Science Buddies: "Ask an Expert"

I am helping my 5th grader with a science fair project based on the Swimming in Acid, Ocean acidification project. We modified it a bit -- we decreased the pH down to 6.8, rather than 7.5, and are using phenol red as an indicator to see what the pH is, because we don't have access to a pH meter (or time to order one before the fair). The big question is this: we set up our jars today using 6 different types of shells, using ocean water with phenol red as a control, and ocean water with phenol red and vinegar as the acidified ocean water. We added vinegar until the ocean water turned yellow (which is supposed to happen at 6.8 ). I am not concerned if it is less than 6.8, just that it is at least 6.8 or less. So we set up bottles with shark eggs, clam shells, scallops, oysters, moon snails, and crab legs. Here's the thing: most of the bottles' acidified ocean water remained yellow after adding the shells, but with the crab legs, the water turned back to pink to indicate the pH went back to a much more alkaline 8.2. Our question is why? What is the chemistry happening with the crab legs that is re-neutralizing the water and actually making it go back to alkaline? I have a background in chemistry so I would understand it if you get into equations....you don't have to simplify for my sake. Is there something inherent in crab shells that contributes alkalinity so as to spend the acid concentration? It looks like the moon snails are slowly changing pinkish as well. My theory is that it must have something to do with air? Because I left out a sample of the acidified ocean water to open air and that sample has also changed to pink. The crab legs initially floated and released bubbles after the acidified ocean water was added. Moon snails (unbroken) also may entrap air. But what is the reaction with air that causes the pH to go back up? I should add that we went to great effort to make sure the bottles had no headspace, but this was particularly difficult for the moon snails and the crab legs. So does this mean that the CO2 is going out of solution into the air? Should we reopen the bottles and add more vinegar to push the pH back down? If we do, then the pH of those bottles may be way less than the remaining samples, compromising comparability. Perhaps add it, diluted, drop by drop until the pH *just* changes (i.e., until we just see the yellow color change)?

Update: I posted this to another thread on the Life Sciences but they suggested posting here. I want to add that over the last 24 hours, the shark eggs, scallops, and oysters all have reverted to a pH greater than the original (all are pink now in the phenol red rather than yellow). The pink is VERY strong in the crab legs (indicating a pH of 8.2, while the other shell types are indicating pH's more like 7.6-8.0). After looking at some carbonate chemistry webpages, I do think that maybe the CO2 is possibly leaving solution and escaping to a gas phase. I'm thinking because my ocean water was cold (I kept it in the basement and I live in the NE), when we added the vinegar maybe it took time for the CO2 to come out of solution due to temperature, and perhaps we stopped adding vinegar to get the pH down to 6.8 too soon? Or could it actually be that the acid is already reacting with the shells and dissolving a small amount of carbonate, and upon addition of the acidified water the CO2 was produced??

Update: Since we didn't hear from anyone here on AAE, and his hypothesis is that shells exposed to acidified ocean water will lose mass, and we are trying to see which ones will lose the most mass, we decided to go ahead and reopen the bottles and add more vinegar. We added just enough so that the color changed from peachish/pinkish (depending on the starting pH) to yellow (6.8ish). I suspect the ocean water buffered against the addition of water, or the shells have started dissolving and the CO2 releases into the headspace and the pH goes basic again. Not sure which, but both are related. We will wait a day or two and update again.

5thgrade_boy - Just wanted you to know that several experts are reading this, but are as bewildered as you. It's easy to understand the pH and indicator issues, but I don't have an understanding of the makeup of the sea creature shells. This is a case where the chemistry expertise in the Physical Science forum understand the acid/base indications using phenol, but can't explain what is changing the pH. Perhaps that is what we were hoping the Life & Earth Science forum could cover. One other possible assistance, if you can't locate a pH meter, perhaps you can borrow a pool test kit from a friend, neighbor or even your school. That, at least, might get you some better readout of actual pH levels. Good luck, I've been following both your posts since they went up trying to understand what is happening.

Rick Marz

Rick, Thanks for your comments. We are actually using phenol red from a pool test kit, so that is what is so handy about us using it...we are able to watch the change from 6.8 to 7.2 to 7.6 to 8.2 rather easily and dramatically.

