[blakeb]

Hello, I'm currently conducting a personal research experiment involving calcium carbonate and its reaction with several acids. As of right now, I am setting calcium carbonate, in the form of seashells, into an excess of a 5% solution of vinegar. The reaction to my knowledge is as follows:

2CH₃COOH_(aq) + CaCO_3(s) H₂O_(l) + CO_2(g) + Ca(CH₃COO)_2(aq)

This reaction also has a carbonic acid intermediate that I have ignored since the decomposition is essentially spontaneous in the presence of water.

My dilemma is that the results seem to produce far more gas than anticipated. My scale may be a bit off, but now in several different trials involving 1.0 gram of seashell, which is >98% CaCO₃, much more gas was produced than I would have predicted stoichiometrically.

Using the ideal gas law and the volume of dry gas collected, the moles of gas collected are always at least double the moles of CaCO₃ that should have been dissolved. In trial 1, using 0.010 moles of calcium carbonate, I ended up producing 0.0252 moles of gas. In trial 2, I produced 0.0190 moles of gas. Currently, I am in the process of conducting a third trial, but far more gas has already been produced than would be expected.

If you want to double check my calculations, for trial 1 the temperature of the water was roughly 295.1K, and in trial 2, 292.7K. Atmospheric pressure, in torr, for trial 1 was 760.5, minus 19.95 torr for water pressure gives us a dry pressure of 740.6 torr. In trial 2, atmospheric pressure was 762 torr minus 17.2 torr for 745.6 torr for the dry gas. In trial 1, I collected 626mL of gas. In trial 2, I collected 466mL of gas. These numbers themselves are alarming of the accuracy of my experiment.

The confusion arises since, according to the chemical equation, calcium carbonate and carbon dioxide should have a 1:1 mole ratio. Does anybody have any idea where this extra gas is coming from? Or any ideas of possible common experimental flaws I could be making with my water displacement apparatus? Or perhaps calculation errors or oversights in certain elements of the gas law? Any and all help is appreciated.

[Borek]

Not seeing details of your experimental setup nor procedure it is hard to comment on possible error sources. Logic behind your calculations looks OK though. Definitely 1 g of CaCO₃ (~0.01 mole) should produce around 240 mL of CO₂ at STP (you are not at STP, but deviations are rather small and should not change the volume by more than few percent at most), so your result of over 600 mL for the first trial is seriously off.

[MNIO]

looking over your lab. 
  (1) The balanced equation is correct
  (2) The assumption that the seashell is 100% CaCO3 is probably close 
  (2) Your sat. vapor pressure of H2O is correct
  (3) The theoretical yield calculation is correct


In the second line, I've reduced the theoretical yield equation to values and operations.  I've colored the values you measured in RED.  If your actual yield is 626mL for a theoretical yield of 246m, then either...
  (1) your mass measured is really 2.53g not 1.0g
  (2) your actual temperature is  83°C = 183°F (at that temp, PsatH2O = 401torr)
  (3) your actual atmospheric pressure = 310 torr
  (4) the actual volume measurement is incorrect.


I suspect you would notice if you were working in a 183°F lab or a lab at 310 torr.  So I'm guessing we can rule those out for now.  So either (1) or (4) is the most likely root cause.


how did you measure mass of seashell?
how are you measuring actual volume?
what does your apparatus look like? 

[AWK]

The solubility of CO₂ in the water at RT is approximately 1.5 g/L. The solubility of CO₂ in pure acetic acid is about 10 times higher. The assumption that CO₂ in 5% acetic acid does not dissolve at all is nonsense, and so is collecting CO₂ over water. The only explanation for these suspicious results is a gross mistake in the sample weight.

Admittedly, blakeb writes "the volume of dry gas collected" do not give any evidence of dehumidification of the gas, but then there is no need to take into account the water vapor pressure.

[Borek]

blakeb writes "the volume of dry gas collected" do not give any evidence of dehumidification of the gas

I assume what he meant was that he subtracted saturated vapor pressure, this is equivalent in terms of the final result.

