Sciencemadness Discussion Board

calcium carbonate composition in seashells

CHRIS25 - 18-11-2014 at 12:27

I have just finished a lengthy experiment with seashells and am surprised by the results.

Experiment:
Exactly 100g white seashells (see image) were dissolved in an eventual total 7 moles HCl, the resulting CaCl2 solution was then treated with NaHCO3.

The expected outcome was approx 90g or less of Calcium Carbonate based on various topics that referred to seashells as being mostly Calcium carbonate minus 2 - 5% in protein and other ingredients.

These are now my results:

375g CaCO3 fully dried and weighed in as such.

Working backwards then:

375g CaCO3 = 3.75mol
CaCl2 solution must have been 2.8mol
and the actual needed HCl stoichiometrically would have been 5.6mol, (Which is quite close to the 7mol added good enough knowing that I added a little to much to allow for loss of HCl gas)

This then means that I got 3 times as much CaCO3 than what was should have been. The only mistake I could have made was in the original weight of the dry seashells, but this is really highly unlikely - honestly, not by 250g?




IMG_1750.jpg - 83kB

[Edited on 18-11-2014 by CHRIS25]

gdflp - 18-11-2014 at 12:36

Is it possible that the shells you began with were mostly calcium oxide, and not calcium carbonate? If so, then you would get around 1.8mols of calcium in solution. In addition, sodium bicarbonate will not precipitate calcium carbonate, it will result in calcium bicarbonate which is soluble in water. The solution must then be evaporated, which causes the calcium bicarbonate to decompose, and any soluble salts will redissolve. If calcium oxide was the starting material, this would result in 180g of CaCO3 which is closer to what you describe. Is it possible that there is still sodium left in your final product? I would try a flame test on the suspect compound.

unionised - 18-11-2014 at 12:37

Interesting.
How carefully did you wash the precipitated CaCO3?

The first thing you need to do is repeat this experiment.
Incidentally, the chances of sea shells being made from something which reacts vigorously with water is, shall we say, slim.

[Edited on 18-11-14 by unionised]

CHRIS25 - 18-11-2014 at 13:24

Hi, Firstly:
CaCl2+2NaHCO3=CaCO3+CO2+2NaCl

(by the way, I tasted the filtrate and it was deliciously salty, to confirm).

No Seashells are not made of Calcium Oxide, but the carbonate. So in effect I am reacting calcium carbonate with the HCl. Actually no, I did not wash the precipitate at all actually, since my wife uses it to make chalk paint, (very expensive product in the shop), it actually works just as good as the bought product.

@Unionised: bit confused on why you said this: ".....the chances of sea shells being made from something which reacts vigorously with water is, shall we say, slim....."

Yes I will repeat with a very small amount, but I did this about 6 months ago and actually got the same results, though I never posted anything since I was, again, making something quick for my wife's hobby and I was not concerned with analysing everything.

[Edited on 18-11-2014 by CHRIS25]

Little_Ghost_again - 18-11-2014 at 13:51

those are limpits by the look of it most the limpets I have used contain alot of magnesium salts, limpits are exposed to the sun and air for many hours a day (they are tidal) personally I dont know the what but I do know the why. because of this exposure limpets excrete slightly different salts for the shell as it needs to be harder than many, both because on environmental factors (tides,air,sun) and to prevent excessive predation.
I havnt got around to testing what the shell is made of but its definitely different to non tidal molluscs I will ask in the in house biologist if you like.
I noticed this difference a while back when collecting shells for use, the very best for calcium carbonate are oysters or if you lucky like I am and live near a beach where they wash up cuttle fish bones are very soft and mostly calcium carbonate. Other shells contain a mix of calcium,magnesium and sodium salts. Give me a couple of weeks and I will post some results from different shells using a GC.

Little_Ghost_again - 18-11-2014 at 13:52

I forgot to add, there are some good papers on shells and there variation in chemistry, join something like academic gateway

DistractionGrating - 18-11-2014 at 13:53

Seashells are most certainly made of calcium carbonate, as any knowledgeable marine aquarist can tell you. Both the amount of resulting CaCO3 and the 7 moles of HCl required point to an error in the original weight, don't you think?