Actually, I think I am beginning to understand what is happening:

We added additional vinegar this afternoon, just until we reached a pH of 6.8 (i.e., just until, with a "drop" using a medicine dropper), the solutions turned yellow. In only a few hours, we are back to the colors prior to addition of the additional vinegar. What I think is happening is that for the Swimming in Acid project, although the title suggests that the shells are in acid, they are in fact still in alkaline/neutral ocean water, as they advised using a pH of 7.5. I advised my child to use a pH of 6.8 since the decreased pH will increase the kinetics of the reaction, and we actually might have a chance of measuring a mass difference (I looked in other posts of people who have tried the experiment as is, and they noted no measurable difference between the mass before and after leaving the shells in 7.5 ocean water after about a month. Someone on AAE posted a way to calculate what mass difference you could expect, and it was not possible for her to measure using a scale that only measured to the 0.1 g.) We also only have a balance that measures to the 0.1 g. The only thing is, the person who conducted the experiment before did not measure the pH at the end of the experiment. That is why I thought it would be nice to use phenol red so we wouldn't have to keep opening/closing our sample bottles to find out what was happening.

So back to what we are seeing...I think that 6.8 is, for the shells, hugely different than 7.5 and the calcium carbonate is, in fact, dissolving in the 6.8 acidified ocean water. No matter how careful we are about trying to close the bottles to eliminate any headspace, we are ending up with some headspace, albeit a small amount, in all sample bottles (controls have no headspace). So I hypothesize that as the shells are dissolving, the carbonate is converting to CO2 and is creating a headspace. This drives the pH back up to 8.2 as the CO2 leaves the solution.


CH3COOH +CaCO3  CO2 + Ca(C2H3O2)2
Acetic acid and calcium carbonate react to form carbon dioxide and calcium acetate

The calcium carbonate dissolves in the water, so the mass of the shells in the vinegar should decrease and they should get a little smaller. The carbon dioxide is released as a gas.
The ocean will never become as acidic as acetic acid solution, but the lower pH in this activity allows the effects to be seen over a shorter amount of time. For a longer-term experiment that more accurately models ocean chemistry, we should place shells in soda water (water with CO2 bubbles) for days to months and observe the effects. (It would be fun to add phenol red, too to see when things are happening more precisely.)

But if I am correct, doesn't this beg the question: If CO2 is in fact released into the air as a result of the oceans becoming more acidic, does this mean that the reaction itself could actually be driving itself??? Maybe upon first inspection, but when you consider all other factors going on (temperature, organic matter decomposition, and many other processes going on in the ocean), one must know you can't extrapolate what is occurring in a complex system like the ocean compared to a small bottle with shells in it and nothing else. In any case, the ocean and this process does seem to buffer itself.

Any thoughts?

Is there a way to contact the writer of the project directly? The sample bottles are back to pHs of 7.2 and up, in the same way as a couple of days ago prior to the addition of additonal vinegar. We are tempted to add more vinegar but are not sure we are on the right track.

Hi 5thGrade_boy,

You're on the right track. The short answer is that your experiment is working as it should. Don't worry if your bottles don't stay acidic - they won't because the shells, etc. will dissolve until equilibrium is reached. That's what is "supposed" to happen

I've been involved in some geochemical research about ocean acidification, so I'll try to provide some context from a seawater chemistry point of view. Since you've said you're comfortable with chemistry, I'll be somewhat more technical than I normally would for a K through 5 project!

When people talk about "ocean acidification", they are not talking about the ocean getting to a pH < 7 (i.e., becoming acidic on the pH scale). Instead, they are talking about the pH of ocean water decreasing by a few tenths of a pH unit. Seawater will still be alkaline, but it will have a lower pH. Although 0.1 pH units may seem like a small number, decreasing the ocean's pH from its current value (8.1) to 8.0 corresponds to a 30% increase in H+ concentration! And, seawater pH has only varied by ~0.2 pH units over the past 800,000 years.

In the ocean (and, in reality, for your experiment), it isn't as simple as CO2 gas being released as shells dissolve. There is a complex equilibrium between the concentrations of H(+), HCO3(-), H2CO3, and CO3(2-). [Note: I'm using () to denote the charge on a charged species.] The concentration of each of these species depends on pH. At the ocean's current pH of 8.1, for example, ~9% of the dissolved inorganic carbon is present as CO3(2-), 90% as HCO3(-), and 1% as H2CO3. Decreasing ocean pH increases the concentrations of H(+), HCO3(-), and H2CO3 while decreasing the concentration of CO3(2-).

