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Sulaiman
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How to measure sulphuric acid concentration ?
Other than by careful, tedious, wasteful titrations,
how can I easily and quickly determine the concentration (% w/w) of H2SO4 in the range 93% to 98% to the nearest 1%?
Density measurements are not useful as they are very sensitive to errors near azeotropic concentrations,
Electrical conductivity similarly seems useful only from 96% to 100%
Sonic velocity measurement could work, but too complicated.
so
what have I missed ?
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j_sum1
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How is a titration wasteful?
To get 1% accuracy you need to be able to measure to +/- 1%. Let's double this target for initial weighing -- that means drawing out about 2g (or
just over 1mL) of acid and measuring its mass to the nearest 0.01g. Dilute to a reasonable volume using whatever volumetric flask you have. In
general, titrations give better accuracy the more dilute they are. If you have a 500mL or 1000mL flask you are probably in the zone. Throw some in
your burette and titrate away. Good technique and repetition should get you a result within 0.5% accuracy fairly easily.
Running the errors:
minimum:
0.995(weighing)×0.995(titration)=0.990025
maximum:
1.005×1.005=1.010025
So, result within 1% for the expense of 1mL of acid.
Getting error to below 0.1% will take a bit more effort and expense. But I am pretty happy with 1% for anything I do at home.
(All this assumes that you have something decent to titrate against. NaOH is no good since it absorbs moisture from the air. Probably the easiest
standard is to prepare some anhydrous Na2CO3 from bicarbonate. This does introduce weak base issues but should be manageable since H2SO4 is a strong
acid.)
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Sulaiman
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D'oh .... thanks j_sum1,
I had a mental lapse - I forgot that I can dilute before titration 
(so no longer wasteful, just tedious ... I can accept that 
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JJay
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Nobody knows what kind of needle to use to handle highly corrosive substances with a glass syringe?
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clearly_not_atara
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It appears from some Googling that nickel alloys such as Hastelloy and Inconel are suited for this purpose, as is 316 (molybdenum) stainless steel. I
suspect that copper will tolerate bromine to some extent as well:
http://www.balseal.com/sites/default/files/tr60c_02070713261...
[Edited on 29-3-2017 by clearly_not_atara]
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JJay
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That's interesting. This link gives both Hastelloy and 316 steel a D rating for handling bromine, the same rating given to aluminum: http://www.graco.com/content/dam/graco/ipd/literature/misc/c... The substances you mention (with the exception of copper) are mainly intended to
handle high temperatures; I'm not sure that they would resist hydrochloric acid well.
I'm not really sure where to get copper needles and doubt they would hold an edge for very long... I was thinking something more like ceramic or
perhaps platinum-iridium alloy might work well but have never seen needles made of such exotic materials....
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clearly_not_atara
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They are alloy families after all. I guess that Balseal uses grades specially formulated to handle chemicals, whatever that might require. I believe
that copper reacts quite slowly with bromine at room temperature. But from the looks of that you might need a specially-designed alloy.
I think I'd bet on a copper needle before glass. Who knows? These are stainless:
https://www.hamiltoncompany.com/~/media/Files/Syringes%20and...
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yobbo II
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In an all glass syringe what is the plunger made from? glass
Just wondering how is there a good seal for plunger to tube wall.
If you obtain a glass tube and heat in a flame it is very easy to make your own 'needle' by just pulling on the softened tube and extending it out
untill you have a very thin neck. It it very easy to do.
They you have to get a plunger for the tube which might not be so simple.
Having looked them up on ebay I guess glass is the answer.
http://www.ebay.com/itm/GLASS-SYRINGE-1cc-LOT-OF-2-LUER-TIP-...
It would be easy to heat a borosilicate syringe (output end) in a hot flame and pull untill you have a thin enough section for you 'needle'. Some of
the syringe would then be unusable I guess but should not be a problem.
[Edited on 29-3-2017 by yobbo II]
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Sulaiman
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30 months ago I made a 0.99 (+/- <0.5% ) Molar solution of Na2CO3 for titrations.
