Sugar reacts with sodium hydroxide to produce carbon monoxide. It seems like a slow reaction, though may be speeded by heatinig. Can anyone speculate
on what's going on?
There are of course more effective ways of preparing this gas but in principle this seems to have the advantage of using readily available materials.Ozone - 31-1-2009 at 20:44
Sucrose does not do much in alkali (at room temperature). Please use the search engine in the upper left hand of the page (under the flaming sigil).
Try looking up something like formic acid...
Cheers,
O3JohnWW - 31-1-2009 at 21:01
What do you want to prepare CO for? The stuff is deadly poisonous to inhale, comparably to HCN, (CN)2, H2S, and H2Se. If you want to handle fairly
large quantities of it, use a fumehood, and preferably wear breathing apparatus with bottled air. There are three possible chemical uses that I can
think of for it - for the refining of Ni from crushed ores in the form of ultrabasic rocks, with which it readily forms the volatile tetracarbonyl
Ni(CO)4; to produce carbonyls of other transition metals with partly empty d orbitals; and, being the anhydride of formic acid, to react with alcohols
or phenols to make formate esters.
If you could do with formic acid instead in order to make formates, this can be obtained by the oxidative cleavage of many sugars, particularly
glyceraldehyde, with HIO4 or Pb(CH3COO)4. The formic acid can then, if desired, be dehydrated to CO with H2SO4.chemkid - 1-2-2009 at 18:55
Someone around here was preparing sodium gluconate or sucrate (cant remember which) with sodium hydroxide and glucose or sucrose. Never found much
information on synthesis, but i would be interested if anyone had any.Ozone - 1-2-2009 at 19:38
Sucrates are salts of gluconic or glucaric acids (that is, the hexose acids). In industry, the "sucrates" or "saccharates" are no such things
(jargon). These are usually complexes (with sucrose, anyway), not salts of the original carbohydrates (e.g. tricalcium sucrate). This is easily and
completely reversible (the reducing sugars are destroyed yielding the brown/black tar that colors most impure mixtures of this kind).
Synthesis is easy. Just add lime (or other alkali), a little heat, and wait. Sucrose increases the solubility of Ca(OH)2 by a factor of roughly 12,
relative to water at the same temperature. This is due to the complexation. This mixture is frequently used, in place of lime, to effect the
clarification of cane juice (saccharate liming).
So, non-reducing sugars (like sucrose) don't do much whilst reducing sugars (glucose, fructose, etc.) are easily destroyed. The free carbonyl sets up
the tautomerization required for reverse aldol scission and subsequent condensation reactions to occur.
Cheers,
O3chief - 2-2-2009 at 04:55
The good old chembook states as reaction for CO-preparation : H2SO4 + formic acid ;
The danger is, that CO has no smell ; you don't feel it coming, just fall asleep, and gone you are ... ; that's how accidentally people make it to the
eternal hunting grounds ever again, by malfunctioning ovens etc.woelen - 2-2-2009 at 05:46
It indeed can very easily be made from formic acid. Drip concentrated sulphuric acid into 60% formic acid and while the liquid heats up (reaction of
water and H2SO4), it also starts foaming, giving bubbles of CO. The reaction is quite fast and smooth. In this way I have made a few test tubes full
of CO.
As stated by others, CO is VERY dangerous, much more so than Cl2 and many other toxic gases, because it has no warning properties, nothing at all.
Even if the gas is present in deadly concentration, you won't notice, until you suddenly loose conscience and fall down. There is no smell, no fumes,
no irritation, nothing. So, do this experiment in a very good fume hood, or work outside on a breezy day.Jor - 2-2-2009 at 06:56
I do really think people overrate CO.
Although it is colorless and odourless, you have to inhale quite high concentration to die or collapse.
Typical STEL , maximum exposure for a short time, is 300ppm. Then you won't suffer effects! Now if your works outside, on a windy day, there is no way
that concentration will be reached. You will have to release 300mL of pure CO per m3 to reach an exposure that is not even harmful at short periods!
I read somewhere that on the streets of mexico city, CO concentration is about 100ppm!
Now try that with other common poisonous inorganic gasses.
When you work outside with chlorine, do ever suffer from severe irritation? I did not, when I worked without a hood, and there was not much wind. This
means concentrations were lower than 10ppm, wich is described as unbearable. And CO will dissappear even more quickly than Cl2, as it is much lighter.