Little_Ghost_again - 18-11-2014 at 14:19

Quote: Originally posted by DistractionGrating  
Seashells are most certainly made of calcium carbonate, as any knowledgeable marine aquarist can tell you. Both the amount of resulting CaCO3 and the 7 moles of HCl required point to an error in the original weight, don't you think?

I am not disputing there is alot of calcium carbonate there, what I am saying is there is a definite chemical difference or properties difference with the shells of some species, rather than look a complete arse I will go search etc and come back with some references and results.
Its purely because I have used shells for other things that I noticed a difference myself with different shells.
So until I go get papers to cite etc its going to be pointless me saying much more. while i agree the maths is probably wrong, I am also aware there are chemical differences with the shells he chose.
There is also something to do with the different forms like argonite (sp) and calcite admittedly both calcium carbonate.


[Edited on 18-11-2014 by Little_Ghost_again]

aga - 18-11-2014 at 14:31

There's a parallel experiment waiting for me to go and weigh in my shed.

Staring with ~100g shells in one pot and ~100g quarry stone (chalk) in the other.

So far one's a bit brown tinged, and the other smelt sulphur-y when reacting.

DistractionGrating - 18-11-2014 at 14:38

Aragonite and Calcite are both CaCO3, but with different crystal lattices. I'm not entirely sure about seashells, but coral skeletons will be largely CaCO3, but will also have some MgCO3 and a tiny amount of SrCO3 as well.

Little_Ghost_again - 18-11-2014 at 14:50

Quote: Originally posted by aga  
There's a parallel experiment waiting for me to go and weigh in my shed.

Staring with ~100g shells in one pot and ~100g quarry stone (chalk) in the other.

So far one's a bit brown tinged, and the other smelt sulphur-y when reacting.


Out of interest are the shells all the same?

This could be a fun and interesting thing to look into, I live near several beaches and each is different, it would be interesting to take several different shells and test them for there composition.
I also have a marine aquarium and my corals are pretty heavy users of calcium and magnesium (I have a calcium reactor and add trace minerals by pump daily). Actually I am putting my reef tank up for sale, I have had it 4 years now and its superb! I spent a fortune on it and also designed and built many of the control systems for it.
I doubt I will get a fraction of its value but I need the money for other stuff like my soap business :D.
Its also costing a fortune to run because the electric here at home is very expensive. I will keep my freshwater tanks (for now anyway) and cant see me ever giving up the dutch tank :D

gdflp - 18-11-2014 at 14:57

Quote: Originally posted by CHRIS25  
Hi, Firstly:
CaCl2+2NaHCO3=CaCO3+CO2+2NaCl

(by the way, I tasted the filtrate and it was deliciously salty, to confirm).

No Seashells are not made of Calcium Oxide, but the carbonate. So in effect I am reacting calcium carbonate with the HCl. Actually no, I did not wash the precipitate at all actually, since my wife uses it to make chalk paint, (very expensive product in the shop), it actually works just as good as the bought product.

@Unionised: bit confused on why you said this: ".....the chances of sea shells being made from something which reacts vigorously with water is, shall we say, slim....."

Yes I will repeat with a very small amount, but I did this about 6 months ago and actually got the same results, though I never posted anything since I was, again, making something quick for my wife's hobby and I was not concerned with analysing everything.

[Edited on 18-11-2014 by CHRIS25]


No, actually sodium bicarbonate reacts with calcium in a two-step reaction. Calcium bicarbonate and sodium chloride are the initial two products. Any precipitation which occurs is calcium carbonate since calcium bicarbonate isn't stable in solid form. But this reaction 2CaHCO3 --> CaCO3 + H2O + CO2 doesn't happen immediately, it is a very slow reaction since CaHCO3 is stable in solution.

CHRIS25 - 18-11-2014 at 15:19

@Distraction Grating. Actually the maths seems fine, the yield was 3.75moles Carbonate. The HCl used was 7moles. The initial balanced equation calls for 2:1 HCl to CaCO3 respectively. Due to the fact that this was not an empirically designed process, rather something for my wife to put in her paint, there may well have been excess HCl added without realising, and also I did not wash the precipitate. I am going to repeat the experiment with determined accuracy with a couple of shells. But in any case, the original 100g of seashells is in no way mistaken, since I weighed out 300 grams and could only find time to dissolve 100.

