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CHRIS25
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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.
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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CHRIS25
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Washing and drying anew - New results
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]
‘Calcination… is such a Separation of Bodies by Fire, as makes ‘em easily reducible into Powder; and for that reason ‘tis call’d by some
Chymical Pulverization.’ (John Friend, Chymical Lectures London, 1712)
Right is right, even if everyone is against it, and wrong is wrong, even if everyone is for it. (William Penn 1644-1718)
The very nature of Random, Chance development precludes the existence of Order - strange that our organic and inorganic world is so well defined by
precision and law. (me)
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blogfast25
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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.
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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.
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aga
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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]
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DraconicAcid
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What's that supposed to be?
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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aga
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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.
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DraconicAcid
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| 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.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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| Quote: Originally posted by aga | namely Fe2+, Fe3+, SO4 2- Al(any)+ and Cl-, the results being :-
Shells: None
Rocks: Fe2+, Fe3+, Al+, Cl-
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"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.
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aga
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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.
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blogfast25
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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.
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aga
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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.
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Texium
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| 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.
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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
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blogfast25
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| 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, [...]
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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.
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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]
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MrHomeScientist
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This is a tremendously bad idea.
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