Sciencemadness Discussion Board - Producing Hydrogen by partial oxidation

Producing Hydrogen by partial oxidation

AndersHoveland - 3-12-2011 at 13:46

I am not sure I believe it!
Hydrogen gas can be liberated by partial oxidation of formaldehyde using either H2O2 or CuO under certain reaction conditions.

Quote:

Formaldehyde, Action of Hydrogen Peroxide on. A Harden. Proc. Chem. Soc. 15, [212], 158-159

When solutions of hydrogen peroxide and formaldehyde are mixed, no reaction appears to take place, but when the liquid is made strongly alkaline with soda, hydrogen is evolved. The reaction occurs according to the equation
H2O2 + 2 CH2O + 2 NaOH = 2 HCO2Na + H2 + 2 H2O
Hydrogen peroxide, therefore, when treated with alkaline formaldehyde, gives a volume of hydrogen exactly equal to the volume of oxygen which it would give with potassium permanganate and sulphuric acid. The reaction proceeds slowly and incompletely unless a large excess of alkali be present. When hydrogen peroxide is treated with excess of formaldehyde, the reaction takes place rapidly and completely, and the hydrogen which is evolved is pure. When, on the other hand, formaldehyde is treated with excess of hydrogen peroxide, the reaction is incomplete and proceeds very slowly, whilst the gas evolved contains oxygen.

Cuprous oxide and soda give a somewhat similar reaction with formaldehyde. This reaction was described by Loew (Ber. 1887, 20, 145) as a catalytic reaction, but it appears in reality to be a quantitative one, expressed by the equation
Cu2O + 2 NaOH + 2 CH2O = Cu2 + H2 + 2 HCO2Na + H2O
Cupric oxide also gives a similar reaction, two atoms of hydrogen being liberated for each atom of oxygen in the oxide

When caustic soda and then formaldehyde are added to a solution of copper sulphate and the liquid gently warmed, the cupric hydroxide is reduced to cuprous oxide without evolution of hydrogen, and when the temperature is subsequently raised, the cuprous oxide reacts as described above. When, on the other hand, caustic soda is added to a boiling solution of copper sulphate, the liquid cooled, and formaldehyde then added, no reduction of cuprous oxide occurs on warming, but metallic copper is formed, and twice as much hydrogen is evolced as the previous case.
Manganese dioxide does not appear to be reduced by formaldehyde, whilst the oxides of mercury and bismuth are reduced without evolution of hydrogen.

Journal of the Society of Chemical Industry, Volume 18, p716-717. [July 31, 1839]

The fact that hydrogen and oxygen gas can be liberated together from the same solution simultaneously may also be of interest, for making gas explosions inside of closed plastic bottles.

AndersHoveland - 4-12-2011 at 01:54

Hydrogen is more electronegetive than carbon. Does anyone have any idea why the carbon gets preferentially oxidized (to formate) rather than the hydrogen?

I would imagine the reaction might proceed through the intermediate H2C(OH)O[-]. But how does molecular hydrogen get liberated ?

kmno4 - 5-12-2011 at 13:57

Generating H2 in this and similar reactions is interesting in itself.
Proposed machanism can be found, for example, here: Russ. Chem. Rev. 29 193

(I pretend that I do not see this insane "making gas explosions inside of closed plastic bottles" by insane user)

[Edited on 5-12-2011 by kmno4]

AndersHoveland - 5-12-2011 at 17:20

Quote:

the formaldehyde reduces [cupric oxide] not only to cuprous oxide, but to metallic copper [in colloid form], but as this happens a second reaction ensues in which the metallic copper acts upon the formaldehyde and decomposes it with the liberation of hydrogen.

Potassium cyanide, when added [to the above reaction] will inhibit or stop further reduction and liberation of hydrogen.
[text describes potassium cyanide as an example of an 'enzymatic poison']

Science, Volume 45, American Association for the Advancement of Science, p507

Quote:

Loew recorded the highly significant observation that the interaction of aqueous solutions of formaldehyde and sodium hydroxide yielded a very small amount of hydrogen, and that if cuprous oxide were added to the reaction mixture, it was reduced to metallic copper.

Quote:

As expected, formaldehyde is directly oxidized [by H2O2] to formic acid, and the latter in turn to carbonic acid.

