UC

Aaaaaaaaah!

 

You should've posted that YESTERDAY so I don't look like such an april fool!

 

Sorry, I'm just SERIOUSLY not thinking right. I've got a hardcore case of brainfry from doing calculus all day for some site engineering junk on a couple projects. I'm leaving for a vacation soon, so naturally I've got to finish all my work before I leave. It's 8:15 PM, I'm still @ work, and I've got about 2 days more work to finish by 5 PM tomorrow.

 

I need to go home & drink some distilled potato juice with an olive in it or something!

 

My WIFE would really love me if I could figure out a way to make Crystalline Carbon via Electrolysis!

You can actually, but nothing you'd be able to make a ring from. Diamond-like carbon coatings are available under certain conditions by electrolysis of alcohol solutions, and poly-(hydridocarbyne) (a precursor to lonsdaleite (hexagonal diamond)) can be prepared by electrolysis of bromoform in acetonitrile with a salt (LiBF4 perhaps) added to aid conduction.

I knew that:mad:

Sorry, that wasn't directed at you. I'm know you know, but should I have left it uncorrected?

Ah, you are right... I'm missing the hydromium ion, as well as vinyl, in regards to nonmetallic cations. Anything else? I'm at a loss. Its a shame that herein lies an exception, as these exceptions are more than rather annoying. Totally unrelated, but I'm actually rather surprised you haven't heard of double- displacement reactions, it is one of the basic kinds. I suppose you probably learned of their separate classes, neutralization, and precipitation reactions. Shame...

thanks for the help,

 

,Theo

You seem to be confusing simple high school level inorganic chemistry with all of chemistry. There is no sharp line. There are carbocations of all kinds, polyatomic chalcogen cations, metal clusters, substituted and quaternary ammoniums, ylides, phosphoniums, all manner of metal complexes, etc.

There is also no hard line between ionic and covalent. Things can have partial positive charges and solvent choice makes a world of difference.

a "double displacement" reaction is better described as a redox reaction.

Actually, a "single displacement" reaction is a redox reaction using an element. A double displacement is a metathesis reaction, and as far as inorganics go, it's only considered to have happened when there is a precipitate. If you wanted to do the reduction of say hexachloroplatinate anion by hydrazinium sulfate, I think the high school way to look at these reactions would crap its pants. Then there are olefin metathesis reactions using things like grubbs' catalyst, but let's not get into that.

Ah, to make a magnet, get an iron nail, wrap some wire around it in a helix (just toward one direction), then connect the setup to a power source (eg 12 volt battery). Now you have an electromagnet. If you want to magnetize the nail, heat the nail and let it cool inside the coil connected to a DC power source.

Insulated wire, please. Otherwise, it just shorts out.

In terms of how well, or easily two compounds react its true that reactivity is really an inexistent concept in that modifying temperature, pressure, addition of a catalyst, etc. Cand easily modify this, however, when I generally think of reactivity, I think of the metal reactivity series, which can also be thought of as an inverse of the standard reduction potentials E(insert degrees symbol here) red (reduction). This is often used to predict what occurs, in the case of a redox reaction, however, as I have learned, even this does not bear 100% accuracy.

 

(Well, science, in nature, being a "real," and not an abstract concept (like mathematics), can never be proven. For example, let us say that I am to make the statement that all sheep are white. No matter how many white sheep I find, I can never be proven correct, while the appearance of but one black sheep, shall be enough to disprove my theory) (The example I used is a rather famous one. I would state it as a quote however, I cannot remember who to attribute this to.)

I would read up on equilibria and standard enthalpies of formation.

Yesterday, since I'm teaching electrochemistry at the moment, I decided to try to replicate THIS experiment.

 

I had some pure lead sheets, which I tore up to make electrodes. It worked beautifully, and I had none of the troubles with the strange material from the anode, which to me suggests that Woelen's hypothesis that the material is to do with the impurities in his electrodes is correct.

 

I'll take some photos one day perhaps maybe.

I did this with some stannous chloride solution with tin anode and graphite cathode and it works great as well.

I got something similar to that strange stuff, which i think had to do with the breaking of electrical contact with some crystals (the solution was still acidic to prevent hydrolysis) and small amounts of antimony in solution.

So i decided to grow some silver crystals.