[MNIO]

AWK.. this is a standard problem in general chemistry. 
  (1) run a rxn and collect gaseous product "A" by displacement of water in
       an inverted water filled graduated cylinder.  The change in water level before
       and after the reaction = volume collected.
  (2) assume the gas collected is
        (a) at the same pressure as atmospheric pressure
        (b) at the same temperature as ambient temperature
        (c) composed of gas A saturated with water vapor (because we bubbled
             gas A up through water to collect it and it's in a sealed inverted cylinder).
  (3) we determine the pressure of just gas A by subtracting the
        saturated vapor pressure of water at that ambient temperature
        we call this value the "pressure of the dry gas A"
        it's not really dried, but it is the pressure IF the gas were dried
  (4) use use that P dry gas A + Volume + Temp to generate moles of gas A
       then calculate % yield or whatever.

This type of problem is covered in every gen chem textbook out there, every student sees this type of problem on exams, and every student runs this type of experiment in a lab at some point.  We've all seen it.  The procedure and calculations BlakeB used are spot on.

To your point.. IF the gas dissolved in the water (which some does) the actual mass recovered < theoretical mass ---> % yield < 100%.  BlakeB is not observing that.  He's observing 250% yield. 

*********
The reason I walked through the calcs the way I did was to get BlakeB thinking about possible sources of error.  1 to check is mass.  Another is water level measurement.  How was it measured?  were there LEAKS that allowed the water level to drop in the graduated cylinder even without adding any CO2?  Did he use a graduated cylinder or something else?.  The typical setup looks like this.



what was his setup?

[AWK]

This standard setup is good for measuring the volume of nitrogen emitted during e.g. the decomposition of nitrite, but not for measuring the volume of carbon dioxide.
The dissolution of aragonite of seashells in dilute acid will be slow enough for air to mix with carbon dioxide by diffusion and therefore it can be estimated that even about half of the carbon dioxide can be dissolved in water + in the reaction mixture.
By simplifying the problem to STP, we will make a smaller error than that resulting from such a setup (but this error is difficult to estimate more precisely). Taking pressure and temperature as well as CaCO₃ content into account, the theoretical volume with respect to STP can be expected to be around -2% (carbonate content) + 3% (H₂O vapors) + 7% (temperature). Any larger volume is due to gross error in the experiment.
Calculating what the pressure or temperature would have to be to obtain such volumes only obscures the problem.
blakeb rightly noticed that he got unrealistic results, but most likely due to the size of the samples. For such a setup you should expect much smaller volumes than theoretical.

[MNIO]

AWK.. (1) This setup is used for a variety of experiments.  Would you like to see examples from general chem textbooks?   (2) Low % yield can be explained by dissolution of gases in the water (among other stuff). (3) BlakeB's theoretical volume was 246mL.. his actual volume was 626mL.  That's a 250% yield.  If dissolved gases were a factor, his volume would have been < 246mL.

[blakeb]

Hello everyone. I'm pleasantly surprised at everyone who took the time to think through my experimental issue, and I do have an update. I believe the problem was a rather gross oversight in my apparatus, but first, a quick description of my apparatus. I used an inverted 500mL erlenmeyer flask for the collection of the gas. I know that this isn't the most precise instrument, but for the theoretical gas yield its the most convenient. To measure volume, I mark the bottom of the meniscus whenever the reaction has stopped, i.e., there isn't a seashell in the reaction flask, and shaking releases no further gas. I then fill the erlenmeyer up with water to the mark I previously made, and measure using a graduated cylinder in intervals of 100 milliliters.

I'm pretty sure that the misstep in my apparatus was having the tube near the top of the inverted erlenmeyer. With this setup, once some gas has been produced, and the end of the tube is no longer submerged under water, the pocket will suck in more air to displace the water due to gravity, in the event a leak. I spent around 3-4 hours today simply toying with my apparatus, running many simple baking soda and vinegar reactions, since they are both abundant, release CO₂ gas, and the reaction happens much faster. If there were to be a leak, the gas in the erlenmeyer would be significantly made up of air that was sucked into the pocket to alleviate the potential energy of the water.

To solve this, I connected a two-holed rubber stopper to the bottom of the inverted erlenmeyer, with one of the holes connected to the tube that leads to the reaction vessel. With a small glass tube entering the inverted erlenmeyer, The tube is always going to remain under water, meaning that if there were a leak, say in the tube or at the rubber stopper of the reaction flask, then the reaction gas would be lost, rather than air being sucked into my inverted erlenmeyer. Doing this, I found my yield to be lower than expected on some trials, which is much better than a 250% yield, and easier to explain. An image of my apparatus has been attached below. It's just a simple apparatus set up in my bathroom.

