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blogfast25
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Also, when using oxalic acid, do you use the dihydrate or anhydrous as a standard? If the latter how do you dehydrate? One source mentions possible
brown or black spots if you get the dehydration at just below 100 C (IIRW) wrong...
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entropy51
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The constant boiling HCl method is in the literature. Essentially a somewhat diluted HCl solution is distilled until a constant boiling point is
reached and that fraction is collected. The BP depends on the barometric pressure and there are tables for finding the exact concentration of HCl in
the distillate. A weighed amount of the distillate is diluted to produce a standard solution. If Google doesn't turn it up, I can find a citation
for you. I'd be surprised if it hadn't been posted on the forum already.
I used the oxalic acid dihydrate. IIRC I dried the recrystallized oxalic acid in a vacuum dessicator because it can indeed decompose if heated too
much. But so can Na2CO3. If the Na2CO3 method works for you that is good, but I would cross check it against another primary standard just to be
sure. At one point I did a detailed comparison of the methods until I found agreement between at least two methods and that point I began to trust
the results.
Beware of Na2CO3 contamination of your NaOH. Apparently a good method, which I use, is to make a 50% solution of NaOH. Over time the Na2CO3 will
precipitate out of this and the clear supernatant can be diluted and standardized.
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blogfast25
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entropy51:
Yes, the HCl azeotrope method re-occurred to me last night. It's within my capability envelope.
Na2CO3 is a tricky one. Hydrolysis of the carbonate is the main danger. I'm revising a new procedure based on decomp. of bicar (rather than starting
from commercial washing soda) and new recryst. procedure of the thus formed Na2CO3. I'll keep you posted.
Carbonate in NaOH is a worry: precautions are taken but you can't do much against stuff that's already in the product. I will definitely try the 50 %
method, almost immediately actually.
[Edited on 12-8-2010 by blogfast25]
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kmno4
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Value ΔG = 6,49 kcal/mol strictly corresponds to thermodynamic constant K=1,754x10<sup>-5</sup> for acetic acid
( ΔG= -RTlnKa) and you can calculate it by yourself.
K contains activities and for diluted solutions a<sub>H2O</sub>=1.
Parts ΔH and ΔS are calculated on the basis of measurements of temperature coefficient of K .
Unfortunately this is possible only for weak acids, for which K is small and independent (almost) on concentration.
Author of my book gives reference to unavailable electronically article:
http://adsabs.harvard.edu/abs/1949ZNatA...4..171B
for more datails.
BTW.
Na2CO3 can be estimated in NaOH by double titration of the same sample: first with phenolphthalein and next with methyl orange.
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blogfast25
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A new 0.2 M HAc solution was prepared and standardised.
This afternoon, three runs for the heat of neutralisation of that solution with 0.2 M gave -55.6; -56.4; -56.9 kJ/mol, average ΔH = -56.3 kJ/mol,
SD = 0.7 kJ/mol. That's again much in line with previous values and literature. I believe the previous faulty previous runs to be caused by an unknown
standardising error, or at least that sounds like the most plausible explanation.
Assuming kmno4 is correct, then the mystery is essentially solved: if Ka is indeed the thermodynamical K for aH2O = 1. The heat of neutralisation of a
weak acid like HAc isn't much different from that of a strong acid and most of the deprotonation Free Energy manifests itself as entropy.
I will try and confirm that by running three neutralisation tests with HAc at near zero C, at 50 C and at 75 C, a total of four estimates.
What's still quite amazing is that for very, very weak acids like NH4+ the actual neutralisation enthalpy is practically zero: there the entropic
effect must be even higher...