It has been kept in a (nearly) airtight hdpe bottle since then, with no visible change.
Opinions on the likelyhood of the solution still being 0.99M (+/- <0.5%) ?
....................................
I would go for an all- glass syrynge and a piece of glass tubing that just fits over the Luer tip, with the end stretched to a point ... as above.
Epoxy resin is not immune to bromine, but a thin sealing layer between the glass tubing and the Luer nozzle would expose only a tiny area to the
bromine.
I guess 
I assume that even if a little glass accidentally gets into the reaction it will not matter, as the reactions take place in glassware.
__________________________________________________
EDIT: based on the reply below, I shall make up a fresh solution.
[Edited on 30-3-2017 by Sulaiman]
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byko3y
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All basic solutions absorb CO2 from air and water. ph of sodium acetate solution practically is 0.5 units lower than theoretical value.
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clearly_not_atara
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Can't you just add eg calcium acetate soln and weight the ppt?
JJay: Check your link again. Hastelloy C gets an A rating for anhydrous bromine and bromine water it appears. However 316 steel does poorly.
[Edited on 30-3-2017 by clearly_not_atara]
[Edited on 30-3-2017 by clearly_not_atara]
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Texium
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Would 95% formic acid and paraformaldehyde be suitable to use in an Eschweiler-Clarke reaction? I'm looking to make some dimethylaniline. Thanks.
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JJay
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Quote: Originally posted by clearly_not_atara  | Can't you just add eg calcium acetate soln and weight the ppt?
JJay: Check your link again. Hastelloy C gets an A rating for anhydrous bromine and bromine water it appears. However 316 steel does poorly.
[Edited on 30-3-2017 by clearly_not_atara]
[Edited on 30-3-2017 by clearly_not_atara] |
Huh, that's interesting... I don't remember seeing Hastelloy C before, but it must have been there... perhaps I posted the wrong link... anyway, it
still gets a D rating for hydrochloric acid... well... actually, it gets a B rating for 37% hydrochloric acid but a D rating for hot 37% hydrochloric
acid and a D rating for cold 37% hydrochloric acid. So I guess Hastelloy C works ok but suboptimally for lukewarm hydrochloric acid??
[Edited on 3-4-2017 by JJay]
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Eddygp
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Is there any simple way to dehydrate ketones to the corresponding alkyne?
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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Texium
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Only way I know of is chlorination with PCl5 and then double elimination of the gem. dihalide
with a strong base, but I doubt that qualifies as "simple"
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CuReUS
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| Quote: Originally posted by zts16 | | Only way I know of is chlorination with PCl5 and then double elimination of the gem. dihalide
with a strong base |
how do you know that it wouldn't give you an allene instead of an alkyne ?
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Geocachmaster
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I was attempting the oxidation of toluene to benzoic acid, this is pretty well known. The reaction should proceed something like this:
C7H8 + 2 KMnO4 --> C7H5KO2 + KOH + 2 MnO2 + H2O
I relfluxed 3.64g of KMnO4 and 45ml of water with an excess of toluene (9 or so milliliters) for a few hours, untill all the purple color had
disappeared. I filtered off the mixture from all the manganese dioxide and then separated the extra toluene in a sep funnel. I then acidified the
~40ml of solution that remained with 6g of sulfuric acid that was mixed with a little more than 20ml of water. Theoretical yield, 1.40g benzoic acid.
This should have been more than enough to neutralize all the KOH and precipitate benzoic acid. There was slight effervescence and the solution
remained clear. The pH had gone from ~10 down to 1. No precipitate of acid, which should have happened given the solubility of benzoic acid in water
and the low pH. It was left for 30 minutes with no change.
The MnO2 left and loss of purple color means the permanganate oxidized something. Most of the toluene was gone too. I'm really confused, why no
benzoate? Has this happened to anyone else? Any help is appreciated.