@Little Ghost something. Yes, this would be interesting actually, I appreciate the extra info. By the way, I did not know about the oysters, I can get an un-ending supply of these as well close by. I will have a look at Gateway.

@Aga. Looking forward to your results.

@Gdifip. Ok I did not know that, but that is probably why it is not seen in any of the balanced equations, it really does not seem to me to be important in attributing any help to solving the initial question. While I respect the chemistry here and am glad to be told, I don't see how it helps. Although yes, it is definitely a slow process, it took hours for every 200 mLs of Calcium chloride to get the precipitate, foam foam foam, tons of it, CO2 and Water naturally. Constant stirring with stir bar, but I emphasize absolutely no heat necessary.

[Edited on 18-11-2014 by CHRIS25]

unionised - 19-11-2014 at 11:19

Quote: Originally posted by CHRIS25  

@Unionised: bit confused on why you said this: ".....the chances of sea shells being made from something which reacts vigorously with water is, shall we say, slim....."


[Edited on 18-11-2014 by CHRIS25]


Because Gdflp wrote "Is it possible that the shells you began with were mostly calcium oxide, and not calcium carbonate? " and CaO reacts vigorously with water.

Etaoin Shrdlu - 19-11-2014 at 13:18

Oooooookay. Let's assume the mass of calcium alone in your shells was 100g, for the fun of it. That only nets you ~250g calcium carbonate with perfect conversion.

You measured wrong, or you cleaned it wrong. How much NaHCO3 did you add?

[Edited on 11-19-2014 by Etaoin Shrdlu]

blogfast25 - 19-11-2014 at 14:26

There has to be a weighing error somewhere, the discrepancy between theor. and actual yield is just ridiculously too high.

When you precipitate the CaCO3 make sure to do it from a fairly dilute solution of CaCl2, e.g. max 1 M. Na2CO3 is better here than NaHCO3 because Ca(HCO3)2 is soluble in water
(calcium in bottled or tap water for instance. Ca bicarbonate is also the cause of scaling in kettles because on heating it decomposes, so over time scale gets built up).

Allow the freshly formed CaCO3 to digest (stand) overnight before filtering. Wash with plenty hot water, finally with a bit of deionised water.

Dry thoroughly at at least 200 C for 2 hours.

aga - 19-11-2014 at 14:36

Quote: Originally posted by unionised  
Because Gdflp wrote "Is it possible that the shells you began with were mostly calcium oxide, and not calcium carbonate? " and CaO reacts vigorously with water.

I think CHRIS25 is pointing out that a CaO-shelled Underwater-dwelling mollusc would be homeless pretty soon.

New Test Completed

CHRIS25 - 20-11-2014 at 01:32

One sad seashell, crushed and weighed: 10.1g Assumed a 0.101m for this.
HCl 36% 0.202m made up in 8M sstrength, total soln 25mLs

Covered and warmed, dissloved and filtered.

Resulting solution treated with Sodium bicarbonate, filtered and washed twice, dried and weighed.

Stoichiometric expectations:

0.101mol CaCl2 (11.2g) + 0.202mol NaHCO3 (16.9g) = CaCO3 0.101mol (20.2g)

Actual NaHCO3 required to complete precipitation: 13.13g ( 0.16mol) OR 17.35g (0.2mol) Was difficult to tell. But seeing that I did not get 20g of carbonate I have to assume a figure of 13.13g completed the reaction.

Actual Yield CaCO3: 11.91g (0.12mol)


The Taste test:
The suspect batch is very salty, but this new test is completely free of any salty taste, tastes like perfect chalk.

Conclusion: I will wash 50g of the suspect batch, dry and weigh and taste. If there is a significant difference in weight then this is where my problem was.

@Etaoin Shridu. """You measured wrong, or you cleaned it wrong. How much NaHCO3 did you add?"""" Total was 452g / 5.4 moles, but there was consistency in the 200mL batches in that each batch took 73,75,84g sodium bicarbonate respectively. Yes I know, total seems overkill, but the final precipitations were difficult to spot since I did this in batches of 500mLs then 200mLs, also I recognise now that I should add more water to each batch. The CaCl2 solutions were far too concentrated.