Oooops... I've just noticed that kmno4's values are in kcal/mol, not kJ/mol:
Quote: Originally posted by kmno4  |
Thermodynamic data for reaction
CH3COOH(aq) -> CH3COO(-)(aq) + H3O(+)(aq)
are (in kcal/mol, rewritten from book by G.Kortum) :
ΔG = 6,49 ( = -RTlnKa) ; ΔH = -0,09 ; TΔS = - 6,58
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Converting:
ΔG = +27.1 kJ/mol ( = -RTlnKa) ; ΔH = -0.38 kJ/mol ; TΔS = - 27.5 kJ/mol
Due to the small value of ΔH this changes very little...
[Edited on 12-8-2010 by blogfast25]
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kmno4
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Yes - the book I mentioned ( translated) is “Lehrbuch der Elektrochemie” by G. Kortüm from ~1960 and in older references cited there they use
kcal.
I have noticed this:
6) The neutralisation enthalpy of NH3 with HCl was experimentally found to be -9.7 kJ/mol of water, so also much lower than - 57.3 kJ/mol. The
strength of both the NH3 and HCl solutions have been verified with acid/base titrometry.
Is this your experiment ?
From another book this is 12,72 kcal/mol (25 C, 0,3 M).
BTW.
There should be only small heat effect in reaction NH4(+) with NaOH.
It is because neutralisation of CH3COOH gives almost the same effect as neutralisation of HCl.
It would correspond to reaction NaCl + NaOH (zero heat).
BTW 2.
For NH4(+) acid:
ΔG=12,56 kcal/mol; ΔH=12,4 kcal/mol; TΔS=-0,16 kcal/mol
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Lambda-Eyde
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| Quote: Originally posted by blogfast25 | | Also, when using oxalic acid, do you use the dihydrate or anhydrous as a standard? If the latter how do you dehydrate? One source mentions possible
brown or black spots if you get the dehydration at just below 100 C (IIRW) wrong... |
http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv1...
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blogfast25
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| Quote: Originally posted by kmno4 |
6) The neutralisation enthalpy of NH3 with HCl was experimentally found to be -9.7 kJ/mol of water, so also much lower than - 57.3 kJ/mol. The
strength of both the NH3 and HCl solutions have been verified with acid/base titrometry.
Is this your experiment ? |
Yes, that was my experiment. I wouldn't put my hand in the fire on that value as it's a single data point, obtained at 0.5 M, with a slightly
different method. But it's not far from the value you quote.
For NH4+ and NaOH I found about DH = -5.0 kJ/mol.
All in all, it seems that entropy is explaining a lot here... Good, I like a bit of disorder ;-)
[Edited on 12-8-2010 by blogfast25]
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blogfast25
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Lambda-Eyde:
That was actually the chemtext I was referring to...
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blogfast25
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Method for primary base standard from washing soda
Firstly, some fundamental data about Na2CO3's interesting solubility behaviour can be found here:
http://www.genchem.com/properties.asp
Don't forget to click on the Fig. 2-1 hyperlink, it's the phase diagram for Na2CO3/H2O
Useful factoid: the Na2CO3 content of a mixture of A parts Na2CO3.10 H2O and B parts of water can be calculated from:
% (anh.) Na2CO3 = 37.1 x A / (0.629 x A + B)
1. 200 g of washing soda were dissolved into 200 ml of tap water (about 23 %, based on Na2CO3.10 H2O) and brought to about 80 C quickly. Do NOT boil
to avoid loss of CO2 (hydrolysis). Filter till clear. I preheat my filter to about 100 C to avoid any crystalisation on the filter. In my case there's
a fine, white powder left on the filter which I presume is mainly NaHCO3. I will verify that.
2. Allow the filtrate to cool or force cool it, then cool further on ice bath. Crystals of relatively pure Na2CO3.10 H2O form and grow visibly. I
stopped cooling at about 14 C which gave a nice crop of wholesome crystals.
3. Decant off the supernatant liquid and collect the crystals on a clean tea strainer or ceramic filter. I use a SS fine mesh (+/- 0.5 mm) strainer.
Allow to drip, then wash with small amounts of ice cold deionised water (DIW). These crystals are undoubtedly already sufficiently pure for much
analysis and synthesis, so set some aside.
