worlds, you aren't likely to learn much chemistry in your biology classes -- the place to learn chemistry is through your chemistry department :)
sunning
09/27/2013, 01:37 AM
It's easy, just Carbon dioxide (toxic) and water to form carbonic acid, am I right?
bertoni
09/27/2013, 01:21 PM
Yes, carbon dioxide will dissolve in water and some will form carbonic acid.
JonJon82
09/29/2013, 08:10 PM
I'll be back Friday with chemistry questions
JonJon82
10/03/2013, 03:04 PM
Well I was alone in the country for three days and didn't get to any chemistry lol. I was too busy looking at plumbing for filtering multiple side by side tanks. In all honesty everything we've gone over in terms of chemistry and alkalinity and acids and bases, and equilibrium, has really made me very confident in making water quality decisions. There are different ways to hook multiple tanks up together and that's what I'm looking at now. Maybe this will start a new thread. This information in this thread is so good for carbonic acid that it should probably stay as carbonic acid
Jonathan
JonJon82
10/07/2013, 03:11 PM
OK I feel my mind turning back into mush. Need another chemistry topic to study that relates to aquariums.
Does anyone have a good one?
bertoni
10/07/2013, 11:23 PM
Mush? Hmm, how about ORP? That'll drive almost anyone crazy. Have you looked through this page:
http://reefcentral.com/forums/showthread.php?t=2127046
JonJon82
10/08/2013, 06:07 PM
Yes! Thank you these articles look like just what my brain needs. And it looks like an understanding of pH isn't complete without knowledge of ORP. Anyway, I want my fish to have plenty of anti oxidants to clear the toxins from their body. (kidding)
Thanks again!
Looks like brain candy for the week =)
Interesting fact, a fishes stomach is about the size of their eye, according to this book
blanden.adam
10/08/2013, 07:38 PM
Good luck with your quest to understand ORP. To me it's still a box of magic that gives you a number I can't relate to a physical process. Sure, it's a number, and empirically it's the ORP the water generates between two metal electrodes, BUT WHAT DOES IT MEAN!?!?!?!?
JonJon82
10/08/2013, 09:35 PM
lol maybe its time to move on to quantum entanglement
I always wanted to know where those electrons go when they jump orbitals
bertoni
10/08/2013, 09:47 PM
Quantum entanglement is very strange. I can't find an application in reefkeeping, though. :)
JonJon82
10/12/2013, 12:14 PM
UNLESS unless you're moving a couple thousand miles away and want to bring your fish with you. Then you could just teleport them...but...would they be the same fish? Anyway, you guys are so great, I'm going to give you my number xxx.xxx.xxxx OMGerd you gave out your number. Well, there are worse things I could do. We could probably get alot more learning done that way. Plus you could share you're own personal updates that I'd love to hear. Have a good week guys
billsreef
10/13/2013, 09:39 AM
Interesting fact, a fishes stomach is about the size of their eye, according to this book
Don't know what book your reading, but that one is really a myth that has it's roots in trying to prevent people from overfeeding their fish ;)
JonJon82
10/14/2013, 03:16 PM
Thanks for dispelling that myth. The internet is full of conflicting information, but I come here because the information is unparalleled. And the way you accept new people is very special. Without this site I wouldnt be confident applying to a LFS that I'm going to do this week. BTW I was at my neighbors chili cook off and they were letting me drink and I was so jubilant I gave out my phone number. I would really rather not have it posted on a public forum for very long.
JonJon82
10/14/2013, 04:19 PM
nm, that was from a library book. Can't believe a library book can lie too you. You were right on, the author was talking about over feeding
billsreef
10/14/2013, 06:02 PM
I removed your number for you ;)
I bet that book was aimed at hobbyist, no shortage of myths in that market.
eyesinthedrk
10/15/2013, 10:30 AM
since orp was brought up....
on another non fish specific forum, i was in a debate with a fellow about weather orp was a test for oxygen. my stance being that testing for an oxidant/redox ratio was not the same thing as testing for oxygen. with O2 being the most abundant yet least powerful oxidant in the water, the results wont tell you how much oxygen you have
kind of like trying to find out how many corvettes are in your town by counting all the sports cars
he pointed to an industrial patent where orp was being used to detect oxygen in a fermentation process trying to replicate aerobic conditions. i agreed that ORP reading 0 oxidants means zero O2. but in a reef tank there are just too many variables.