@Blogfast. Hallo, I don't get why calcium bicarbonate would be a problem at all, just can't see it unless it prevents some of the Ca from the CaCl2 from being joined to the carbonate because it hogs the carbonate in a disolved state in solution? Using Sodium Carbonate does look a much more efficient idea; (I was going to do this, I have a ton of the stuff), just assumed that for some reason since everyone uses the bicarbonate that must be the correct way. I certainly appreciate the heads up on doing this process more efficiently with your explanation, I need to know how to improve the things that I do all the time, so thanks..

[Edited on 20-11-2014 by CHRIS25]

blogfast25 - 20-11-2014 at 04:49

Seems it was an extreme case of poor washing/occlusion. You'll need to BOIL the old batch with clean water for a bit, allow to stand overnight and filter and wash again.

"just assumed that for some reason since everyone uses the bicarbonate that must be the correct way"

No: there are reasons to use bicarbonate and there are reasons to use carbonate: it's case dependent. For CaCO3 I'd use carbonate each and every time.

Request: don't abbreviate 'mol' to 'm'. It's like abbreviating 'cm' to 'c' or 'kg' to 'k': 'c' what? 'k' what? Respect our language, por favor! :D To try and be as unambiguous as possible makes good science.

[Edited on 20-11-2014 by blogfast25]

CHRIS25 - 20-11-2014 at 07:44

Quote: Originally posted by blogfast25  


No: there are reasons to use bicarbonate and there are reasons to use carbonate: it's case dependent. For CaCO3 I'd use carbonate each and every time.

Request: don't abbreviate 'mol' to 'm'. It's like abbreviating 'cm' to 'c' or 'kg' to 'k': 'c' what? 'k' what? Respect our language, por favor! :D To try and be as unambiguous as possible makes good science.

[Edited on 20-11-2014 by blogfast25]


What would be the reasons then for carbonate and bicarbonate?

As for abbreviation, I actually thought that it was acceptable to use a small 'm' for Mole and large 'M' for molarity in that this was an accepted abbreviation. Was I wrong?

gdflp - 20-11-2014 at 08:47

'M' for molarity is an accepted abbreviation. However, 'mol' is the accepted abbreviation for mole, not 'm'. It's not an uncommon mistake, lots of people often come to the same conclusion you did. In all honesty, it's not that big of a deal, everyone will most likely understand what you mean, it just looks more professional the proper way.

blogfast25 - 20-11-2014 at 10:37

Quote: Originally posted by CHRIS25  
[
What would be the reasons then for carbonate and bicarbonate?



Assuming it's a carbonate you want to precipitate, which (carbonate or bicarbonate) to use depends on solubility of the corresponding hydroxide and also on the bicarbonate if it exists. In the case of calcium the hydroxide is fairly soluble, the carbonate very insoluble and the bicarbonate fairly soluble: using Na2CO3/K2CO3 is then recommended but bicarbonate also works.

If the hydroxide is very insoluble (e.g. Cu<sup>2+</sup>;) the bicarbonate tends to precipitate the hydroxide or a hydroxycarbonate ('basic carbonate'), not the actual carbonate.

In some cases no carbonate can be precipitated: see Al<sup>3+</sup> and Fe<sup>3+</sup>, in those cases carbonates, bicarbonates, hydroxide or ammonia all precipitate Al(OH)3 or Fe(OH)3.

These subtleties are caused by the fact that carbonate or bicarbonate solutions are alkaline, so they also contain OH<sup>-</sup> ions, which kind of 'compete' (with the carbonate/bicarbonate ions) as to what exactly precipitates.


[Edited on 20-11-2014 by blogfast25]

DraconicAcid - 20-11-2014 at 10:56

How did you dry it?

I'd guess that it was still damp or hydrated with a yield like that.