The conventional way is dissolve [ce]AgNO3[/ce] in water and put in copper, then the copper dissolves to make copper nitrate and the silver precipitates out.

I figured that i would just dissolve silver in nitric acid myself... So my question is this: do I need to separate the [ce]AgNO3[/ce] from the nitric acid or could i saturate the acid with silver then just put in the copper?

Well, depends how much nitric acid is left. If you aren't dissolving any more silver with hot acid, I think it would be fine to just cool, pour off the liquid and use it directly. Otherwise the copper and very fine silver crystals will dissolve.

I didn't mean exactly as a metal, but rather as a cation, as the ammonium ion is the only nonmetallic cation, with the exception of hydrogen.

Since when? You're missing a vast amount of chemistry somewhere.

"Ammonium" as a metal doesn't have much of an existence, but a solution of it in mercury does. Here's a reference.

http://www.vias.org/encyclopedia/chem_ammonia_salts.html

 

Since it decomposes in water to give hydrogen I guess it can be viewed as reducing the water but that's not clear.

That sounds exceedingly dated. I suspect the main component is sodium amide.

there is a big bang inside of us all..

That's what she said.

I'm trying to make thermite; what do I need to make it? From what I've read it said just mix aluminum powder with iron(III) oxide (rust). Is this correct? I just want to watch it burn.

 

PS-also, which equation is better? Both are for thermite, but which one is the better... um... thermite?

 

Fe2O3+2Al -->2Fe+Al2O3+heat

3Fe3O4 + 8Al --> 9Fe + 4Al2O3 + heat

The internet is a good place to find ways to hurt yourself. If you just want to watch something burn, make a campfire. Thermite is not *that* impressive. However, it is extremely hazardous, especially for someone who 1) clearly has no grasp on the concepts behind the reaction and 2) seems to just want to goof around. If you do something wrong with thermite, it'll sling molten iron around, maybe all over you. Water won't come fast enough to save the body part it lands on and it certainly won't stop the main reaction. Try explaining that to the emergency room staff.

Thermite can be handled safely, but if you need to ask here or anywhere else about how to make it, you shouldn't.

I have a way of doing this but I feel like there's got to be a quicker route. Here's how I did it:

 

1. HCN

2. LiAlH4

3. H2O

4. H2SO4, -H2O

5. H2, Pd

 

(i.e. cyanohydrin formation, reduction, acid-catalyzed elimination to get rid of OH, then H2/Pd to get rid of double bond)

 

Could it maybe be done with a Wittig reagent?

You have two reductions. I'm not sure, but you may be able to do both with the hydrogen and Pd on carbon.

Alternative to cyanohydrin: http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv4p0221

Ultra pure water has a very high resistance. Any electricity you try to pass through it either doesn't flow at any meaningful rate or has effectively all of it's energy dissipated as heat. When you add ions that can ferry charge, you allow current to flow and the energy is instead (some of it at least) used up breaking H2O into H2 and O2. When you burn the H2 and O2 again, you get back *all* of the energy that went into splitting them (but none of the energy lost to the water's still present resistance, or dissipated in the wires or internal resistance of the power source). No engine can harness all of that energy, in fact, it will lose the majority of it as heat to the surroundings. Overall, you lose most of the energy you put in with the original electric current.

Stronger acid in which way, more corrosive or donates more hydrogen ions?

 

Because for example hydrochloric acid is a stronger acid then hydrofluoric, but hydrofluoric is much more corrosive.

 

Nitric acid dissociates slightly worse, but the nitrate ion has more oxidizing potential then the chloride ion, meaning nitric acid is more powerful not just due to the oxidizing power of nitrates...(i.e. even nitrate salts are oxidizers).

 

I agree YT2095, the oxygen from hydrogen peroxide helps strip the hydrogens off the acid, so that they are not left unpaired.

The corrosiveness of HF is only due to the powers of the fluoride anion. Chloride ion has just about 0 oxidation potential. The oxidizing power of nitrate should have nothing to do with how "strong" it is. The statement you make about the H2O2 is a load of bollocks. he's saying that if you make HCl and H2SO4 oxidizing, you will see that you are mistaking oxidizing power for acidity.

Please look up the Hammett acidity function.

At any rate, the strongest acid that can exist in aqueous solution is [ce] H3O^+ [/ce]. To what degree the acidic protons of a given acid choose to protonate water to make hydronium ion is the main concern.

