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BryceFitz
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[*] posted on 23-12-2015 at 06:56
What am I doing wrong?


I am trying to make a rocket similar to an Estes rocket, but a little more powerful. So far in my project, I have made and tweaked small 25 gram samples in paper casings. My best results thus far are from 1 gram sulphur, 5 grams aluminum powder, 16 grams ammonium nitrate, and 3 grams silicone caulk. I have compacted the fuel in a casing with a hammer and drilled a hole through a clay nozzle and fuel. I have then put a sparkler in through the nozzle and lit the end as a fuse, and my best result was sparks and flame shooting out the back at least 5 feet, yet the rocket produces little to no thrust. Can anyone tell me why this is, or what I could do to improve my rocket?
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[*] posted on 23-12-2015 at 07:04


First off you're not providing a descriptive subject. That alone should be enough to have the post deleted in my opinion. Other than that, use potassium nitrate. Ammonium nitrate is not well suited for simple rockets.



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[*] posted on 23-12-2015 at 07:11


^^^as usual, great answers. Gonna do any of these OUTSIDE. With my kinda luck, if it can burst into flames, it will. My rooster might just get startled today.



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[*] posted on 23-12-2015 at 12:46


Quote: Originally posted by Deathunter88  
Are there any soluble barium compounds that can be used to test for the sulphate ion other than barium nitrate and barium chloride?

I ask because I just recently found out that barium chloride is listed as an "extremely toxic" chemical in China and thus illegal for an individual to possess. Barium nitrate on the other hand is listed as an explosives precursor and also illegal for the individual to own.

OR if no such compound of barium exists then suggestions of any other compounds that can be used to identify the existence of sulphate ion in solution would be welcomed.

Ba acetate
Ba hydroxyde
Ba perchlorate




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[*] posted on 23-12-2015 at 20:24
Formaldehyde


Is there any way to isolate formaldehyde from products that contain it or to synthesize it?I heard about the catalytic oxidation of methanol,but I think there is too much of a risk of igniting the formaldehyde/methanol,so is there any other way to synthesize it?
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[*] posted on 23-12-2015 at 21:29
Green nitric acid


Yesterday i was trying to concentrate nitric acid.

My plan was to bubble NO2 through dilute nitric acid because NO2 hydrolises in water to give HNO3 and HNO2.

but when i bubble NO2 nitric acid turned dark green. Anybody knows why this happend.
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[*] posted on 23-12-2015 at 21:39


Quote: Originally posted by CitricAcid  
Is there any way to isolate formaldehyde from products that contain it or to synthesize it?I heard about the catalytic oxidation of methanol,but I think there is too much of a risk of igniting the formaldehyde/methanol,so is there any other way to synthesize it?
There are many threads about synthesis of formaldehyde, as well as plenty of information elsewhere on the internet... search and ye shall find.



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[*] posted on 24-12-2015 at 03:54


Quote: Originally posted by idrbur  
Yesterday i was trying to concentrate nitric acid.

My plan was to bubble NO2 through dilute nitric acid because NO2 hydrolises in water to give HNO3 and HNO2.

but when i bubble NO2 nitric acid turned dark green. Anybody knows why this happend.

In one of the other posts I read that you bubbled NO2 through 60% HNO3. This is not dilute anymore. What happens is formation of HNO3 and NO. With more NO2 you get N2O3, which is deep blue. Together with the dark brown of NO2 you end up with a green color. Your acid becomes contaminated with N2O3. You can remove that by bubbling a lot of air through the acid.

This effect can also be observed when dissolving a metal, giving colorless metal ions. E.g. dissolve some mercury, bismuth or lead in conc. nitric acid. You'll see that the liquid turns green. This is caused by formation of a mix of NO and NO2 when the metal dissolves. On dilution with water, this color disappears, the N2O3 then is hydrolyzed and bubbles of NO escape from the liquid and the liquid itself becomes colorless.




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[*] posted on 24-12-2015 at 04:04


Quote: Originally posted by woelen  
What happens is formation of HNO3 and NO. With more NO2 you get N2O3, which is deep blue. Together with the dark brown of NO2 you end up with a green color.


Don't you only get N2O3 at low temperatures? I remember holding a chunk of dry ice above a reaction mixture of copper and nitric acid, and getting solid blue condensing on the ice.