4. For the second crytallisation, we take advantage of the fact that soda's maximum solubility is 33.2 % at about 35 C and drops steadily towards 0 C.
Because the slush contains so much liquid water, not much needs to be added. I've had success with adding 1, 2 or even 0 parts of DIW to 10 parts of
Na2CO3.10 H2O/water slush. I think 2 parts probably works best. Heat that sludge au bain marie to about 35 C, making sure all crystals
dissolve. Allow to cool or force cool on ice bath. Due to purity, often the solution is reluctant to crystallise and you may need to add a crystal
from the first crop: the solution then beautifully crystallises all out at once.
5. Repeat 3. Allow time to drip to almost dryness. Dehydrate to constant weight in oven at 110 C. Use a plastic (PP) or silicone cup for this:
anhydrous sodium carbonate sticks like mad to glass. Use immediately. I do believe it can be stored in CaCl2 desiccator for some months.
6. Quality test: dissolve some in DIW, a clear solution has to be obtained, without turbidity or insoluble bits. I've become so paranoid about this I
leave the test solution standing overnight for a final check.
[Edited on 14-8-2010 by blogfast25]
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entropy51
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blogfast25, when you oven dried the decahydrate to constant weight, did it lose the theoretical amount of water?
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blogfast25
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| Quote: Originally posted by entropy51 | | blogfast25, when you oven dried the decahydrate to constant weight, did it lose the theoretical amount of water? |
In a certain sense the decahydrate doesn't really exist: the sludge from the second crystallisation is impossible to dry without the some of the
crystals losing some water. I see this on all utensils used in contact with the Na2CO3.10 H2O solutions: apart from some clear, well formed crystals
forming, in air quite a bit of them effloresce. The commercial 'washing soda' also looks like a mixture of decahydrate and lower hydrates and
commercial products usually mention 'min. x % sodium carbonate decahydrate', never 'pure sodium decahydrate'. Holleman confirms the effloresence.
Pure Na2CO3.10 H2O may possibly obtained by gentle drying at low temp. under mild vacuum? It's a shame because the decahydrate has a much higher
gravitational lever, better for primary standards..
So the decahydrate second crop is Na2CO3.10H2O/liquid water sludge. This is then oven dehydrated. Ergo I cannot know whether the obtained weight loss
is in accordance with the theoretical weight loss.
[Edited on 15-8-2010 by blogfast25]
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blogfast25
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Well, well. At first glance the white insoluble matter found in my washing soda doesn't appear to be NaHCO3.
I took 6 g of the still wet stuff and reacted it with 10 ml 20 % HCl. Much effervescence AND considerable heat evolves but not everything dissolves.
The heat is strange: it would seem to exclude NaHCO3 and CaCO3 (I checked: a bit of limestone with 20 % HCl doesn't generate much heat on dissolving
in HCl).
The solution was filtered and diluted a little and strong (5 M) NaOH was added: a white, gelatinous precipitate dropped out, quite a bit of it too.
The alkali metals wouldn't do that and Ca(OH)2 isn't particularly gelatinous as a precipitate, IIRW. It was filtered and washed slightly, then scraped
of the filter. Strong Na2CO3 was added to the filtrate, which didn't respond.
To the precipitate was then added strong Na2CO3 and the solution was heated to test the displacement Ca(OH)2 (s) + Na2CO3 (aq) --> CaCO3 (s) + 2
NaOH (aq) but apart from the gelatinous precipitate dispersing through the solution, nothing much else happens.
I'm now drying some of the white stuff for some further testing...
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watson.fawkes
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In the US at least, much washing soda is derived from trona, so you might consider mineral-type contaminants such as borates,
silicates, sulfates etc.
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arsphenamine
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| Quote: Originally posted by watson.fawkes | | In the US at least, much washing soda is derived from trona, so you might consider mineral-type contaminants such as borates, silicates, sulfates etc.