The key question is whether or not seawater is saturated with respect to calcium carbonate or magnesium carbonate, two minerals that the shells of many sea creatures are made of. If the water is undersaturated, then shells may dissolve. As the concentration of CO3(2-) goes down, seawater that was formerly saturated or supersaturated can become undersaturated. Seawater's saturation state also depends on temperature and pressure, so there is a "horizon" in the ocean called the carbonate compensation depth (CCD for short). Decreasing the ocean's pH moves CCD to shallower depths, exposing creatures that previously lived happily in supersaturated waters where their shells were stable into undersaturated water where their shells can start to dissolve. At the ocean's pH, the concentration of CO3(2-) drops extremely fast as pH decreases, so the CCD can potentially shallow by quite a bit.

That's the basic seawater chemistry story. Here are my thoughts about your particular experiment:

Gastropod shells are mostly calcium carbonate, which, as you pointed out, reacts with acid to produce Ca(2+), H2O, and CO2. The CO2 will either outgas (if there is enough of it for bubbles to nucleate) or disassociate into carbonate, bicarbonate, carbonic acid, and hydrogen ions. Yes, the CO2 that is produced does act to lower pH.

When you put the shells in the acidic water, which is undersaturated with respect to calcium carbonate, they begin to dissolve until the concentrations of dissolved species reaches equilibrium. It doesn't surprise me that the acidic solutions quickly became alkaline - remember that pH is a log scale, and reaction (dissolution) rate depends on how far out of equilibrium the solution is. A pH 6.8 solution will result in a much larger and faster response than the pH 7.5 solution. Once equilibrium is established, the mass of the shells will no longer change with time.

There is the issue of atmospheric equilibration, but I suspect the dominant effect in your experiment is from dissolution of carbonate, not outgassing of CO2 from the solution.

This has been a long post - let me know if I can clarify something or if you have additional questions. I'll get an email when you respond to this thread, so I'll get back to you quickly.

Terik,

Thanks so much for your reply!

So now he just completed his local school fair and is moving on to the District Fair this weekend. He did very well and tied for 3rd place in his school!

Just to summarize, after only 15 days of leaving the shells in the acidified ocean water and "re-injecting" more vinegar every 2 days, he was able to measure a drop in the mass of oysters, as opposed to the scallops, clams, and moon snails.

To explain what was happening in the bottles further, we need some input. The pH of the moon snails had swung the most over the course of 2 days, going from about 6.8 to about 8.0-8.2 every 2 days. However, after I think it was the 3rd vinegar addition, we went on Winter Break and could not bring the bottles with us (we have a small car and a large family and went to visit my elderly mother). Anyway, when we got back, the pH of the moon snails bottles were only up to about 7.8, definitely not 8.0-8.2. And the other bottles, instead of going up to about 7.6, they held at about 7.2. None remained 6.8, however.

To explain this, I think this may have something to do with the alkalinity of the water in the bottles. The moon snails contained the most mass per bottle, at about 42 g, as opposed to the other shells. They also took up way more volume within the bottles. So I'm thinking that there was less ocean water in those bottles so whatever was happening chemically, we were able to see it at a greater concentration, maybe? And that with every addition of vinegar, perhaps the source of alkalinity was becoming spent? Unfortunately we had to weigh the shells in time for the fair so we couldn't go further and stopped at 15 days. But what is actually causing the bottles to turn back to pink/pinkish orange? Is it simply the release of CO2 into the headspace of the bottles, or is it something else? What about the conversion of whatever is contributing alkalinity? Another thing is, what stumps me if it is indeed just the release of CO2, is why doesn't the reaction in some ways feed itself? In other words, could the concentration of CO2 in the headspace ever get so high from the dissolving of the shells that it could re-acidify the water? But we're not seeing that from the pH indicator fluid changing from yellow to pink as the reaction progresses from 6.8 to 7.2 and higher.