CHRIS25 - 20-11-2014 at 12:39

@gdifp. Ok, thanks. Good to know that then. mol it is.
@Blogfast. Thankyou, into my notebook then.
@DraconicAcid. Believe it or not it dries by the open fire for hours and hours, then the room stays heated all night due to neighbour's boiler against the wall plus all the next day and so on, I usually allow two to three days depending on amounts and temp of room and taste buds! (Yes I have a habit of eating my chemicals just to make sure I made the right one):D

blogfast25 - 20-11-2014 at 13:02

Quote: Originally posted by CHRIS25  
@DraconicAcid. Believe it or not it dries by the open fire for hours and hours, then the room stays heated all night due to neighbour's boiler against the wall plus all the next day and so on, I usually allow two to three days depending on amounts and temp of room and taste buds! (Yes I have a habit of eating my chemicals just to make sure I made the right one):D


That might not even be enough, you know. Most of us dry using directly applied heat, such as from an oven. For homemade CaCO3 a domestic, kitchen oven will not be a problem for instance, as it's totally non toxic. An old ceramic kitchenware oven dish can serve as support.

CHRIS25 - 20-11-2014 at 13:32

Ah, if the wife allows...but yes, I should then try this with what I have and weigh it before and after. maybe that this is necessary then.

Washing and drying anew - New results

CHRIS25 - 21-11-2014 at 14:44

Took 80g of carbonate not washed, just filtered once, dried under old conditions and then weighed: 80g
This was then washed thoroughly and dried under old conditions (house fire for 3 days): 70.4g
This was then dried in oven 180c for 1 hour: 54.6g


Took the rest of the "dried over three days by fire", then put that into
the oven for 1 1/2 hours 180c
Before oven: 240g after oven: 188.6g

So what people were saying about drying and not washing were correct. Not only was a lot of salt removed, but despite the thoroughly dry chalky feel of the carbonate I was surprised by the huge amount of water still left in there.

[Edited on 21-11-2014 by CHRIS25]

blogfast25 - 22-11-2014 at 04:22

Quote: Originally posted by CHRIS25  


So what people were saying about drying and not washing were correct. Not only was a lot of salt removed, but despite the thoroughly dry chalky feel of the carbonate I was surprised by the huge amount of water still left in there.



The last bit of water is 'hiding' in all the microscopic nooks and crannies of the CaCO3, at temperatures below the BP of water that water is very slow to evaporate. Heat over 100 C and drying becomes much, much faster.

aga - 24-11-2014 at 12:22

Parallel Experiments

A sample of mostly white clam-like sea shells and fragments were collected from the sea shore.
Rock samples from a Limestone quarry were also collected.

100g each of Rock and Shell were dissolved in HCl solution.
The HCl was added in lots of around 50ml. When reaction ceased (no bubbles) another lot of 50ml was added to each sample.
Eventually 300ml of the unknown HCl concentration were added to each sample over a period of 3 days.
The Rock sample smelt distinctly sulphurous during the dissolution.

Each sample was then filtered.
The Shells sample left a residue of small shell fragments and brown sludge, weighing 2.84g.
The Rock sample residue was 19.60g of mainly larger rock pieces and a fine grey powder, which made up 2.74g of that weight.

The CaCl2 in the Shell solution was calculated to be 107.7g, requiring 163.05g of NaHCO3 to convert back to the CaCO3.
The Rock solution calculated to have 89.2g of CaCl2, requiring 135.04g of bicarbonate
No sodium carbonate was handy, and there was just 250g of bicarbonate on hand.

Both solutions were tested for various ions, namely Fe2+, Fe3+, SO4 2- Al(any)+ and Cl-, the results being :-

Shells: None
Rocks: Fe2+, Fe3+, Al+, Cl-

First the bicarbonate was added to the Rock solution, and was added too quickly.
The stuff fizzed out of the pot, destroying it's further relevance in the experiment.

The Shells solution has the bicarbonate added much more slowly, and the frothing was contained in the pot. It was found that the bicarbonate could be added faster later on in the reaction.

Due to selfless deprivation, only 125g of bicarbonate was available for this reaction.
After the reaction ceased (left it overnight) the resulting calcium carbonate was dried for an hour at 200C in the oven, using an appropriate container (dog's stainless steel water bowl).

After an hour, the white slush had solidified with a white crust on top.
When the crust was broken, fine white powder was found underneath.

After removing from the bowl with a hammer, the resulting mass was 75.75g

With sufficient carbonate, i calculate that the Yield should have been (75.75 / 125) * 163.05 = 98.81g

Edit

Well, similar.
I forgot to factor out the CaCl2 remaining in the 75.75g of white stuff.