So myself, I would not worry too much. If you generate small amounts of CO gas, outside, you are not likely to get injured or worse, unless you stick
your nose right above the flask (where it is concentrated). But in that case, chlorine will kill you as well, having no time to respond.
Even though it is one of the number 1 chemical killers, it always kills people that are unaware of it's presence, and concentration may be very high,
because houses are not ventilated, especially when it's cold. Now if you generate some in a controlled fashion, and you KNOW how much is released (so
you can calculate potential concentration), I do not see any accidents happen. And headache/tiredness are symptoms, so you are not totally unaware you
are inhaling too high concentrations, if you KNOW it is there.
I would work inside with small volumes of CO, without hood. We all stand in front of a barbeque as well, for an hour. Ever thought of CO concentration
you are inhaling when you barbeque?
Ofcourse, my advice is still to be very careful, and treat it as potentially dangerous gas. But treating it like it is HCN is exaggerating.
A table:
Concentration Symptoms
-35 ppm (0.0035%) Headache and dizziness within six to eight hours of constant exposure
-100 ppm (0.01%) Slight headache in two to three hours
-200 ppm (0.02%) Slight headache within two to three hours
-400 ppm (0.04%) Frontal headache within one to two hours
-1,600 ppm (0.16%) Dizziness, nausea, and convulsions within 45 minutes. Insensible within two hours.
-3,200 ppm (0.32%) Headache, dizziness and nausea in five to ten minutes. Death within 30 minutes.
-6,400 ppm (0.64%) Headache and dizziness in one to two minutes. Death in less than 20 minutes.
-12,800 ppm (1.28%) Unconsciousness after 2-3 breaths. Death in less than three minutes.
So for example after breathing 200ppm for 2 or three hours gives you a slight headache. Breathing chlorine, NO2, HCN, H2S, etc would have killed you
by then.
[Edited on 2-2-2009 by Jor]
[Edited on 2-2-2009 by Jor]chief - 2-2-2009 at 07:43
Quote:
And CO will dissappear even more quickly than Cl2, as it is much lighter.
CO has nearly the same dansity as air ... so it mixes perfectlywatson.fawkes - 2-2-2009 at 10:22
Quote:
Originally posted by Jor
Even though it is one of the number 1 chemical killers, it always kills people that are unaware of it's presence, [...]
An excellent point, one that I'll reinforce by pointing out that CO detectors are inexpensive (< USD 20) and they can go right
next to the fire detector in the lab.zephram - 7-2-2009 at 14:04
My thanks to all for the contributions. All good stuff though unfortunately mostly fairly wide of my initial question. Yes, I had used the search
engine before posting, both on this site and the web in general. I found no shortage of the very familiar stuff on toxicity and standard methods or
preparation, but only brief mentions of my original enquiry which was about the preparation of carbon monoxide from sugar (sucrose) and
sodiun/potassium hydroxide.
"Another laboratory method to generate carbon monoxide is reacting sucrose and sodium hydroxide in a closed system."
I'd be interested to hear if anyone has more practical details.Picric-A - 7-2-2009 at 14:13
That probably requires high temperatures and pressures, well beyond the ability of a home chemist.
If you need CO stick with the ideas mentioned in the thread ( formates + conc H2SO4)Formatik - 7-2-2009 at 17:00
I haven't seen this method in any of the older literature, which describe using HCOOH or C2H2O4.2H2O with conc. H2SO4 for laboratory methods of
preparation.zephram - 7-2-2009 at 19:10
> That probably requires high temperatures
> and pressures, well beyond the ability of a
> home chemist.
Well maybe ... or maybe not. Heating a couple of grams of sugar with a simar amount of NaOH and water in a test tube produces gas that burns with a
blue flame.
Also some Materials Safety Data Sheets for caustic cleaners warn of using it in tanks used that may contain traces of sugar solutions.garage chemist - 7-2-2009 at 19:31
The cheapest method for making CO would be to put charcoal in a tube furnace, heat to 1000°C and lead in CO2 at one end. At the other end you get a
mix of CO (mostly) and CO2. Wash out the latter with aqueous NaOH and you have your CO.