But hydrogen peroxide may, on the one hand, oxidize formaldehyde to formic acid with the liberation of hydrogen.
2HCHO + H2O2 --> 2HCOOH + H2

while on the other hand it will reduce carbonic acid to formic acid with the liberation of oxygen
H2CO3 + H2O2 --> HCOOH + H2O2 + O2

Since hydrogen peroxide may thus react either as an oxidizing agent or as a reducing agent, and in some instances in both capacities concurrently…

“The Action of Hydrogen Peroxide upon Simple Carbon Compounds
H. Shipley Fry and John H. Payne, Journal of the American Chemical Society (1931)

for more details about this reaction see Journal of the American Chemical Society, Volume 29, Issue 2, p1233,
http://books.google.com/books?id=5FQ2AQAAIAAJ&pg=PA1233&...

which explains that the hydrogen is formed only as the formaldehyde is being oxidized to formic acid
2HCHO + H2O2 --> 2HCOOH + H2
but not when the formic acid is subsequently oxidized to carbon dioxide. And some of the formaldehyde is oxidized to formic acid without any matching formation of hydrogen. About twice as much carbon dioxide is evolved as hydrogen gas.

The reaction between formic acid and hydrogen peroxide must be more complicated. Simply mixing the two chemicals apparently results in "performic acid", HOOCH=O, in equilibrium. The oxidation of formic acid to carbon dioxide apparently requires alkaline conditions.

Other than the mention in the above source, I cannot find anything else about hydrogen peroxide being able to reduce carbonic acid to formic acid, and such a reaction is very doubtful. I assume that the author made a mistake. I also found mention in another source that "Reduction of carbon dioxide by hydrogen peroxide... in the production of oxygen and formic acid from carbonic acid and hydrogen peroxide", but this was in the context of hypothesized uptake of carbon dioxide by plants in the year 1919, before the complete modern understanding of photosynthesis, so I would be very hesitant to accept it as fact.
Chemical abstracts, Volume 13, Issue 2, American Chemical Society. p1483 subarticle by Carl L. Alsberg.
http://books.google.com/books?id=8cW2AAAAIAAJ&pg=PA1483&...

Other Reactions involving liberation of Hydrogen from Organic Compounds
Wurtz found that ethylene glycol fused with solid potassium hydroxide yielded mostly potassium oxalate and hydrogen gas. The yield of hydrogen was 58% of that required by the below proposed reaction
C2H6O2 + 2KOH --> K2C2O4 + 4H2

Rosorcinol (1,3-dihydroxy-benzene) is similarly converted to phloroglucinol (1,3,5-trihydroxy-bezene) by fusion with sodium hydroxide, with the evolution of hydrogen.

These are both unusual examples of disproportionation reactions, where a compound is simultaneously oxidized and reduced.

source:
"THE LIBERATION OF HYDROGEN FROM CARBON COMPOUNDS", Shipley Fry, Else L. Schulze, Helen Weitkamp
J. Am. Chem. Soc., 1924, 46 (10), pp 2268–2275

[Edited on 6-12-2011 by AndersHoveland]

AndersHoveland - 6-12-2011 at 16:04

I think the reaction mechanism might involve the formation of an
HO-O-O-CH=O intermediate, and cylization of this intermediate to an unstable square ring, which could then decompose into O2, CO2, and H2. See the picture below:

The driving force for the typically unfavorable formation of H2 would be the simultaneous formation of O2 and CO2, which are highly favorable.

The transient square ring would contain 3 atoms of oxygen and 1 atom of carbon, with an additional fourth oxygen, with a negetive charge on it, bonded to the carbon. A single atom of hydrogen would also be bonded to the carbon. The hydrogen atom on this ring would encounter a positively charged hydrogen ion from outside, which would pull out the negetive charge on the oxygen atom through the molecule. I am mostly writing this paragraph in the event that the accompanying picture stops working at some time.

If this is in fact correct, it suggests that the formation of hydrogen may actually be due to the oxidation of formic acid to carbon dioxide, rather than the direct oxidation of formaldehyde to formic acid.

It is known, for example, that dihydrogen trioxide, H2O3, has a transient existence in alkaline solution of hydrogen peroxide, and this could potentially be the reason that alkaline H2O2 acts as a stronger oxidizing agent.
H2O2 <==> HOO[-] + H[+]
H2O2 + HOO[-] <==> H2O3 + OH[-]

I really do not know if this is the correct reaction mechanism, or whether there is some other more predominant mechanism, but for now it seems like the best explanation for the liberation of hydrogen gas.