[Edited on 4/16/2017 by Geocachmaster]
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j_sum1
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I was just watching an interesting video on Thallium chemistry https://www.youtube.com/watch?v=jw5S6mtx338
(I think this is fluorescence's work.)
It got me wondering, what exactly is the difference between Thallium (I) triiodide and Thallium (III) iodide? I get that there are different
oxidation states for both the Tl and the I, but that would seem to me to be largely academic once the product was isolated. Can anyone enlighten me
at all?
(I remember asking a similar question about PbO2 which, in times past, was referred to as lead peroxide. But it is now more accurately referred to as
lead (IV) oxide.)
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anewsoul
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I finally found a procedure for oxidizing o-nitrotoluene to o-nitrobenzaldehyde but it says it uses sulfuric acid that is 30-40° Be'. What is that supposed to mean? Is it some
kind of unit for concentration? I've never seen it before.
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Sigmatropic
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Be' or Baume is a density scale. https://en.m.wikipedia.org/wiki/Baumé_scale.
30-40 baume H2SO4 thus is 1.26-1.38 g/mL which would be about battery acid. For a conversion table see https://www.google.nl/url?sa=t&source=web&rct=j&...
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anewsoul
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Thanks for the help! Hopefully the reaction works out.
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clearly_not_atara
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| Quote: | | I relfluxed 3.64g of KMnO4 and 45ml of water with an excess of toluene |
A great way to synthesize benzyl alcohol. You need an excess of oxidant to get complete oxidation to the acid. You may have produced some
benzaldehyde as well.
| Quote: | | I finally found a procedure for oxidizing o-nitrotoluene to o-nitrobenzaldehyde but it says it uses sulfuric acid that is 30-40° Be'.
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I recommend isopropyl nitrite instead:
http://acta-arhiv.chem-soc.si/52/52-4-460.pdf
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anewsoul
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I've seen this method but it looks like too much trouble for the yield it gives. Also it requires sodium methoxide and I don't have sodium and also I
would have to make the isopropyl nitrite and I don't have any nitrite salts. Might be a better method, but I have plenty of sulfuric acid and
manganese salts to prepare the dioxide.
[Edited on 17-4-2017 by anewsoul]
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Panache
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When reducing oximines back to their respective ketones or aldehydes aqueous dithionite is said to work well at rt.
Anyone used this and if so what were your reaction conditions, is time dithionite conc, dionithionite excess?
[Edited on 21-4-2017 by Panache]
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clearly_not_atara
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https://erowid.org/archive/rhodium/chemistry/redamin.dithion...
| Quote: | General Procedures
Method A.
The oxime (20 mmol) was mixed with water (15 ml) containing sodium dithionite (28 mmol). The suspension was stirred overnight at room temperature.
(Warming to 40°C reduced reaction times to several hours.) In some cases, a precipitate formed.This product was very high melting and, on treatment
with 2 M hydrochloric acid, liberated the carbonyl compound and sulfur dioxide. It was therefore assumed to be the bisulfite addition compound of the
carbonyl compound and was not isolated. A slight excess of 2 M hydrochloric acid was added to the reaction mixture and nitrogen was bubbled through
the mixture to expel the sulfur dioxide. Solid sodium carbonate was added carefully to alkalinity; the aqueous mixture was allowed to stand for 30 min
and was extracted with ether (2x10 ml) which was dried (MgSO4) and evaporated. The residue was essentially pure carbonyl compound (by t.l.c.)
Method B.
The reaction described under Method A was performed in the presence of sodium hydrogen carbonate (28 mmol). Cleavage by means of this modification
appeared to proceed considerably faster. The usual workup gave the carbonyl compound in comparable yield. |
Notably there are some claims that sodium dithionite reduces oximes, however, upon inspection, dithionite is only reported to reduce alpha-keto
oximes, produced by nitrosation of the ketones. It's worth noting that alpha-keto oximes tend to be present as the nitroso compound rather than the
oxime, unlike all other oximes, and the nitroso compound is more susceptible to reduction.
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