The reaction of CO2 with C is endothermic, hence the tube furnace, and an equilibrium.Picric-A - 8-2-2009 at 10:40
@ zephram - There are many gasses that burn with a blue flame,
Do a CO test on the gas produced, and find out for sure, I am currently away from my lab so its up to you.
Personally, however, i try to stay well clear of CO whenever possible,Eclectic - 8-2-2009 at 10:53
Reaction of NaOH with sucrose to produce CO is likely a misinterpritation of this, mentioned in a wiki, and widely copied without review.Picric-A - 9-2-2009 at 00:04
Great link Eclectic!
how exactly did you produce a 'gas that burns with a blue flame' ,zephram ?
Is seems unlikely to work with normal sugar (sucrose), or did you use another sugar?Formatik - 9-2-2009 at 11:12
A burning pale-blue flame is suggestive of CO (note its LEL: 12.5, UEL: 74.0 - that's almost like hydrogen!), but this is a pretty primal test, though
still not entirely indicative. More qualitative, and still simple tests rely on its reducing power. It will blacken silica gel (or PdCl2-soaked
paper). A blue-green color from silica gel which is soaked with a solution of I2O5 in conc. H2SO4. Light purple aq. KMnO4 solutions are said to go
colorless. If there is CO, then it might be impure. CO2 might be one of those impurities, testing for CO2 is as simple as leading the gas through a
clear Ca(OH)2 solution, where CaCO3 precipitates.
The reason why CO is often placed with compounds like HCN, etc. is because they are all blood poisons, note that does not mean as deadly as HCN, just
going by toxicological action. In the case of CO, it will react with the hemoglobin to form carbon oxide hemoglobin (HbCO), this reaction has a far
higher affinity than that competing one of O2. CO is an insidious poison, because of its odorless nature. It should not be underestimated, but also
not overestimated.chief - 9-2-2009 at 14:00
The cheap detectors seem all to only beep at a limit-value, but reading the conc. is not possible.
A good one, with 1% accuracy, I couldn't find below > 100 EUR ; such a one would be much more useful ...
From conc. CO one breath may be deadly, within short time ...Jor - 9-2-2009 at 14:07
Same for chlorine, bromine, etc.
The fact is that, at concentrations where CO would be lethal, it's detectability is not important anymore as any other toxic gas would knock you down
at the same concentration as well, without you being able to respond. When CO is lethal, other gasses are as dangerous as well, as they kick you down
in one breath.
In neither cases you have time to respond. At least you get headache symptoms with CO.
I'm not saying CO is something to play around with carelessly, but please look at the concentration table I posted, those concentrations are not
easily reached. I would be FAR more afraid of Br2, NO2, H2S, HCN, to wich CO seems like a toy to me.
It surprises me that some gasses are overestimated by people and some are underestimated. I think many people overestimate things like CO, while H2S
is greatly underestimated (although not anymore on this forum, because of Klute's scary story).
[Edited on 9-2-2009 by Jor]FrankRizzo - 10-2-2009 at 05:06
Wasn't there a fad in Japan not too long ago where people were mixing bath salts and some other OTC product in order to commit suicide by CO
poisoning?
EDIT:Nevermind. It was hydrogen sulfide.
[Edited on 2/10/2009 by FrankRizzo]chief - 10-2-2009 at 06:30
But CO is breathed in without notice - just like air. At the same time chlorine etc. will probably hurt the nose quite badly and give some warning
that way ...
Probably it's not easy at all to breathe in the deadly chlorine-dosage ; long before one would get away from it, just by the smell ...
[Edited on 10-2-2009 by chief]Jor - 10-2-2009 at 10:13
I agree that CO is breathed without notice but yo didn't get my point. We compare these gasses, so compare same concentrations...
Chlorine gas at say 100ppm is very dangerous, but you can indeed flee. But at the same conc. CO is not even dangerous... You can breath these levels
for hours and get a slight headache.
Now at concentrations where CO is lethal, say 1000ppm+ (and that even takes a very long time, hours, before you die), chlorine doesn't warn anymore
and kicks you down on the ground in one breath. So it doesn't give warnings anymore as well.
Therefore I can easily conclude Cl2 is more hazardous. To be honest, when you KNOW you generate CO, there is very little chance you die, when I look
at the ppm data/health effects posted above. I would generate a liter of CO inside, no problem. Just try that with chlorine, and compare the effects.