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I would guess primarily silicates,
similar to sodium silicate (aka "water glass") which
has been used to clarify turbid waste water.
Particularly, it precipitates metals from solution.
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blogfast25
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watson and arsphenamine:
Sure but what I' seeing isn't really consistent with those...
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watson.fawkes
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Shale is another impurity of trona, but I think that's even less likely. Worth a guess, I suppose. The only
other thing I've got is that if you've got output from a Solvay process, there's a different class of impurities.
[edit] P.S. It's also conceivable that your washing soda isn't even notionally pure, that it's been compounded for laundry. I mention this with the
full disclosure that occasionally I have forgotten to check that something is plugged in.
[Edited on 15-8-2010 by watson.fawkes]
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blogfast25
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| Quote: Originally posted by watson.fawkes | Shale is another impurity of trona, but I think that's even less likely. Worth a guess, I suppose. The only
other thing I've got is that if you've got output from a Solvay process, there's a different class of impurities.
[edit] P.S. It's also conceivable that your washing soda isn't even notionally pure, that it's been compounded for laundry. I mention this with the
full disclosure that occasionally I have forgotten to check that something is plugged in.
[Edited on 15-8-2010 by watson.fawkes] |
It's possible. It's also possible it's straight Solvay soda. It's only 1 £/1 kg in retail.
About half of the impurity dissolves in strong HCl, the other half doesn't. I'm now investigating the first half...
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blogfast25
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The packet states 'min. 30 % decahydrate', which is definitely referring to the (poor) state of hydration. Endearingly it also states:' no
petrochemicals', so in all likelihood it's not compounded.
Some of the dried water insoluble residue was treated with strong HCl and about half of it does dissolve in it with effervescence, the other half
doesn't. Filtering and diluting a bit, the filtrate was then split into 2 equal portions. One was neutralised with strong Na2CO3: no precipitation
even at pH = 11. The other was neutralised with strong NaOH and a whitish, gelatinous precipitate forms. I'll check the latter for amphoterism and
complexation with NH3.
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kmno4
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| Quote: Originally posted by blogfast25 | | Quote: Originally posted by kmno4 |
6) The neutralisation enthalpy of NH3 with HCl was experimentally found to be -9.7 kJ/mol of water, so also much lower than - 57.3 kJ/mol. The
strength of both the NH3 and HCl solutions have been verified with acid/base titrometry.
Is this your experiment ? |
Yes, that was my experiment. I wouldn't put my hand in the fire on that value as it's a single data point, obtained at 0.5 M, with a slightly
different method. But it's not far from the value you quote.
For NH4+ and NaOH I found about DH = -5.0 kJ/mol.
All in all, it seems that entropy is explaining a lot here... Good, I like a bit of disorder ;-)
[Edited on 12-8-2010 by blogfast25] |
This is rather very big difference between your -9,7kJ/mol and
literature -12,72 kcal/mol (=53,2 kJ/mol) value.
BTW
For reaction in kind of "salt1+salt2 -> something" you can estimate enthalpy of rection from available tables of thermodynamic data for aquaions
and pure substances. In this way you can predict (to some extent) heat of rections: NH4(+) + OH(-) -> NH3 + H2O; CO3(2-) + 2H(+) -> CO2 + H2O
and many other.
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blogfast25
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| Quote: Originally posted by kmno4 |
This is rather very big difference between your -9,7kJ/mol and
literature -12,72 kcal/mol (=53,2 kJ/mol) value.
BTW
For reaction in kind of "salt1+salt2 -> something" you can estimate enthalpy of rection from available tables of thermodynamic data for aquaions
and pure substances. In this way you can predict (to some extent) heat of rections: NH4(+) + OH(-) -> NH3 + H2O; CO3(2-) + 2H(+) -> CO2 + H2O
and many other. |
I'm inclined to stand by my value, as I remember the experiment yielding a very low ΔT. But I'm willing to repeat it at 0.5 M. It will take a few
days because of the need to restandardise NH3 and other commitments. I'll also check the raw data of the experiment.
