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Author: Subject: Rocket Fuel from UV Photolysis of Alcohol/Aqueous Nitrate?
AJKOER
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[*] posted on 19-9-2017 at 09:54
Rocket Fuel from UV Photolysis of Alcohol/Aqueous Nitrate?


Well, may be not so much as rocket fuel as a fuel spiker for testing on the old lawn mower!

Lets consider the hypothetical UV photolysis of a methanol (or ethanol) and water solution containing nitrate. First, I could expect some radical formation, which is reported to occur in the UV (defined bands, see https://en.m.wikipedia.org/wiki/Ultraviolet) photolysis of aqueous nitrates:

NO3- (aq) + hv → •OH + •NO2

Source: "Mechanism of Nitrite Formation by Nitrate Photolysis in Aqueous Solutions:  The Role of Peroxynitrite, Nitrogen Dioxide, and Hydroxyl Radical", by Sara Goldstein and Joseph Rabani, published in Journal of the American Chemical Society, Vol. 129: , Issue. 34, 2007, Pages. 10597-10601. To quote from the first page :

"Photolysis of aqueous NO3- with λ ≥ 195 nm is known to induce the formation of NO2- and O2 as the only stable products. The mechanism of NO3- photolysis, however, is complex, and there is still uncertainty about the primary photoprocesses and subsequent reactions. This is, in part, due to photoisomerization of NO3- to ONOO- at λ < 280 nm, followed by the formation of •OH and •NO2 through the decomposition of ONOOH (pKa = 6.5−6.8)."

Link: http://pubs.acs.org/doi/abs/10.1021/ja073609+

Further reactions in say a water methanol solution containing nitrate under UV photolysis could include:

CH3OH + •OH → •CH3 + H2O

NO3- + •OH → •NO3 + -OH

•CH3 + •NO2 → CH3NO2. (Caution, per Wikipedia on Methyl nitrite, it is a toxic asphyxiating gas and a potent cyanotic agent with exposure possibly resulting in methemoglobinemia, see https://en.m.wikipedia.org/wiki/Methyl_nitrite )

•CH3 + •NO3 → CH3NO3 (Caution, per Wikipedia on Methyl nitrate, it is a toxic, volatile, and a sensitive HE, see https://en.m.wikipedia.org/wiki/Methyl_nitrate )

•OH + •NO2 → HONO2

•OH + •NO3 → HONO3
......

With the possible indicated formation, in small amounts (maybe a good thing here) of some interesting energetic/toxic compounds.

Miscellaneous comments: CH3OH + CH3NO3 has been used as a rocket fuel (Wikipedia reference above), pure methyl nitrate is listed as a high explosive (see demo at https://www.google.com/url?sa=t&source=web&rct=j&...). Also, the photolysis of any produced nitrite is said to be superior to aqueous nitrate, which could increase the rate of product formation. I would not advise heating the product brew as CH3NO2 is a heat sensitive explosive which could likely detonate any formed CH3NO3. Final comment, if you dismiss the feasibility of my chemistry above, please note this comment on Wikipedia on incidental methyl nitrite creation:

"Methyl nitrite is also present in aged cigarette smoke. Here it is presumably formed from nitrous oxide (itself formed by autoxidation of nitric oxide) and methanol.[5]"

which I would alternately explain as the result of thermally induced methyl radical creation reacting in the presence of NxOy.

[Edited on 19-9-2017 by AJKOER]
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[*] posted on 19-9-2017 at 11:09


How high does the alkyl nitrate concentration need to be before the rate of photolysis is as high as the rate of production?
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[*] posted on 19-9-2017 at 11:48


Quote: Originally posted by unionised  
How high does the alkyl nitrate concentration need to be before the rate of photolysis is as high as the rate of production?


I would not claim to be much of an authority in this area. However, my general comment is that the rate of photolysis is generally low, and per my recollection depending on the reactants, may be further even subject to a threahold limitation (meaning, after a concentration point, the reaction rate may be independent of further increases/decreases in reactant concentration).

As here the rate of product production is due to fast reactions with short lived radicals, I would not expect any of these radical reactions to constitute a rate limiting step.

Here is a reference, for example https://www.google.com/url?sa=t&source=web&rct=j&...

