# UC

Senior Members

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## Posts posted by UC

### find the equation for

Informed guessing is a surprisingly reliable technique. Given that weight is density multiplied by volume, what type of relation do you think you'll have between weight and length?

<nitpick mode>

*mass* is density multiplied by volume. Technically, newtons would be the SI equivalent of weight, not grams.

</nitpick mode>

### Question bout alkoxides

This is purely for my own curiosities sake...

If one were to take an ionic solution in an alcohol, and then run current through it, would the product be the alkoxide in solution?

Something along the lines of sodium chlorode in methanol with a few volts using graphite electrodes. Would this produce hydrogen at one electrode (from the alcohol) and chlorine at the other electrode (from the salt) leaving sodium methoxide in solution?

What if the anion was something more difficult to oxidize, such as sulphate?

Thanks!

Well, sodium chloride doesn't dissolve in alcohol, but that's besides the point. The easiest way to do this is by azeotropic drying or using molecular sieves. An alkali metal hydroxide dissolved in alcohol exists in the following equilibrium: [ce] KOH + MeOH <-> KOMe + H2O [/ce]. If the alcohol is truly dry, the equilibrium will lie to the right unless the concentration of KOH is very high (and therefore the concentration of water would be high). If you can remove the water from the reaction mix, then Le Chatlier's Principle states that the equilibrium will push to the right. Either you use molecular sieves to adsorb the water or you add a solvent that has an azeotrope with water, but not MeOH, which will boil off below the boiling point of methanol. simply distill off the azeotrope to drive the proportion of methoxide up.

Also a problem with the original idea is that alcohols are not nearly as polar as water and as such, even ionic compounds that dissolve in them don't disassociate into ions nearly as well. These solutions are in most cases not good conductors.

### Electrolysis Problem

salter, use the chem notations: [ce] CuCl4^2^- [/ce] for example, delivers [ce] CuCl4^2^- [/ce]

Magnesium hydroxide would be formed in aqueous solution, not the oxide, and it is white colored. The black may be nonreactive contaminants in the Mg though.

I also think you mean cathode instead of diode Very different things.

### extracting magnesium

No, all you will produce in aqueous solution is magnesium hydroxide. You need molten salts of some kind or a mercury cathode to get metal.

### nitration of p-nitrobromobenzene

Hi guys,

I was wondering if a second nitration of p-nitrobromobenzene would work using a catalytic amount of H2SO4. I was thinking it wouldn't since both the bromine and the Nitro group are deactivating substituents, and since its a catalytic amount of acid it wouldnt add a second nitro group onto the ring. Is this reasoning correct?

No, there is a reason nitrating mixtures for polynitrated benzene rings contain a vast amount of concentrated or fuming H2SO4. The ring is extremely deactivated, especially in this case. Any chance or reaction would require a high concentration of nitronium ion, and probably a fair amount of heat.

### Sulfur Dioxide

Recently, I've been thinking a lot about preparation, and methods of synthesis for various compounds, and amongst these, I've thought of producing sulfur dioxide, in some way, and then bubbling it, through a dilute hydrogen peroxide solution. As a result, this ought to yield a dilute solution of sulfuric acid, (about 5%, based upon the type of hydrogen peroxide I have) which I can then boil down, to a more feasible concentration of approximately 15%. (roughly) My only question, is in the generation of the sulfur dioxide. The obvious way, would be simply to burn sulfur crystals, however, I am unsure where these can be obtained. I do have various sulfides, (FeS2), which I suppose, upon heating, will oxidise to SO2 and Iron Oxide, but what are the temperatures required, for a such a reaction, to take place? If the temperatures are too great, are there any alternative ways to produce the SO2, necessary for the reaction to occur?

Ps: To prevent any furor, or unnecessary worry, I am well aware of the toxicity of this compound, and shall plan accordingly.

Roasting metal sulfides takes an enormous amount of heat and is appropriately done in a furnace. H2O2 is generally stabilized with assorted phosphate buffers or phosphoric acid, which means that these will be contaminants in yoru final product. By far, the easiest way to make lots of SO2 is to acidify and gently heat bisulfite or metabisulfite (available fairly cheaply from wine making supply places) solutions. Ideally you'd use a non-volatile acid, perhaps cheap drain cleaner H2SO4 that is too dirty for direct use. The gas will have water vapor in it, but I doubt that is an issue if you're trying to just make aqueous acid in the end.

If the seals on your apparatus aren't good enough, you will definetly smell it. The SO2 causes a choking, burning sensation and always makes me sneeze a lot.