Here is the likely mechanism for the normal oxidation of formaldehyde to formic acid by H2O2,

[Edited on 7-12-2011 by AndersHoveland]

AndersHoveland - 16-12-2011 at 00:15

Doing a search of this forum, I found this:

Quote: Originally posted by KABOOOM(pyrojustforfun)

HCHO + H2O    qe   H2C(OH)2
Formaldehyde peroxidation
H2O2 + H2C(OH)2   qe   HOCH2OOH + H2O
HOCH2OOH + H2C(OH)2   qe   HOCH2OOCH&l t;SUB>2OH + H2O

One approach was to use another difunctional peroxide that those already in use (hydrogen and adipoyl peroxides were already patented). I looked at some texts and noted one could make an aldol peroxide from formaldehyde and hydrogen peroxide. I put together some concentration variations of butadiene, isopropanol, formaldehyde, and hydrogen peroxide, to make the initiator in- situ. We set them up in the standard sealed polymerization bottles at 75 degrees Centigrade. Nothing happened for about two hours, and I looked in one them hourly. Just after the 2nd hour inspection, I went back to my desk for coffee, and one of the technicians raced up saying "Did you hear the explosion?". I hadn't, but went with him to the polymerization lab external door just in time to hear another one "cook off".
The lab was cleared of all personnel already, and one just barely missed being inundated by the 75 degree water as the first explosion went off. I put on face mask and rubber apron and crawled into the lab to turn off the steam and fill the bath with cold water. The lab floor was littered with sand sized glass particles, and one of the polymerization bottle guards was about 15 feet from the polymerization bath. After about a half hour of cooling, I donned apron, long gloves and face mask, and took out all the remaining bottles, and emptied those that still were intact- not many. The force of the explosion ripped the stainless caps off some of the bottle shields, and one of those shields remains on my desk to remind me I'm fallible.
This was the last day before my vacation was to start, so when I left after cleaning up, I was gone for at least a week. When I got back, I found that there had been an inquiry. Gerd Lenke had been assigned top look into the matter, and he found an ancient Gmelian book set which showed details of the formaldehyde/hydrogen peroxide products. The one I wanted (and probably got) was dimethylol peroxide, which melts about 60 degrees C, and is a "brisant" explosive at 72 degrees. It would not have detonated at 70 degrees, but I was five degrees higher. I still think it would be a good di-functional initiator, BUT...Handle With Care!

There where some links in the post, but unfortunately they stopped working.

Essentially the reaction between H2O2 and CH2O apparently can potentially be dangerous.

This paper is revealing:
http://pubs.acs.org/doi/abs/10.1021/ac60297a031

Another interesting, but more distantly related, reaction involves the reaction between methyl alcohol and air, catalysed by the enzyme methanol oxidase, to produce formaldehyde and hydrogen peroxide. The reaction should preferrably be done at [minus] -22 degrees Celsius, because otherwise the enzyme is only stable for 2 days at ambient temperatures. The enzyme can be produced by the yeast Hansenula polymorpha DL-1.
http://www.che.cemr.wvu.edu/publications/projects/large_proj...

[Edited on 16-12-2011 by AndersHoveland]

ScienceSquirrel - 16-12-2011 at 05:25

I remember mixing hydrogen peroxide and formaldehyde solution together a long time ago.
The mix would remain completely quiescent for some time and then a fast evolution of gas would shoot most of the liquid out of the tube. I was only using 3% hydrogen peroxide. I suspect that a sealed bottle with 35% hydrogen peroxide would be an unpredictable bomb!

AndersHoveland - 16-12-2011 at 11:26

Do you think the hydrogen and oxygen simultaneously evolved from the reaction in somehow spontaneously igniting? Could this a potential cause of the reported explosions?
Or could the gas explosion above the solution be the cause of detonation of dimethylol peroxide, or some other related formaldehyde-peroxide explosive, in the solution?

Adas - 16-12-2011 at 11:43

Quote: Originally posted by ScienceSquirrel

I remember mixing hydrogen peroxide and formaldehyde solution together a long time ago.
The mix would remain completely quiescent for some time and then a fast evolution of gas would shoot most of the liquid out of the tube. I was only using 3% hydrogen peroxide. I suspect that a sealed bottle with 35% hydrogen peroxide would be an unpredictable bomb!

What the hell? 3% peroxide? This is simply unbelievable!!! :O

AndersHoveland - 14-1-2013 at 02:59

The reaction may be much simpler. Under alkaline conditions, the formaldehyde may convert to
HO-CH₂-O^-

In the presence of something that can act as a mild oxidizer (such as CuO or H2O2), a hydrogen atom on the carbon atom would be very vulnerable. Only the interesting thing is that the hydrogen atom is merely freed from the carbon, since it is the extra electron in the molecule that is being taken. And of course this would leave behind CO₂

Formaldehyde in the presence of a base could make a selective reducing agent for other chemical reactions.