Another source comments:

"For NO2 and O3, photodissociation quantum yields do not drop immediately to zero below the dissociation threshold. This effect is explained by channeling the internal energy of molecules into the dissociation process."

Link: https://www.google.com/url?sa=t&source=web&rct=j&...


[Edited on 19-9-2017 by AJKOER]
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[*] posted on 19-9-2017 at 11:58


"However, my general comment is that the rate of photolysis is generally low".
That's the problem.
You only photolyse nitrate slowly, so you only make alkyl nitrate slowly.
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[*] posted on 19-9-2017 at 12:29


Quote: Originally posted by unionised  
"However, my general comment is that the rate of photolysis is generally low".
That's the problem.
You only photolyse nitrate slowly, so you only make alkyl nitrate slowly.


True, I did mention 'small quantities'.

Perhaps a list of favorable/unfavorables would help. My comments would be:

) Process is slow but yield increases with time.

) Total expected products, assuming using something other than a weather balloon, would still be small.

) No demand for strong acids or expensive reagents.

) Unconventional equipment required for the photolysis (UV lamp).

) Limited need for standard lab equipment.

I think whether the unfavorables out weigh the favorables may depend on personal preferences and restrictions. I cannot operate a legal lab and reagents/glassware purchases raise red flags. I even get questioned buying over the counter stuff (like, what are you doing?, what do you need it for?, ...)
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[*] posted on 19-9-2017 at 12:53


" Process is slow but yield increases with time."
Only until you start to destroy it as fast as it's made.
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[*] posted on 19-9-2017 at 13:24


Too slow for any practical use... also at a certain point CH3-ONO2 and CH3-ONO or CH3-NO2 also enter the photolysis process and are destroyed...

To speed things up... and circumvent those problems...
CH3-Cl (or even better CH3-I) + AgNO3 or AgNO2 would be more suitable...
CH3-I is a good source of CH3° and I° while AgNO3 and AgNO2 are good sources of Ag° and NO3° or NO2°...
Excluded the photochemical reaction; the reaction also occure into solution...because of the driving force of precipitation...
AgNO3 + CH3-Cl --> AgCl(s) + CH3-ONO2
AgNO2 + CH3-I --> AgI(s) + CH3-NO2 (66%) + CH3-ONO (34%)

The speed of that reaction is nearly instantaneous with primary alkyl halides... it is a known reaction to determine qualitatively the order of an alcohol function
==> primary, secondary or tertiary...
This by passing via its halide and reaction with aqueous AgNO3 solution ... immediate precipitate/trouble is a sign of primary... delayed means secondary and none means tertiary...

Another fuel than CH3OH would be aceton (propanone)... propanone has a very good quantum yield...
It is a way to obtain CH4 and CO from propanon

[Edited on 19-9-2017 by PHILOU Zrealone]




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[*] posted on 20-9-2017 at 14:52


Quote: Originally posted by unionised  
" Process is slow but yield increases with time."
Only until you start to destroy it as fast as it's made.


Different compounds have differing quantum yields. There is also an unmentioned (to be determined based on a yield time study for products of interest boosting fuel performance) reaction time frame. This is akin to when a farmer decides to harvest his crop.

I do find it curious that the some people who speak highly of electrolysis (which is inherently slow by nature also) point singularly to time issues with photolysis. Also, there are instances in electrolysis where the electrode's surface area (or lack thereof) is an important issue (for example, when the recommended electrode is a platinum mesh) . If the corresponding factor in photolysis is the vessel size or light source access, this is usually more readily addressable, in my opinion, for photolysis.
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A correction: CH3OH + •OH → •CH3O + H2O

But then possibly: •CH3O + •NO2 → CH3ONO2 (methyl nitrate)

The above suggests to me possibly replacing nitrate with nitrite (which is also more photo active in creating •OH) given the reaction between the hydroxyl radical and the nitrite ion:

NO2- + •OH → •NO2 + OH-

given the above interaction with •CH3O. Also:

CH3OH + hv → •CH3 + •OH (for other possible branch products, see http://chemistry.emory.edu/faculty/widicusweaver/photolysis.... )

The above is one possible basis for the photodegradation of compounds in methanol in sunlight. For example: "Photolysis of fluchloralin in aqueous methanol", by Tapas Saha and Anjan Bhattacharyya, 2001, DOI: 10.1002/ps.425. To quote the abstract:

"The photodegradation of fluchloralin by UV irradiation or sunlight in aqueous methanolic solution has been examined. In the presence of titanium dioxide five photoproducts were obtained, but only four in its absence. One photoproduct, 2, 2′-azoxy-bis(α,α,α-trifluoro-6-nitro-p-toluidine) is reported for the first time as a metabolite of fluchloralin. In natural sunlight the rate of degradation was higher than in UV light and titanium dioxide had almost no effect on the rate of degradation."