### Aqua Regia, HCL, NH4NO3?

No need to dig up dead posts, plus, when you do a thermodynamic calculation the fact that the enthalpy is positive means that it is most likely NOT going to occur. For reactions to be spontaneous, the Gibbs Free Energy MUST be negative. Only in cases where there's a significant entropic difference will a reaction that is positive in enthalpy be spontaneous. So no. Your proposed reaction will not work. If it did, people would use HCl and NH4NO3 for making aqua regia since NH4NO3 is much easier to store than HNO3.

This is actually well known, although NaNO3 is usually used for other reasons. In solution, you have ions. The salts are not very important. As long as there are NO3-, H3O+, and Cl- ions around, the mixture will work. It is known as "poor man's aqua regia." Using NH4NO3 is a bad move since chlorine might be formed in the mix and this would make chloramines or NCl3 and also because the ammonium cation precipitates things like Hexachloroplatinate ion, preventing dissolution.

Merged post follows:

Consecutive posts merged
yes, HI is colorless in aqueous solution. colored solutions tend to have transition metals.

This isn't the issue so much as iodine stains. HI is very easily oxidized by air to I2 and solutions often have a brown tinge from this.

### NH4Cl > NH3 + HCl?

interesting problem

electrolysing probably won't work, it will probably get you

2NH4Cl > 2NH3 + H2 + Cl2

BAD idea. This nets you NCl3, a very powerful, dangerous, and unstable compound. Think NI3, but liquid and not safe when wet.

You can get one or the other out, but not both at once. Either use H2SO4 for the HCl or use NaOH solution for the NH3.

### Food combinations that cause "short circuit" ?

Certain wild mushrooms can be extremely dangerous if there is also alcohol in your system.

### Anodes and Cathodes

mmm, i could have been more specific. I mean aqeous NaCl, with a view to obtain Sodium Hydroxide, Chlorine and Hydrogen.

Do a google search for chlorate cells and read about the electrode choices (especially the anode, which is under much harsher conditions) for them. These are the rugged kind of electrodes you want.

### copper (I) oxide

it has very strong red tones, but the substance is still submerged. when i remove the cathode from solution, the substance clings to it. kinda looks felt like in texture as one would expect, but it is very much more red at that point. i thought that maybe the maroon look was the fact that it was still 'wet', the same way stones get deeper in tone when wet. i have not yet been able to dry it out completely don't want it exposed directly to air for fear that it would oxidize more...

to dry i was thinking of using everclear and a homemade vacuum pump or something, if you have any good ideas about how to dry it without letting it oxidize please tell me. i tend to make things more complicated at times.

is there a way to test if it is indeed copper or copper(i) oxide?

Throw it in concentrated HCl. If it dissolves, it's Cu2O. If it doesn't dissolve, it's Cu. If it looks like this: http://en.wikipedia.org/wiki/File:Synthesizing_Copper_Sulfate.jpg it's definetly copper.

### copper (I) oxide

copper(I) oxide isn't maroon or red. I suspect the red build up is copper.

Actually Cu2O (which is what is being discussed) is red (maroon, however is most likely copper). If the buildup is actually *on* the cathode, it is certainly copper metal.

Please see here for making Cu2O by electrolysis:

http://webpages.charter.net/dawill/tmoranwms/Chem_Cu.html

### HNMR and Synthesis Need Some Help Please!

Yeesh, lots of questions. For proton NMR spectra, all you need to consider are 1: how many equivalent hydrogen are there for the peak (gives the integration, this is where symmetry matters) and how many hydrogen is it coupled to (ignore hydrogen attached to the same carbon, symmetry does not matter).You're first guess is correct, in that it is coupled to 8 hydrogens and thus is a nonet.

For part D, you do need to count both methylene groups and thus you have a pentet.

For questions 2 and 3, I want you to tell me why you are picking the reagents you picked. what does each do? what are the intermediate products?

For 4, because acidified dichromate is a *powerful* oxidant. You can make aldehydes with dichromates, but it's fairly tricky.

For 5, http://upload.wikimedia.org/wikipedia/commons/b/b9/PBr3_alcohol_rxn.jpg What kind of mechanism is that and what does it do to stereochemistry?

For 6, because neopentyl groups are so hindered that they don't really undergo SN2 reactions. The alternative is SN1, which involves a carbocation and alkyl shifts can occur.

For 7, look up zaitsev's rule.

### eucalyptus

Looking at the structure, I don't see any way to modify the molecule itself to make it water-soluble. Your best bet would be to formulate it with a surfactant or lipid, to make (respectively) an emulsion or a liposomal dispersion. It won't be a true solution. However, liposomal dispersions can be lyophilized to a dry powder (often mixed with mannitol), and instantly reconstituted by adding water.

