SplendidAcylation
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Strange behaviour in the reaction of calcium carbonate with propionic acid
Hi,
I have recently been experimenting with the pyrolysis of a mixture of calcium salts, for the preparation of ketones.
I required calcium propionate, so I decided to make it!
Having propionic acid and calcium carbonate (chalk), the reaction of these compounds seemed like the sensible approach. I have carried out this simple
reaction twice, obtaining the desired calcium propionate.
However, in the course of these two reactions, I have observed some phenomena which are somewhat odd to me, but which I suspect some of the geniuses
here might be able to explain.
My procedure was as follows:
1) A stoichiometric* quantity of calcium carbonate was placed in a flask
2) Two molar equivalents of propionic acid were weighed out in a beaker, and diluted with enough water to ensure the calcium propionate product
remained dissolved in the solution
3) The acid solution was then added to the calcium carbonate, with swirling. This resulted in some gentle fizzing, and lots of foamy bubbles.
4) Once all of the acid had been added slowly, the mixture was refluxed on a hot-water bath, until no more bubbles were visible, and audible fizzing
had ceased (around 1 hour).
5) The reaction mixture was then left overnight (whereupon some solids, assumed to be unreacted calcium carbonate, remained at the bottom of the
flask).
6) Having stood overnight, the reaction seemed to be complete, with no further fizzing, however the mixture still smelled strongly of propionic acid.
7) The reaction mixture was then filtered, and the filtrate, smelling of propionic acid, was evaporated down, yielding calcium propionate.
*In my first attempt, I used the stoichiometric quantity of calcium carbonate, and obtained a ~82% yield of calcium propionate in the end.
In the second attempt, I used a 1.5x excess of calcium carbonate, whereupon I obtained a ~86% yield of calcium propionate.
So, the mysteries are as follows:
At step 6, why does the reaction mixture smell strongly of propionic acid, even with a large excess of base being used?
Why is the yield sub-quantitative? (Obviously the answer to both of these questions seems to be: Because the reaction was not complete!)
So we can therefore replace these two questions with one question:
Why does the reaction progress fairly quickly to near-completion, and then appear to stall, or greatly decrease in rate?
I did some equilibrium calculations (shown below), which lead me to the conclusion that the problem did not lie in the realm of thermodynamics, as I
obtained K=4.36E9 for the equilibrium:
2C2H5COOH + CO3(2-) <--> 2C2H5COO- + H2O + CO2
This indicates the practically no propionic acid should be present at equilibrium, right?
Calculations:
CO3(2-) + H2O <--> HCO3- + OH- (K=10^-3.67)
HCO3- + H2O <--> H2CO3 + OH- (K=10^-7.7)
H3O+ + OH- <--> 2H2O (K=10^14)
C2H5COOH + H2O <--> C2H5COO- + H3O+ (K=10^-4.88)
H2CO3 <--> H2O + CO2 (K=588)
Firstly, adding the carbonate and bicarbonate equilibria, we obtain:
CO3(2-) + 2H2O <--> H2CO3 + 2OH- (K=10^-11.37)
Then, balancing the acid-dissociation reaction for propionic acid:
2C2H5COOH + 2H2O <--> 2C2H5COO- + 2H3O+ (K=((10^-4.88)^2) = K=10^-9.76)
Adding this to the previous equilibrium, we obtain:
2C2H5COOH + CO3(2-) + 4H2O <--> 2C2H5COO- + H2CO3 + 2OH- + 2H3O+ (K=10^-21.13)
Balancing the self-ionization reaction for water:
2H3O+ + 2OH- <--> 4H2O (K=10^28)
Adding these together gives:
2C2H5COOH + CO3(2-) <--> 2C2H5COO- + H2CO3 (K=10^6.87)
Finally, adding the hydration of CO2:
2C2H5COOH + CO3(2-) <--> 2C2H5COO- + H2O + CO2 (K=4.36E9)
I'm sure I probably made some errors in these calculations, but common sense also tells us that this reaction should proceed almost completely in the
forward direction!
The problem, therefore, must lie in the realm of kinetics?
This doesn't make much sense to me, as the reaction seems to reach ~80% completion within around one hour, and then suddenly seems to stall, which
doesn't seem consistent with the usual progression of a simple reaction like this.
Note: I also tried heating the reaction mixture further after standing overnight, but this did not seem to get the reaction started again.
Am I missing something obvious, or is this a bit weird?
Thanks!
[Edited on 3-8-2023 by SplendidAcylation]
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Fery
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Maybe this could be helpful:
Propionic acid is a weak acid, there is the reaction:
Ca(CH3CH2COO)2 <-> Ca2+ + 2 (CH3CH2COO-)
CH3CH2COO- + H2O <-> CH3CH2COOH + OH-
The solution of Ca propionate is slightly alkaline and there is always some very low concentration of the free acid in its solution in water. The free
acid is present even if you prepare a solution from very pure Ca propionate and boiled distilled water free of CO2.