Link: http://onlinelibrary.wiley.com/doi/10.1002/ps.425/abstract

Also, borosilicate glass is inferior to quartz in UV photolysis with methanol, at least in one case, see http://www.sciencedirect.com/science/article/pii/00456535859...

[Edited on 21-9-2017 by AJKOER]
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[*] posted on 21-9-2017 at 04:13


Came across an available online photolysis study ("Solvent Effect on the Photolysis of Riboflavin" by Iqbal Ahmadin, et al, in AAPS PharmSciTech, 2015 Oct, 16(5): 1122–1128) involving aqueous methanol with some interesting points.

For example, Table I displays the molar concentrations x E05 of RF and the corresponding photo products in methanol over a 2 hours period where RF denotes riboflavin, FMF is formylmethylflavin and LC is lumichrome. The 120 minutes reported values (which I have rounded to one decimal place) are RF 1.9 (from a starting value of 3.0), FMF .8 (from zero) and LC .4 (from zero).

An extract that presents an interesting summary of results:

"A plot of k obs for the photolysis of RF as a function of solvent dielectric constant is presented in Fig. 3. It shows that the rate constants are linearly dependent upon the solvent dielectric constant. Similarly, a linear relation has been found between the values of k obs and the solvent acceptor number indicating the degree of solute–solvent interaction (Fig. 4). In order to observe the effect of viscosity on the rate of photolysis, a plot of k obs versus inverse of viscosity was constructed (Fig. 5). It showed a linear relation between the two values indicating the influence of solvent viscosity on the rate of reaction. These results are supported by the fact that a plot of dielectric constant versus inverse of viscosity of organic solvents is linear."

Link: https://www.ncbi.nlm.nih.gov/pmc/articles/PMC4674638/

While the starting concentration of RF was small, the percentage degradation (or the percent of new products created) in just 2 hours (nearly 37%), I found impressive.
----------------------------------------------------------------

Related work concerning the dielectric properties of alcohol-water mixtures, please see "Effective permittivity of alcohol + water mixtures as influenced by concentration", available at http://www.jocpr.com/articles/effective-permittivity-of-alco... .

My takeaway on this paper is that increasing alcohol concentration in water lowers the dielectric constant of the mixture and, per the above study, likely reduces the rate of photolysis depending on the offset due to any reduction in viscosity (which is impacted by salt concentration).

[Edited on 21-9-2017 by AJKOER]
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[*] posted on 21-9-2017 at 05:33


Interesting.
But such large colorizer molecules display high absorbances into various wavelenght ranges... what is not the case of methanol.

Would be nice to have a colorizer that is a kind of catalyst of your proposed reaction... but that remains integer...
So the colorizer catches the light energy at a specific wavelenght and then transfers its "radical-ization" energy to the solvent for reaction...
This is often used into photochemistry

About your initial post...
You have to take into account the various reaction of radicals (initiation, propagation and termination)

For example:
HO°+ °OH --> HO-OH
CH3° + °CH3 --> CH3-CH3
CH3O° + °CH3 --> CH3-O-CH3
HO° + NO3° --> HO-O-NO2 --> HOO° + NO2°
CH3O° --> CH2=O + H°
HO° + H° --> H2O

Your choice of molecules (Methanol and HNO3/ NO3(-)) is not the best since they allow for a lot of possibilities (read side reactions)




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[*] posted on 21-9-2017 at 06:35


PH Z:

Yes, colorizers are an option, but as mentioned above in one of my sources experimenting with CH3OH/TiO2, I was thinking of ZnO in sunlight to activate N2O (actually mentioned in an online ebook), with the laughing gas displaying high solubility in alcohol. The mechanics I would express as:

ZnO + hv → [ZnO]* → ZnO + e-

e- + N2O → N2 + •O

•O + H2O → OH- + •OH (except at very high pH, then just •O)

This reaction also reduces water content, a good thing from a combustion viewpoint. Also, apparently ZnO, which is not quite as a good photocatalyst as TiO2, can be inexpensively prepared, I suspect, by pouring NaOCl (bleach) into a container containing zinc metal, a little sea salt and a piece of copper. Jump start the electrochemical reaction (actually a battery cell) in a microwave oven. Scrap off any ZnO not settled to the bottom.