Edit
This was never edited, ever.

[Edited on 24-11-2014 by aga]

DraconicAcid - 24-11-2014 at 13:11

Quote: Originally posted by aga  
Both solutions were tested for various ions... So2-


What's that supposed to be?

aga - 24-11-2014 at 13:16

What ?
I see no So2- anywhere, and it says in the post that it has never ever been edited, ever.
Must be a software glitch.

DraconicAcid - 24-11-2014 at 13:18

Quote: Originally posted by aga  
What ?
I see no So2- anywhere, and it says in the post that it has never ever been edited, ever.
Must be a software glitch.


Must be.

blogfast25 - 24-11-2014 at 14:23

Quote: Originally posted by aga  
namely Fe2+, Fe3+, SO4 2- Al(any)+ and Cl-, the results being :-

Shells: None
Rocks: Fe2+, Fe3+, Al+, Cl-



"Al(any)+" can 100 % safely be said to be Al<sup>3+</sup>, as Al(I) only exists in very exotic circumstances. So how did you test for Al<sup>3+</sup>? I'm curious because I only know of one fairly difficult test (Aluminon).

Re. ferrous ions, it's safe to say that in such Very Olde Things only ferric iron can exist, due to the inevitable oxidation of the former to the latter over Eons of Time.

aga - 24-11-2014 at 14:40

It was a New rock. I only just found it ... Doh !

I forgot that Fe3+ will prove positive to BOTH potassium hexacyanoferrate *and* ammonium thiocyanate.
Thanks for the reminder.

I got an 'Aluminon' testing kit from www.oxfordchemserve.com, the bestest chem website ever.

The kit is reasonably tricky to use, as you have to keep the pH within a narrow range all the time.

I don't know of any other readily available test.

blogfast25 - 24-11-2014 at 14:49

It's possible that the HCl insoluble residue of both the limestone and shell could have contained some alumina but to find it you'd have to practically alkali fuse the residue to solubilise the alumina.

Since as Al has no prominent role in metabolisms it's unlikely to be found significantly in the Earthly remains of living things like shells and limestone.

aga - 24-11-2014 at 15:13

Was a limey kinda rock from an area which previously was highly volcanic, so most deposits are mostly igneous or metamorphic, with the odd entire Rocky Beach sticking out at an angle.

Perhaps i was slightly overstating the Limestone nature of the rock ...

Ok. Hands up. Fair Cop.

It was just some random rock i found in a disused quarry.

Texium (zts16) - 24-11-2014 at 15:17

Quote: Originally posted by blogfast25  

If the hydroxide is very insoluble (e.g. Cu<sup>2+</sup>;) the bicarbonate tends to precipitate the hydroxide or a hydroxycarbonate ('basic carbonate'), not the actual carbonate.
Huh, I was under the impression that for copper, using Na2CO3 is more likely to precipitate basic copper carbonate than NaHCO3 is, since Na2CO3 is more basic than NaHCO3, and therefore a solution of it contains more OH<sup>-</sup> ions.

Source: http://www.aqion.de/site/191
Lists the pH of a 1mM solution of NaHCO3 as 8.27 and that of a 1mM Na2CO3 solution as 10.52

blogfast25 - 25-11-2014 at 05:57

Quote: Originally posted by aga  
Was a limey kinda rock from an area which previously was highly volcanic, so most deposits are mostly igneous or metamorphic, [...]


There's no 'igneous limestone', limestone is always sedimentary. Metamorphic limestone is marble: limestone that's been melted under very high pressure, hence the pretty streaks and higher hardness.

What you used was probably plain limestone.

Quote: Originally posted by zts16  
[since Na2CO3 is more basic than NaHCO3, and therefore a solution of it contains more OH<sup>-</sup> ions.



Yes, but what else is much higher in concentration: carbonate ions. In bicarbonate solutions there are almost no carbonate ions.

[Edited on 25-11-2014 by blogfast25]

MrHomeScientist - 25-11-2014 at 06:29

Quote: Originally posted by CHRIS25  
(Yes I have a habit of eating my chemicals just to make sure I made the right one):D

This is a tremendously bad idea.