If you used CaCO3 there is interfering Ca(HCO3)2 and if you let the mixture to stay overnight in the air, there is dissolved CO2. For the effect of
CO2, e.g. read this:
https://pubs.acs.org/doi/pdf/10.1021/ed079p29.1
propionic acid pKa = 4,86
H2CO3 pKa1 = 6,35
CO2 from air frees some more propionic acid from Ca propionate.
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teodor
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The similar result you will get also with strong acid and alkali earth carbonate. Reaction of diluted H2SO4 with excess of BaCO3 will go several
weeks. Ba(OH)2 reacts with moderate speed - slow withow stirring. Based on this observation I would say: the solubility of carbonate is a speed
limiting factor.
Why it works fast with the excess of concentrated acid? My guess is: in concentrated acid those carbonates are much more soluble.
[Edited on 4-8-2023 by teodor]
[Edited on 4-8-2023 by teodor]
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SplendidAcylation
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Thanks for the replies!
@Fery, that publication is quite interesting, it is hard to believe the pH could vary by such a large degree, simply by the action of a tiny amount of
CO2!
I must admit, I'm a bit confused, because CO2 is a much weaker acid than propionic acid, and it is also less soluble, so surely CO2 would liberate
practically zero propionic acid from the solution?
However I see your point, the CO2 is being continuously supplied from the air, while there is no propionic acid in the air, so the reaction would
eventually go in the direction you say. Hmmmm.
@teodor,
Wow, I didn't expect that!
So you suppose maybe there is a sort of logarithmic decrease in reaction rate, as the concentration of carbonate ion decreases not only because of the
progression of the reaction, but also due to the decreasing solubility of the calcium carbonate itself? Interesting.
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clearly_not_atara
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The equilibrium pH of water in normal air is like 5.7 or so IIRC. The pKa of propionic acid is around 4.9. So when half of the propionic acid is
deprotonated you have a pKa of 4.9. When you have 80% consumption of propionic acid the pKa is probably approaching air equilibrium wrt evaporation of
CO2. The equilibrium will shift a little due to the "reservoir" of carbonate but this should give you some intuition about what is happening.
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Boffis
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When I prepared calcium propanoate and calcium butanoate recently I used calcium hydroxide rather than the carbonate for the simple reason that it
does generate foaming. I disperse the slaked lime into a large volume of water and then while stirring vigorously slowly run in the appropriate acid.
Calcium propanoate is sparingly soluble so in the presence of a high concentration of acid and propanoate salt that occurs locally around the calcium
carbonate grains causes a protective layer to form around the residual carbonate preventing or slowing further reaction. For some acids such as
butanoic and phenylacetic acid I found it best to convert the acid to the sodium salt and then mix with calcium chloride solution. These reactions are
bet carried out hot and then allowed to cool, the product is more easily filterable.
In the end I purchased a bag of calcium propanoate off Ebay because it is cheap. In your case, SplendidAcylation, I suggest you add water sufficient
to reduce the final concentration of calcium propanoate to less than 100g per litre and heat almost to boil for a hour or so or until foaming ceases
(have a squeezy bottle of isopropanol to hand to kill the foam if it gets out of hand). Cool and filter off the product. You can evaporate down the
liquor but there isn't much salt left in solution if have cooled it in the fridge to 4C overnight before filtration.
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unionised
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If WIKI is to be believed, calcium propionate is quite soluble (about 50% w/v).
(My elderly copy of Merck says it's soluble toom but doesn't say how soluble)
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Boffis
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Hi unionised, you are right. I have just looked back at my notes from the work I did recently with the pyrolysis of this material.
I prepared it by from 100g a calcium hydroxide dispersed into 600ml of cold water and slowly added 200g of pure 98-99% propanoic acid. The mixture
became warm and the calcium hydroxide rapidly dissolved. The pH of the resulting solution was rather high so I added a little more acid and treated
the slightly cloudy solution with 5g of charcoal and vacuum filtered it. The solution stud overnight but no crystals formed so I evaporated it down in
a pan to start with but this requires constant stirring once solids start to separate as it bumps so I transferred it to a large shallow ceramic bowl
and heated on a steam bath until almost dry and then finished it of in an oven at about 100C. The yield was 230g of almost white solid that was almost
odourless and easily powdered.
It was the calcium butanoate that was more problematic, it is significantly less soluble (I estimate about 17-18g per 100ml of water) and has almost
no change of solubility with temperature.
So calcium propanoate is pretty soluble (>230gm in 600+ml of water) but the solubility doesn't change much was temperature so it doesn't
crystallise out on cooling saturated solution. This is a common feature of propanoates and acetates too.
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