Note, I also edited a comment above leaning to nitrite in lieu of nitrate.
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With respect to your reaction chain comment, perhaps not such a big issue (I have already revealed this with one of my reference link above, which provides a graphic depiction of a long chain with the photolysis of CH3OH, repeated here http://chemistry.emory.edu/faculty/widicusweaver/photolysis.... ). My reasoning is how many of possible organic nitrates or nitrites created are not likely to be energetic? The problem remaining products for fuel boost purposes are likely water and H2CO (also a health issue).

Warning: If one elects to infused fuel with N2O, be aware that it can be detonated with a high grade primer or possibly with an electrostatic discharge. Not yet clear if such a scenario is even possible here, but is worth noting as to why this thread is placed in the Energetic Materials forum and experimenting is recommended only in small quantities.

[Edited on 21-9-2017 by AJKOER]
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[*] posted on 10-10-2017 at 10:57


Quote:
My reasoning is how many of possible organic nitrates or nitrites created are not likely to be energetic?

First of all, alkyl nitrites are pretty unstable, and can oxidize other things or cause their decomposition quite easily. Also, methyl nitrite is a gas.

I'm not sure why you think that EVERY reaction has to be photocatalyzed. Free radicals aren't usually a good thing to have in a reaction, since they're often not selective about how they react. I beileve bromine is probably the main, if not only exception.




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[*] posted on 8-12-2017 at 17:08


Melgar:

You may be right on the lack of need for a photo catalysis here as the burning of even a standard fuel can produce high temperatures, and said temperatures, used to pre-heat a fuel mix, provide a natural path for thermally produced radicals.

Interestingly, I came across the following statement in Wikipedia (https://en.wikipedia.org/wiki/Methanol), to quote:

"As a fuel for mud racers, methanol mixed with gasoline and nitrous oxide produces more power than gasoline and nitrous oxide alone."

This statement may possibly be indicative of thermally based radical formation of a higher energetic compound, from the less energetic methanol fuel interacting with N2O, since the thermally treated mix of CH3OH/N2O/gasoline has unexpectedly more power per unit volume than the pure gasoline/N2O pre-heated mix.

Or, the pre-heating of methanol, N2O and possibly water as a compound capable of absorbing the heat energy, or other endothermic reactions (including generally, for example, reagents reacting to form explosives), and said compound(s) undergo a phase change or subsequently decomposed on ignition, thereby releasing, as added energy, the previously absorbed thermal energy.

Or, the high temperature methanol including its impurities (like water, formic acid,..) enables a corrosion reaction with the stainless or carbon steel engine (see, for example, http://emsh-ngtech.com/methanol/our-methanol/ ). The consequence of this process in the presence of N2O could proceed as follows:

M(n) = M(n+1) + e-

N2O + e- = N2 + .O- (See Reaction [45] at https://www.google.com/url?sa=t&source=web&rct=j&...)

.O- + H2O = .OH + OH-

H+ + OH- = H2O

Example of a possible net reaction with iron in the presence of a corrosive methanol mix (composition, see https://www.google.com/url?sa=t&source=web&rct=j&... ) assuming a single electron reduction of N2O occurring in acdic conditions:

Fe2+ + N2O + H+ (aq) --> Fe3+ + N2 + .OH

Related sources as author assumes a two electron reduction of N2O occurring in neutral to basic conditions: https://www.ctahr.hawaii.edu/huen/tpss435/redox.pdf and http://wwwuser.gwdg.de/~kuzyakov/Geoderma_2016_Wang-Milan_Fe... .

The radical creation could then induce a radical chain reaction with the gasoline. Bottom line, any limited power performance increase in the presence of a N2O/methanol mix may be actually instigated by the corrosive nature of the methanol mix with respect to certain metals.

If others have alternate explanations, please expound.

Thanks to all for the quality of the review to date.

[Edited on 10-12-2017 by AJKOER]
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