It is still rather confusing though, since when I did the experiment, it would take at least an hour for bubbles to even form in the inverted erlenmeyer. My guess was that pressure needed to build in order to counteract the pressure of the water in the tube, although the atmospheric pressure in the reaction flask should be the same as the water. Upon shaking the reaction vessel, the gas in the tube would move along slightly, until eventually, it released into the inverted erlenmeyer, which would chain the releasing of what I assumed to be the pressurized CO₂ formed from the reaction. If anything here sounds incorrect, I'd like to hear.

Aside from that, thank you to everyone for giving my some insight, especially with the CO₂ being dissolved in the vinegar and water explanation. This was something I completely overlooked, and would help very much in explaining the decreased yield of gas.

[Borek]

Just leave it without reagents for several hours and see what happens. If it sucks, it sucks

[AWK]

As Borek pointed out, a large error may cause leakage of the apparatus. Due to the underpressure caused by the water column, air can be sucked into the system. A similar error can cause the reaction mixture to heat up.
This kit is not suitable for water-soluble gas measurements - from the information provided, a CO₂ solubility of approximately 0.5 g/L (partial pressure of CO2 ~ 1/3 atm) can be expected.

[AWK]

AWK.. (1) This setup is used for a variety of experiments. 

I know.

Would you like to see examples from general chem textbooks? 

I've seen a lot. Quite a lot with errors.

(2) Low % yield can be explained by dissolution of gases in the water (among other stuff).

This is how it should be in this experiment.

(3) BlakeB's theoretical volume was 246mL.. his actual volume was 626mL. 

My estimate without using a calculator: 1.08 x 224 = 242 - quite reasonable - less than 2 % error.

That's a 250% yield.
_{No, that's almost a 300% error.}
 If dissolved gases were a factor, his volume would have been < 246mL.
_{Agree - it should be so.}

[Corribus]

As AWK has noted, I do not think this is a good experimental setup for measuring gas produced by a reaction when the reaction rate as slow, because your gas generation competes with partitioning of the (pressurized) gas into the water (and diffusion out of the system completely). This has been noted in an unrelated thread recently, albeit in a different context: https://www.chemicalforums.com/index.php?topic=104544.0

Gas diffusion would a source of error for any gas produced, not just carbon dioxide, although carbon dioxide is particularly egregious because it actually reacts with water, not just diffuses through it. For that matter, it would happen regardless of what the gas generation rate is - just that if the gas generation rate is fast, you have a good chance of measuring the volume accurately before it all diffuses away. Don't wait three days before you measure though.

Ehlrenmeyer isn't great, accuracy and precision is bad for one thing, and the more surface area you have at the gas/liquid interface, the faster your partitioning of gas into the liquid will be. I guess, better than a beaker though.

In any case, I guess the effect would cause an under-estimate of the gas produced, which is not what was observed here. The leak identified by the OP describes the immediate problem better. I guess the positive error would be even greater except the gas diffusion counteracts it to some degree.

By the way, with due respect to MNIO, I would also point out that just because something is described in a general chemistry textbook does not mean it is the best (or even a good) way to do things. Experiments in general chemistry lab are designed to show basic concepts with idealized reaction systems. When trying to adapt these types of experiments to more complex reactions, this is what happens. I don't think anyone would do this kind of experiment in a real research laboratory because of some of the issues pointed out. There are much better, more precise ways to measure gas concentrations. This isn't meant to discourage the OP - even with all the problems, you are learning and you can acquire useful information. Just something to keep in mind.

[blakeb]

Thank you for your concern, Corribus. Now that I have sorted the supposed problem of leakage and the placement of the end of the tube in the inverted collection erlenmeyer, I am facing this issue of carbon dioxide diffusion within the water. My question would now be what alternatives are there for gas collection of this sort? I am aware of an upward displacement apparatus, but how would I be able to measure the gas collected? A couple ideas I have would involve either limewater or a pressure chamber. Thoughts?

Also, that linked thread you sent is loosely related to my research. In general, my research is revolving around global warming and ocean acidification, so thank you for that. This experiment in particular is designed to gauge the potential reaction between acid rain and chalk deposits, and how that could affect calcifying species in the ocean, and ultimately the oceans ability to buffer atmospheric acidity and temperature.

[Borek]

First things first, I would go for a stronger acid. That should make the reaction faster so less time for losses, plus lower pH will shift the CO₂/H₂O reaction equilibrium to the left - all in all less possible problems to deal with.