the counter point was that all OPOs are removed from the system once they pass through the carbon, any other super oxidants are rapidly consumed or burned up. meaning that in the end the oxidant side of the ORP measurement is all O2. if it wasn't the fish would literally melt before our eyes
Randy's article clearly states otherwise. as does this thread i found when checking the facts of this claim. that validated both my points
http://reefcentral.com/forums/showthread.php?t=2317076&highlight=orp
the part that throws me off is that based on some of his arguments he clearly has a better grasp of chemistry than i do, and the conviction of his argument is giving me doubts.
am i missing something?
dkeller_nc
10/15/2013, 02:29 PM
Good luck with your quest to understand ORP. To me it's still a box of magic that gives you a number I can't relate to a physical process. Sure, it's a number, and empirically it's the ORP the water generates between two metal electrodes, BUT WHAT DOES IT MEAN!?!?!?!?
Ha - to me, ORP is lot less byzantine than a multi-component buffering system. Or memorizing the Kreb's cycle to spit it back onto a biochem test.Funny how even in the scientific community, certain specific things appeal to certain specific people. I find the solution calculations for general relativity in the limit of r --> 0 (i.e., a black hole) to be fascinating. Most of my colleagues get stuck at the expression of tensor mathematics in curvelinear coordinates. ;)
blanden.adam
10/15/2013, 05:22 PM
Ha - to me, ORP is lot less byzantine than a multi-component buffering system. Or memorizing the Kreb's cycle to spit it back onto a biochem test.Funny how even in the scientific community, certain specific things appeal to certain specific people. I find the solution calculations for general relativity in the limit of r --> 0 (i.e., a black hole) to be fascinating. Most of my colleagues get stuck at the expression of tensor mathematics in curvelinear coordinates. ;)
Here's to scientists who understand what I do not. Especially to those who write wonderful open-source computer programs in languages I understand or with intuitive user interfaces to calculate the things that I cant on my own! :beer:
JonJon82
10/15/2013, 08:23 PM
What's stranger ORP or the reason why species, such as the blue ribbon eel (Rhiomuraena quaesita) can never adapt to captivity, deciding better to starve to death?
bertoni
10/15/2013, 11:54 PM
I don't see why ORP reading zero implies zero oxygen, in general. The level of oxygen probably would be close, but someone probably could concoct a system with a steady state that allowed a bit of oxygen. I'm not talking about a reef tank or anything like it, though. This quote from the ORP article is relevant:
So, there is a subset of oxidizers and reducers that are actually capable of reacting with each other, and moreover for interpretation of ORP, in impacting an ORP electrode. Consequently, a single ORP value measured for a given aqueous solution may not correctly describe the relationship between any given pair of redox species in the solution.
On a more practical note, carbon (as we dose it) won't remove all the oxidizing agents. It won't remove most of them, I'd guess, actually.
This quote is relevant, too:
The unfortunate circumstance with ORP, however, is that we do not have a good understanding of the redox active species in seawater and marine aquarium water.
You could look at the sections on pH and ORP, too. ORP rises when pH drops, but a drop in pH can happen with no change in the oxygen level. In fact, pH dropping generally is a sign of carbon dioxide buildup due to respiration, and respiration consumes oxygen. :)
bertoni
10/15/2013, 11:55 PM
Oops, another point that is mentioned in Randy's article is that ORP is not in equilibrium in our tanks, which is an important issue.
billsreef
10/16/2013, 05:36 AM
On a more practical note, carbon (as we dose it) won't remove all the oxidizing agents. It won't remove most of them, I'd guess, actually.
Indeed, if carbon removed any significant quantity of oxidizing agents we would see ORP crash from use of it ;)
Oops, another point that is mentioned in Randy's article is that ORP is not in equilibrium in our tanks, which is an important issue.
Yup, ORP is not an equilibrium type thing.
dkeller_nc
10/16/2013, 06:40 AM
I don't see why ORP reading zero implies zero oxygen, in general. The level of oxygen probably would be close, but someone probably could concoct a system with a steady state that allowed a bit of oxygen. I'm not talking about a reef tank or anything like it, though.
It doesn't - Oxidation Reduction Potential doesn't necessarily imply any particular oxygen level in the general sense. That's why it's sometimes called a "confounded" variable - there are many possible combinations of different molecular species and concentrations that can give one the exact same ORP reading.
One of the things that often confounds new chemical engineers is not realizing that very few oxidation reactions in aqueous solution depend on the dissolved molecular oxygen concentration at all.
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