I'm sorry if this is the wrong section, perhaps it belongs in physics, but

 

how cold is ice + salt + water? I'm wondering because most of the places I look it says like -10 degrees, which seems about right if you use the freezing point depression equation.

 

However, I was easily able to condense a drop of chlorine (I'll get a better picture soon, but here), whose boiling point is -35 degrees celsius, -30f. I used water + ice + a good amount of magnesium sulfate and sodium chloride.

 

I think maybe because I used two different salts, i was able to circumvent the common ion effect...

 

oh, and kids: don't make liquid chlorine.

no. There's an almost 100% chance that is just cold water with chlorine dissolved in it.

Thanks for the muriatic acid tip. I recently consulted, a friend, only to discover large 34% buckets, being sold at home hardware. Absolutely brilliant stuff.

 

An interesting but unrelated thought: on first handling the mixture, I felt a slight stinging sensation, however, I rinsed my hands thoroughly to be sure, and in retrospect, it was probably only a figment of my mind's early anxiety, just as one feels as if their head is itching, at a checkup for lice, despite the pests' inexistence.

That's what gloves are for.

Never mix Ammonium Nitrate with Copper!

explosive TACN can be formed, which it is doesn`t go off, is certainly Plenty enough evidence to get you arrested and sent away for a VERY long time!

You do realize that this has been beaten to death as a subject on sciencemadness? I believe it's the anhydrous material that is truly dangerous, and I'm not sure that anyone there has been able to produce it, though I don't read energetic materials forum much.

Also, the formation of the ammine complex requires that the solution have a reasonable amount of NH3 in it. Mixed with strong HCl, this is extremely unlikely. Splash some HCl in some dark blue tetraamminecopper (II) solution. The color of the complex (and the complex itself) vanishes upon acidification.

The NOx formed will end up destroying a good portion of the ammonium ion anyway, via nitrous acid. IIRC, NH4NO2 *can* be made with ammonia and hydrogen peroxide, but is only "stable" in the cold, decomposing on heating to nitrogen and water.

The hydrogen would have to be reacted with something like platinum to make it reactive if im not mistaken. There is also a light frequency that will break the bond also so it may be worth a look to see what frequency its at. I have it somewhere but its to late to look. Ill check back tommorow.

for plain hydrogen? The energy is well above visible light and will rip apart all sorts of bonds, not just hydrogen. precious metal catalysts are useful for catalytic hydrogenation. I suspect you make something more along the lines of a mess or poisoned catalyst if you put it in a metal salt solution.

As was said, simple bubbling will not achieve your desired reduction, nor will any easy method in aqueous solution. You're better off using another metal or something like hydrazine to reduce the copper. Alternatively, you precipitate copper hydroxide, roast to the oxide, and heat that in a hydrogen atmosphere to make copper powder. The real question is, why bother. Electrolysis readily plates copper sponge onto the cathode, and bright copper if the solution is acidic and contains appropriate levelling agents. I'm under the impression that you were after H2SO4 though, in which case, give up. If it were that simple, it would be widely known.

I would greatly advise against this. I believe someone over on sciencemadness managed to get clean bromine out of 1-Bromo-3-chloro-5,5-dimethylhydantoin (which is what I think those tablets are) but they used precise amounts of appropriate acid and distilled the bromine off.

Halogens are all pretty horrible and I doubt you have the proper equipment to handle them. On the other hand, there is a fairly simple way to make bromine on woelen's site that even a beginner could probably follow.

The gases are soluble. Halogens are all at least slightly soluble.

I already have this experiment running-

20.2% HCl + excess ammonium nitrate = poor man's aqua regia

 

It is slowly eating a piece of copper pipe (and we all know HCl could not do this by itself)...

 

the solution is becoming a very dark greenish, anybody know what this copper compound is?

 

a little bit of chlorine is being produced- you can see its distinctive color on the surface

 

is it really just copper chloride? then how does ammonium nitrate come into play?

It's some sort of mixed valency Cu(I),(II) complex involving chlorides. It's fairly common to see this. Please do this outdoors, if you aren't already. Aqua regia evolves dangerous NOCl fumes as well as NOx.