CO from Na Nitrate and Charcoal
dann2 - 27-3-2009 at 13:58
Hello,
Some stuff here on making CO from a cartridge of Sodium Nitrate and Charcoal. May be suitable for lab. work.
Would Ammonium Nitrate do instead of the Na Nitrate?
Dann2
Attachment: CO.zip (401kB) This file has been downloaded 583 times
benzylchloride1 - 28-3-2009 at 21:17
I have read about also the use of oxalic acid with sulfuric acid for the preparation of carbon monoxide. Oxalic acid can be obtained cheaply. Formic
acid is better, but I can only obtain it from a scientific supply. I have read about a procedure for making the hexacarbonyls of Cr, Mo, W from the
anhydrous metal chlorides dissolved in ether in the presence of ethyl magnesium bromide by bubbling CO through the solution. The reaction uses about
6-8 liters of CO, and produces about two grams of the carbonyl from about 10g of the anhydrous metal chloride. The yield sucks, but it sure beats
spending $30+ per gram for these organometallic compounds. I am planning on running this reaction once I get the tube furnace that is needed to
prepare the metal chloride from the oxide and gaseous disulfur dichloride. I am fascinated by the chemistry of carbon monoxide complexes of transition
metals (carbonyl complexes) I will post more about my experiments as soon as I conduct them.Chromium, Molybdenum, Tungsten Carbonyls. The reaction
could be conducted by slowly passing dry CO through the reaction mixture until the reaction ceases to absorb CO. A mineral bubbler could be attached
to the inlet and outlet to indicate the flow rate in and indicate when the reaction is finished. A slight excess of oxalic or formic acid would be
used for generating the CO. CO can ve handled safely in a good fume hood. The apparatus should be completely sealed and the outlet lead into the
reaction mixture. The CO should be generated slowly so that a large excess is not vented to the atmosphere of the fume hood .
Procedure from Vogel for generating CO:
Carbon monoxide. This gas is readily prepared by the action of
concentrated formic acid (sp. gr. 1-2 ; about 90 per cent.) upon concentrated
sulphuric acid at 70-80°. The apparatus of Fig. / / , 48, 5 is
recommended. The distilling flask (500 ml.) is immersed in an oil bath
maintained at 70-80°, and is connected to two wash bottles containing
concentrated sulphuric acid. 125 grams of concentrated sulphuric acid
are placed in the distilling flask and 85 g. of the strong formic acid are
slowly added from the dropping funnel; a steady stream of gas is evolved.
The resulting carbon monoxide may contain traces of carbon dioxide and
sulphur dioxide . these impurities may be removed, if
desired, by passage of the gas through a tower filled
with potassium hydroxide pellets.
Carbon monoxide is very poisonous : all operations
involving its preparation and use must be carried out
in an efficient fume cupboard.
Here is the procedure from Preparative Inorganic Chemistry.
Cr(CO),, Mo(CO),, W(CO),
The hexacarbonyls of the chromium group are formed via
reaction of CO with a suspension of anhydrous halides of Cr, Mo
or W in a Grignard solution, followed by hydrolysis. The reaction
mechanism has not yet been elucidated.
The reactor vessel / in Fig. 343 is a one-liter flask fitted
with a two-hole rubber stopper. The dropping funnel t has a
considerably enlarged tip to prevent plugging during the reaction.
It is used for the addition of the Grignard solution (via a), as well
as that of CO (at b). Stopcock h is a gas vent which remains
normally closed during the reaction but which is occasionally
opened to allow flushing the reactor with CO. Flask f is fitted
exactly into the ice bath e, and the whole apparatus is vigorously
shaken on a machine. To monitor the CO consumption, a standardized
gasometer is connected to b via a drying train (whose
last tube is filled with P3O5).The reactor flask / is filled with nitrogen. The metal chloride
(10 g. of fine anhydrous CrCl3 powder; 17 g. of sublimed MoCls;
or 20 g. of WC16 [0.05 moles]) is introduced, and the vessel is
evacuated and filled with CO. A mixture of 50 ml. of anhydrous
ether and 50 ml. of anhydrous benzene is added through the
dropping funnel and the apparatus is then connected to the CO line.
The Grignard reagent is prepared from 12 g. (0.5 moles) of
Mg, 54 g. of C3HsBr and approximately 300ml. of anhydrous ether.