Sulfonating just about anything makes it water soluble, but then it'd be nonvolatile. But you'd also probably cleave the ether if you use oleum.

### Lab Glassware acquisition

You asked this as user thorazine on sciencemadness.org. You got their opinion. Indian and Chinese glassware is in general, garbage and you're better off buying pyrex or kimax stuff if you don't know the quality. It's pricier, but the odds of the pieces breaking in the long run is lower so you'll save money and much frustration.

### eucalyptus

couldnt the ether group be protonated in acid making it more soluble? or would it just decompose?

Oxonium ions are not very happy ions, unlike ammoniums for example. They will slowly hydrolyze in strongly acidic conditions, but the equilibrium for [ce] R2OH^+ + H2O <-> R2O + H3O^+ [/ce] lies far to the right.

### Distillation Observations

Sounds good. Just one question. Why should I turn the adapter right side up? It always seems like more vapor gets through if its horizontal...

I do like the idea of insulation. I hadn't even thought of that.

I had actually used boiling chips, but they were glass, and now that I think about it, that glass is far too smooth to help. I'll use some plaster.

Yes, in the Liebig, I got just fog, but because of the rapid changes in pressure in the system (boiling chips didn't even help that much unfortunately, i'll make them smaller next time), the liebig just helped keep that fog in an area where the gas was constantly cooling (and any condensed liquid would run away from the boiling system instead of back into the boiling flask).

thanks for the tips, seriously

It's not about the size of the boiling chips. It's about surface area. One or two real boiling chips probably have the surface area of hundreds of crushed bottles. It provides a good place for bubbles to nucleate. The sulfuric acid fog *is* condensed, just not sticking well to itself. The shape of the graham condenser forces it to hit walls and coalesce into drops.

The rationale for the upright 3-way adapter is to minimize strain on the glass joints and keep acid from pooling in the 3-way (and possibly seeping past the stopper).

### Distillation Observations

activated charcoal would be better than plaster as plaster would react

No it wouldn't. Plaster of paris at least, which is what most people think of when they think of plaster. It's calcium sulfate. Sulfuric acid isn't going to do anything to it.

### Distillation Observations

Turn the 3-way adapter right side up. Use a smaller heating flask or insulate the larger flask to get more vapor into the condenser. I'm surprized one condenser (or even an air condenser or no condenser) wasn't more than enough. Drop the leibig entirely. The graham just gives more surface area for the H2SO4 to condense on. Were you getting a "fog" of acid in the liebig? Use boiling chips. Activated carbon or maybe even some powdered plaster is probably acceptable as a crude replacement.

### Photography with silver chloride

You should talk to YT about this stuff This is one of his specialties. He's probably asleep right now, so I'll give it 4 or 5 hours before he responds.

### Making Hydrochloric acid + Halogen Chemistry.

add conc sulphuric acid to NaCl and heat. bubble the gas through H2O and roberts your fathers,sisters brother.

test it with conc ammonium hydroxide white mist of NH4CL should appear.

hope this helps, i havent done it yet but i will be trying it soon.

*SLOWLY* add conc. sulfuric to NaCl. Only use gentle heating when no more HCl is evolved by itself. If you dump the two together and heat, you will probably blow the connections on your apparatus open and end up inhaling a cloud of HCl, as will everyone nearby. Lead the gas to the surface of ice cold water. Bubbling probably isn't necessary as the solubility is quite high and you will end up with suckback.

### chemical reactions producing oxygen

[ce]KClO4 -> KCl + 2O2[/ce]

I believe that [ce] 2KClO3 -> 2KCl + 3O2 [/ce] using a bit of [ce] MnO2 [/ce] as catalyst is going to be easier. It is my understanding that [ce] KClO4 [/ce] is quite a bit more stable and resistant to decomposition.

As far as I can tell, this is just homework, so dangerous isn't really an issue. If it is actually being done, I suggest adding washing soda and activated carbon to some plain 3% hydrogen peroxide. The bubbles will be [ce] O2 [/ce].

As for big314mp's suggestion, that's not quite a reaction to some ears, although it will do as advertised.

### How can you tell when a compound undergoes electrophilic aromatic substitution?

As for activating versus deactivating, more electrons make an aromatic ring more reactive. Are fluorines going to give or pull away electron density from a carbon? (look at electronegativities if you aren't sure) If that carbon is attached to an aromatic ring, is it going to give or borrow electron density to compensate for what the fluorines are doing to it? If you have any further questions, feel free to ask them here

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