This has been a wonderful project to go over, but I am somewhat restricted about what my son can relate, since indeed he does not have the benefit of any chemistry background and concepts like "equilibrium" and "balance of reactions" is way over his head, as he is hardly able to understand the reactions. However, I did a bit of research on this online and even ocean geochemists admit it is a very complex problem to consider. WoodsHole did a similar experiment as ours and unexpectedly got results that suggested that crabs and lobsters actually *GAIN* mass and calcify their shells in the presence of increased CO2 (!). They too got that oysters lost the most mass, along with soft clams.

Any insights as to what is actually happening are most appreciated!

Hi 5thgrade_boy,

Neat results! And, I'm glad that your son's project is doing so well.

The varying results for oysters, scallops, clams, and moon snails is probably, in part, related to variations in shell composition. Based on some poking around I did, it looks like oyster shells are mostly calcite while clams and moon snail shells are dominantly aragonite. Calcite and aragonite are both calcium carbonate, but they have different crystal structures: the atoms are arranged differently. I see mixed data about which dissolves more readily, but calcite versus aragonite probably plays a role in what you're seeing. The level of magnesium in the shell plays a roll; magnesium ions often substitute for calcium. The presence or absence of the organic-based protective layer that often coats bivalves could contribute to variations. Another factor is probably at work is the surface area to volume ratio of the shells. The area over which sea water can react with the carbonate minerals in shells is controlled by the shell's surface area. But, the amount of material in the shell is controlled by volume. As a result, shells with high surface area to volume ratios will break down more quickly, all other factor (e.g., shell composition, organic coatings, etc.) being equal.

I can't say for sure why the pH didn't increase as much over Winter Break. I suspect you are on the right track about using up the alkalinity in the shells. The material available for reaction is largely controlled by the surface area of the shell. There will be some thickness (for a given timescale) over which the acidity can penetrate the shell, but once that material is consumed, the reaction will not be able to continue until new material is exposed (e.g., by scraping off the outside of the shell, breaking it, etc.) or equilibrium is perturbed in some other way. Another factor is that the reaction rate rate will rapidly decrease as the system approaches equilibrium.

As far as the mass/volume per bottle goes, if you added the same volume of vinegar to each container, of ocean water, then the concentration of H+ in the moon snail bottle could conceivably be higher. However, I thought that the oysters changed in mass, but the moon snails did not. Or am I misunderstanding something? If the moon snails didn't measurably lose mass, but the oyster shells did, then this line of thinking isn't particularly helpful for explaining the difference.

The pH indicator turns from yellow to pink when the pH increases from 6.8 to 8.2. The pH changes because as the shells dissolve, they release carbonate ions into the water. The increase in carbonate ions leads to a decrease in the concentration of hydrogen ions. Since pH is the -log(concentration of hydrogen ions), the pH goes up. There will be ocean water - head space exchange as the gas in the head space reaches equilibrium with the water. But, I suspect that in your situation the main issue is not the ocean-gas equilibrium but rather the equilibrium between the different carbonate species dissolved in the water. That's where the action is happening.

As far as explaining this on a 5th grade level goes, I think a good way to go about it is to use the vinegar and baking soda reaction, which is very visual, as an analogy. Baking soda isn't calcium carbonate, but the concepts still apply. Essentially, ocean acidification is all about "using up" either the acid or shells (or, vinegar and baking soda). Try doing an experiment with a lot of vinegar and a little bit of baking soda. After the bubbles and foaming subside, the baking soda will be gone and the liquid will still be acidic. You could use phenol red to show this. In this case, you "used up" all of the baking soda, but there is still left over vinegar. Then, do an experiment with a lot of baking soda and a little bit of vinegar; after the reaction is over the vinegar will be gone (water will remain) but baking soda will be left over and the pH of the remaining liquid will be alkaline (again, you can use phenol red to see this). The idea of things getting used up by reactions is at the heart of this project, as is the idea of having one of the things involved in the reaction (baking soda/vinegar or shell/acid) limit how much of the other participant gets "used up".

Let me know if I can help with anything else - clarify things, answer other questions, etc. And, good luck!

Thanks so much for your input! It makes so much sense now!

I'm just glad that we were able to see some measurable change in mass at all. But yes, we did see a slight change in mass with the moon snails, but less than what we saw in the oysters. (we're talking 3% compared to 1%), but since the controls were so steady, we are taking the result as significant.

Hi 5thgrade_boy,

It sounds like your interpretation is sound given the data you collected. I'm happy to hear that things are coming together for you!