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blogfast25
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As regards the white, water-insoluble residue in my washing soda, it's almost certainly MgO or Mg(OH)2, added probably as an anti-caking agent. I
isolated some more and washed it carefully. It dissolves effortlessly in 1 M H2SO4, with mild turbidity and generation of considerable heat but no
bubbles. That's been filtered and the salt will be crystallised tomorrow.
The stuff was also tested with very strong NaOH: it's not amfoteric. It doesn't react with strong NH3 either.
A cation that doesn't form an insoluble chloride, nor insoluble sulphate, nor insoluble carbonate but does precipitate a white, gelatinous hydroxide
(presumed)? That looks very much like Mg2+...
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blogfast25
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Well, using a new bottle or standardised NH3, I now get a value of -64.7 kJ/mol for NH3 + HCl at about 0.2 M! So that's over-shooting but shows the
previous value of -5.0 is erroneous (although I can't figure out why from the raw data).
The value for NH4+ + NaOH will also have to be repeated.
But the value of -64.7 kJ/mol was obtained in slightly dodgy circumstances: it turned out that the concentration of NH3 in the latest bottle of
'household NH3' is only 0.37 w% (0.216 M)! The previous bottle was bad enough at 4 w% but this takes the biscuit... Dry distillation of garden grade
ammonium sulphate with NaOH, here I come!
My local hardware store stocks some citric acid monohydrate ('no additives' it says), so I'll see if that's any use as a primary acid standard.
Interesting solubility (and other) data on citric acid here (third search result from top, a *.DOC download):
http://www.google.co.uk/#hl=en&q=citric+acid+solubility&...
The solids content of my washing soda turns out to be about 38.6 w%, so fairly close to the 'theoretical' value of 37.1 w%.
[Edited on 18-8-2010 by blogfast25]
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blogfast25
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Well, my first efforts at recrystallising citric acid monohydrate haven't been very successful. The raw product is a neat, white, well formed
crystalline product, seemingly of 'good quality'.
Citric acid is extremely water soluble (over 90 w% @ 100 C and still over 50 w% @ 10 C) and forms syrupy solutions. Cooling these down to RT (say from
80 C saturated) yields no crystals but allowing cooling at about 5 C overnight yielded a nice crop. The problem is that the supernatant liquid is
still highly concentrated. Washing with cold DIW is possible but drying the wet crystals (of presumably monohydrate) is hard: at 50 C more syrupy
solution is formed between the crystals and the wash water and as the syrup is highly concentrated, driving off water is nigh impossible at
atmospheric conditions.
I will titrate a known quantity of (raw) citric acid against a secondary NaOH standard, using the first two equivalence points, just to see.
The white residue in my washing soda is definitely MgO or Mg(OH)2. Dissolved in H2SO4 and filtered I obtained nice, big clear crystals. After washing
and redissolving in DIW, adding Na2CO3 yields some fizz and upon heating the white, basic carbonate Mg(OH)2.3MgCO3.3 H2O (acc. Holleman) precipitates
from the solution.
The only test left is for MgNH4PO4.6H2O but I've no H3PO4 at hand.
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blogfast25
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The value of the reaction heat of NH4+ +OH- ---> NH3 + H2O has been confirmed at -5.2 kJ/mol. Using 400 ml of standardised NH4+ (about 0.4 M), 400
ml containing a slight excess of NaOH, the temps. were: TNaOH = 22.3 C, TNH4+ = 17.9 C, Tend = 20.3 C. Mols of NH4+ neutralised = 0.172 mol. ΔH =
-5.2 kJ/mol tallies perfectly with three other values of around -5.0 kJ/mol obtained with NH4Cl at 20 C, (NH4)2SO4 at 20 and at 0 C and now this one.
There's just no heat in it.
Music to my ears because I'll be dry distilling (NH4)2SO4 with NaOH and can do without runaway heats...
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