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[*] posted on 24-12-2015 at 05:22


In aqueous solution you can have it at room temperature as well. It is quite stable in conc. HNO3, in water it slowly decomposes.

A nice experiment is adding a few drops of a concentrated solution of NaNO2 to dilute acid at room temperature (e.g. 10% H2SO4). You get bubbles of NO/NO2, but the solution turns blue as well. Not really deep blue, but easily observable like a fairly dilute solution of CuSO4 in water. This is N2O3. On standing it slowly disappears and the liquid becomes colorless. If the acid is colder, then the color persists for a longer time.




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[*] posted on 24-12-2015 at 05:48


Quote: Originally posted by CitricAcid  
Is there any way to isolate hydrazine from mixing bleach and household ammonia without adding sulfuric acid? What I mean is:Is there any way to separate hydrazine from everything else that is formed when mixing bleach and ammonia? Can it be separated via sep funnel?

[Edited on 22-12-2015 by CitricAcid]


Quote: Originally posted by CitricAcid  
Back to my original hydrazine question: In addition to isolating it without converting it to its salt form,is there any combination of solvents that can be used to extract it from solution?

Methylethylketone azine...is a good way to go
(CH3-CH2-)(CH3-)C=N-N=C(-CH2-CH3)(-CH3)

Addition of Ni(NO3)2 in the wish to precipitate Ni(N2H4)3(NO3)2 (nickel nitrate(II) tris-hydrazino complex) doesn't work!




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[*] posted on 24-12-2015 at 15:42


Quote: Originally posted by PHILOU Zrealone  


@The_Davster,
Do you have a reference, link or document for me?
That reaction is very interesting for the energetic materials field ... and I'm passionnate about those...



t-Butyl azoxy next to an amine on an aromatic system.....nitrate and the nitramine cyclizes with the tBu azoxy giving a 1,2,3,4-tetrazine-1,3-dioxide. Sexy molecules.

Easier method is to take R-N=O and react with tBuNBr2 (Horrible smell :o) giving R-N(O)=N-tBu. Those working with these type of molecules seem to have abandoned the grignard for the nitroso route.


Attachment: how to make THE GRIGNARD.pdf (332kB)
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[*] posted on 26-12-2015 at 17:22
Flourine


I'm aware of the dangers of fluorine, but I was thinking; If someone wanted to do some flourine chemistry, would it be possible to use PTFE plastic or plumbers tape and extract it from that? It's just idle curiosity, I have no current plans to play with anything this reactive.

Bit late, but Merry Christmas ;)
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[*] posted on 26-12-2015 at 18:08


Hey, zts. I had just composed a reply and lost it because the thread was merged when I submitted. Your efficiency is spectacular.

@maker
firstly, "flourine", I presume is a derivative of flour and probably harmless. ;)
Secondly, in case no one has done so already, welcome to SM. Stick around. Read the faq. Learn some stuff -- there is a wealth of great knowledge here.

Thirdly, addressing your question:
Fluorine is nasty stuff. Few have a legitimate reason to work with it and those that do have specially-equipped labs. Those without the technical equipment and proper knowledge tend to have a lower life expectancy. Fluorine killed a number of people on the way to being isolated. One of the issues is the fact that it is so oxidising that it is next to impossible to contain it in anything. Another issue is the tendency to form HF and kill the experimenter. The difficulties ramp up from there.

PTFE is probably the compound that contains the most fluorine on a mass basis that you are ever going to encounter. It is quite ironic that it is so safe. There is a cool experiment that you can do with it if you have some powdered aluminium or Mg. (The powder must be fine however.) Dust some metal powder onto some plumbers' tape and hit it with a flame. You get a nice thermite-style reaction that is reasonably impressive. You do need to do this outside. Avoid breathing any fumes because some of the byproducts are toxic. It is reasonably spectacular and simple to do.




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[*] posted on 1-1-2016 at 16:17
What's this white growth in my Copper Acetate solution?


I just made some Copper Acetate from some copper wires to grow crystals from. I dissolved the wires in vinegar and hydrogen peroxide.
As the copper acetate solution evaporates, dark colored crystals start growing at the bottom. After a few days I also start seeing a white substance growing on my crystals.
Does anyone know what this is? And how do I prevent it from growing so I can get some nice clean crystals?