This solution is added to the metal chloride suspension first in
portions of about 5 ml. each, later dropwise. The initiation of the
CO reaction as well as its progress may be observed via a wash
bottle containing some cone. H3SO4 provided the stopcock of t is
closed. The absorption of CO, which for reasons unknown occasionally
slows down and then accelerates, is continued for about
4-6 hours after the addition of all of the Grignard reagent. The
reaction absorbs on the average 7 liters and occasionally up to
9 liters of CO.The reddish-brown reaction product is hydrolyzed by cautious
addition to a mixture of ice and dilute HsSO^ and the mixture is
then steam-distilled without prior removal of ether and benzene.
The steam distillation is continued for 3-4 hours or as long as
white needles of the carbonyl product are observed in the (descending)
condenser. The organic layer (benzene-ether) in the distillate
is separated and the aqueous phase extracted 3-4 times with fresh
ether. The combined ether extracts are concentrated by distillation,
keeping the temperature below 60 °C, and the residue is allowed
to crystallize in a refrigerator.
The yields of crude carbonyls are quite variable: in the case of
Cr(CO)6 they are 2 g. maximum, while up to 3-4 g. of W(CO)6
may be isolated. Higher yields of Cr(CO)s (up to 67%) are obtained
in an autoclave under high CO pressure (35-70 atm.). To remove
strongly adhering, odorous organic impurities, an immediate
vacuum sublimation of the hexacarbonyls is recommended.Jor - 29-3-2009 at 03:14
Do you do this in a fume hood at home?
As these metal carbonyls are really toxic, at least nickel carbonyl is.
You can also prepare CrCl3 by passing CCl4-vapour over Cr2O3, at very high temperatures in a tube furnace. You will have COCl2 byproducts, wich you
can bubble in toluene for later use.
Im not sure if pottery Cr2O3 will work. It is insufficient in many reactions, I heard, because of inertness. Is this true?
You can prepare very active Cr2O3 by decomposing ammonium dichromate. This way I made some Cr2O3 wich i used as a catalyst (heating it to red heat) to
oxidise ammonia with oxygen.
[Edited on 29-3-2009 by Jor]JohnWW - 29-3-2009 at 06:52
In the formation of Cr(CO)6, Mo(CO)6, and W(CO)6, electron pairs from the C in the CO molecules would occupy the remaining three vacant 3d or 4d or 5d
orbitals and all three vacant 4p or 4p or 6p orbitals, of the zerovalent metal atoms. Because all 6 CO molecules would be equivalent by resonance,
this looks like a case of p-d orbital hybridization. However, they would not be as stable as Ni(CO)4 or Pd(CO)4 or Pt(CO)4, because the 4p or 5p or 6p
orbitals are at relatively higher energy levels than in Ni, Pd, and Pt. In the Ni, Pd, and Pt tetracarbonyls, the d valence orbitals are all full, and
the electron pairs from the COs occupy the vacant s and p orbitals (sp3 hybridization).benzylchloride1 - 29-3-2009 at 20:51
Nickel carbonyl is one of the most toxic substances known. Most other metal carbonyls are toxic, but not to the degree of nickel carbonyl. I have
conducted some work at the university with iron pentacarbonyl which is probably the most toxic common metal carbonyl. This compound is a low boiling
liquid with a bad musty odor. The fume hood i have at home is very similar to the ones at the university. Before I work with Carbon monoxide, I will
install several detectors around the fume hood. The hexacarbonyls of Cr, Mo, and W are solids with high vapor pressures and can be handled safely in a
fume hood. I wonder if the corrosponding anhydrous bromides could be used in the reaction? These could be made by adding bromine to the metal of
choice in a reaction flask. The reaction would be initiated by adding a drop of water to form a small amount of the aqua complex. I have also found a
stable copper (I) carbonyl that could be prepared fairly easily. Copper (I) iodide is treated with sodium tris(3,5- dimethylpyrazolyl)hydroborate in
an ether solution. Carbon monoxide is bubbled through the solution until crystals of the complex form. The sodium
tris(3,5-dimethylpyrazolyl)hydroborate is prepared by heating sodium borohydride with 3,5-dimethylpyrazole until the evolution of hydrogen ceases. The
chemistry of metal carbonyl complexes is fascinating. It would be nice if there was a way of producing iron pentacarbonyl in useable quantities, 10g
without an autoclave and cylinders of CO.