Perhaps Woelen will do some experiments with this...

[Edited on 14-1-2013 by AndersHoveland]

ScienceSquirrel - 14-1-2013 at 05:26

Formaldehyde will act as a reducing agent in a variety of situations eg the crossed Cannizzaro;

http://www.adichemistry.com/organic/namedreactions/cannizzar...

and the synthesis of methylamine from formaldehyde and ammonia;

http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv1...

Adas - 16-1-2013 at 10:03

This idea just came to my mind and I would like to share it, so someone could test if it works. It is a simple experiment.

2 Cu+ + 2 NaHSO4 ----------> 2 CuSO4 + 2 Na+ + H2

All you need is Copper(I) salt and sodium hydrogen sulfate. I know that it's impractical but I just want to know if it can work. It may work, because Cu(HSO4

2 can not form so Copper(I) must oxidize to Copper(II).

EDIT: I am sorry, I didn't check the reaction, Copper can not push Sodium out of the sulfate salt. Maybe it would work with combination of other salts, and maybe other salts + acid.

[Edited on 16-1-2013 by Adas]

ScienceSquirrel - 16-1-2013 at 13:13

Quote: Originally posted by Adas

This idea just came to my mind and I would like to share it, so someone could test if it works. It is a simple experiment.

2 Cu+ + 2 NaHSO4 ----------> 2 CuSO4 + 2 Na+ + H2

All you need is Copper(I) salt and sodium hydrogen sulfate. I know that it's impractical but I just want to know if it can work. It may work, because Cu(HSO4

Metals like magnesium, aluminium and zinc will react with sodium hydrogen sulphate to form the corresponding ions and hydrogen.
Copper metal and copper I salts will react with acids in the presence of oxidising agents, sometimes this can be the acid itself, it can be as simple as atmospheric oxygen or it can be an added oxidising agent eg hydrogen peroxide. Hydrogen gas is never formed to my knowledge.

Morgan - 2-3-2013 at 19:02

Quote: Originally posted by AndersHoveland

Hydrogen is more electronegetive than carbon. Does anyone have any idea why the carbon gets preferentially oxidized (to formate) rather than the hydrogen?

I would imagine the reaction might proceed through the intermediate H2C(OH)O[-]. But how does molecular hydrogen get liberated ?

The other day I was reading an article about iron and decided to memorize the site's electronegative series. I keep finding that carbon is more electronegative than hydrogen not the other way around. Am I missing something?

AJKOER - 2-3-2013 at 22:12

Quote:

Quote: Originally posted by AndersHoveland

Science, Volume 45, American Association for the Advancement of Science, p507
......

Loew recorded the highly significant observation that the interaction of aqueous solutions of formaldehyde and sodium hydroxide yielded a very small amount of hydrogen, and that if cuprous oxide were added to the reaction mixture, it was reduced to metallic copper.

As an interesting sidebar, I find the formation of copper in colloid form and if the cupric oxide is also in colloid form, to be possibly significant. At the time of the writings, the use of Copper oxides in catalyst, and preparation route thereto, may not have been completely appreciated.

Am I right that colloid cupric oxide could be a powerful catalyst and have interesting properties? See for example "Smart copper oxide nanocrystals: Synthesis, characterization, electrochemical and potent antibacterial activity at http://www.sciencedirect.com/science/article/pii/S0927776512... and note the incomplete statement "Among the inorganic nanocrystals, Copper oxides are industrially very important which are widely used in fields such.."

The only other path for home chemists that I know for colloidal CuO (and not Copper) would be (link: http://onlinelibrary.wiley.com/doi/10.1111/j.1551-2916.2006.... , to quote: "colloidal suspensions of copper oxide (CuO) nanoparticles were prepared by an alcothermal method, in which copper acetate was reacted with sodium hydroxide in the presence of acetic acid in ethanol at 78°C."

[Edited on 3-3-2013 by AJKOER]

AndersHoveland - 3-3-2013 at 01:18

I do not think Cu⁺¹ is even stable in aqueous solution. It disproportionates to Cu⁺² and copper metal.

Cu+1 salts can form, however, in the form of coordination complexes, such as CuCl₂^-, which is stable in aqeuous solution. This is often just called (not entirely accurately) copper(I) chloride. Copper cyanide is another.