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[*] posted on 1-1-2016 at 16:26


Basic copper(II) acetate. Add more acetic acid.



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[*] posted on 1-1-2016 at 19:18
Stannous chloride


I use stannous chloride in my glass blowing operation to iridize my glass. After spraying I try to recycle the residue but some contaminants are present so the contaminants slowly build up and affect the quality of the iridescence .I put the sludge or residue in a crock pot and cooked off all of the liquid but got a grayish crystalline rock.I have re-dissolved the rock but got a dark brown liquid with a lot of sludge on the bottom of the jar.I drained to brown liquid from the sludge I think this residue is my earlier contaminates . The brown liquid fumes ok but I would like to try and get back to a clear crystal stannous form that I started with .Any suggestions.
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[*] posted on 1-1-2016 at 20:53


Quote: Originally posted by j_sum1  
Hey, zts. I had just composed a reply and lost it because the thread was merged when I submitted. Your efficiency is spectacular.
Oops, sorry about that!



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[*] posted on 2-1-2016 at 20:28


Quote: Originally posted by DraconicAcid  
Basic copper(II) acetate. Add more acetic acid.


Oh man. This has actually been a "burning question" of mine. "Basic" Copper Acetate?




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[*] posted on 3-1-2016 at 18:39
Valency of manganese


Solution were yellow, precipitated hydroxide, light brown lumps.
Decomposed to oxide, it got darker, and lost weight.
What valency of Mn is this?
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[*] posted on 3-1-2016 at 18:47


Likely IV, at this point.



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[*] posted on 5-1-2016 at 03:46


Soluble starch

I have searched the forum and a few other places but have not found good answers to this question.
I am after some soluble starch for use in redox titrations. I know I can buy some from a chemical supplier but I was wondering about otc sources and preparations. I didn't think it would be too hard -- in fact, it is probably so basic that no one writes much on it. Here is what I have tried and my thoughts so far.

1. I have thus far used white rice and washed it with cold water and collected the water from it. This has been somewhat successful in the past but my last attempt produced something that was not sensitive enough. I have also used hot water but always gotten a cloudy mixture that was not really any better for the task.

2. I have read a few very simple procedures for obtaining potato starch. Basically juice a potato and dry it. I am not sure how good this would be for iodometry. It is usually presented as a the first step in making a simple plastic product. I think that the starch obtained is quite soluble though.

3. There are quite a number of starchy products available at the supermarket. What the Americans call cornstarch is called cornflour here and I am not sure how it is processed or how it will act. Again, it is not that soluble. Arrowroot seems like it would be a better option. Alas, old-fashioned laundry starch has been replaced by ezi-iron.

4. There are numerous papers available that describe ways of making starch more soluble. There are modified starches. There are procedures for separating it from gluten and other proteins. It can be made into pre-gelatinised starch through a laborious process of soaking in DMSO and then baking for a long period. There is also the possibility of treating it with alkali or "a salt" to make it more soluble. I am not sure how one then effectively gets rid of the alkali or salt or even if it matters too much. I guess if I used ammonia then any excess would evaporate away.


So, (and I know I am probably overthinking this) what is the best course of action? Is juicing a potato ok? Or should I buy some tapioca flour? Is it worth treating it in any way? Or is best to spend a couple of bucks at a chem supplier?




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[*] posted on 6-1-2016 at 15:25
Hydrogen peroxide extraction from sodium persulfate


Is it possible to obtain hydrogen peroxide from sodium persulfate

If you know the answer please reply I am cuurently looking at alternative methods for obtaining hydrogen peroxide because it is a bit costly
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[*] posted on 6-1-2016 at 15:48


Quote: Originally posted by arkoma  
Quote: Originally posted by DraconicAcid  
Basic copper(II) acetate. Add more acetic acid.


Oh man. This has actually been a "burning question" of mine. "Basic" Copper Acetate?

Yes. Like Cu(OAc)2*Cu(OH)2
where OAc is the acetate ion. (not actinium stuff). What question do you have about it?




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[*] posted on 6-1-2016 at 16:10
catalyic decomposition of hypochlorite


Is there a chemical which can decompose bleach into oxygen without being consumed like maganese dioxide for sodium chlorate

[Edited on 1-7-2016 by zts16]
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