Due to the ammonium ion, the web of reactions occuring is even odder. Here's a short rundown of some of what's going on:

[ce] 4H^+ + 2NO3^- + Cu -> Cu^+^2 + 2NO2 + 2H2O [/ce]

[ce] Cu^+^2 + Cu -> 2Cu^+ [/ce] (This is aided by chloride complexation of the Cu(I) ion as soluble CuCl2^-)

[ce] Cu^+ + 2H^+ + NO3^- -> Cu^2^+ + H2O + NO2 [/ce]

[ce] 3NO2 + H2O -> 2HNO3 + NO [/ce]

[ce] NO + NO2 + 2H2O <-> 2HNO2 [/ce]

[ce] HNO2 + NH4^+ -> N2 + 2H2O + H+ [/ce]

[ce] HNO2 + HCl <-> NOCl + H2O [/ce]

[ce] HNO3 + 3HCl -> Cl2 + 2H2O + NOCl [/ce]

[ce] Cl2 + Cu -> Cu^2^+ + 2Cl^- [/ce]

[ce] Cl2 + 2Cu^+ -> 2Cu^2^+ + 2Cl^- [/ce]

[ce] 2NO + O2 -> 2NO2 [/ce]

I'm sure I've missed a few.

buy some chalk

Chalk isn't usually chalk anymore IIRC. I think it's almost always gypsum ([ce] CaSO4*2H2O [/ce]) nowadays. I precipitated mine from boiling CaCl2 solution using sodium carbonate, but only because it's an enormous pain in the arse to try to get CaCl2 back out of solution. If you have the time, you can dissolve eggshells or seashells in acid, filter, and add carbonate to precipitate CaCO3.

Good quality limestone is pretty pure CaCO3. The common impurities are MgCO3 (form a complete solid solution, IIRC) and iron oxides.

you know, actually, if you gently electrolyse, the ammonia produced at the cathode can dissolve back into the water faster than it is generated. i think ammonia dissolves easier then chlorine, since Cl2 exists in equilibrium with HCl and HOCl... but at worst for every two steps forward of seperating NH4Cl, you'll be taking one step back of reformulating it...

 

this will eventually give you NH4OH solution, with a trace of NH4Cl... But thats basically the NH3 part...

 

the gases that do escape will be hydrogen and chlorine (with some ammonia and oxygen)... you can collect them and a match or spark will turn them back into hydrogen chloride... with some ammonium chloride, water, and NCl3:eek:...

 

you can leave it as HCl or maybe bubble it into water to get hydrochloric acid... But thats basically the HCl part.

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actually i guess NH3 also exists in equilibrium with NH4OH...

 

none the less, you can use even a very slight difference in solubility to seperate them...

 

and you can, for example, put one electrode in the bottom of the solution, and one almost on the surface, so that one of these- Cl2 or NH3- can dissolve easier...

 

and the ammonium chloride produced would just be electrolysed again... only thing is that if you let chlorine dissolve, you would also get HOCl with the HCl... i dont know if they are easy to seperate...

At least half of this is nonsense. If you electrolyze NH4Cl, you will make plenty of NCl3. Any NCl3 at all is too much. End of discussion, unless you'd like someone to seriously hurt themselves trying any of this.

http://www.cem.msu.edu/~reusch/VirtualText/intro1.htm <- online organic chem text

http://www.orgsynth.org/ <- excellent resource

http://www.periodictable.com/ <- because it's awesome

I assume you mean more than just sputtering...can you give us any more specifics on what you're trying to do?

ah, that's interesting...

 

i do apologize, my experience with chemicals does not exceed what you could buy at the local supermarket, so most of the things i know are looked up...

 

but i do like the idea of bubbling SO2 into H2O2 instead of SO3 into H2O...

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I was down at a hardware store today, and i happened to find a big case of sodium metabisulfite, used for treating rust stains (i guess Fe2O3 + Na2S2O5 > Fe + Na2S2O8 or something similar?)...

 

i read somewhere that sodium bisulfite releases S2O when mixed with water, but i hope you have the equipment to catch it!

no, just metabisulfite is a good reducing agent and is acidic. [ce] Fe2O3 [/ce] is fairly easily reduced to Iron (II) ions under these conditions and dissolves. Metabisulfite is just a dehydrated form of bisulfite and yes, [ce] SO2 [/ce] will escape an aqueous solution of it.