Sciencemadness Discussion Board - The Short Questions Thread (4)

The Short Questions Thread (4)

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Pumukli - 25-7-2014 at 03:50

Papaya,

Are you sure you titrated Na2CO3?
I mean was not your "standard" contaminated with either H2O (crystal water?) or NaHCO3?

Next time try to heat your Na2CO3 in a metal pan for a while say around 120 C before the titration. This way you can elimante both "contaminants", water will evaporate and NaHCO3 would decompose to give Na2CO3.

AlphaDecay - 25-7-2014 at 09:04

My dad broke his mercury thermometer, but the Hg didn't spill anywhere, however it is stuck in the thin glass tube.
How can I extract it from the glass? Perhaps heating it? Im asking for help due to mercury toxicity, I don't want to poison anyone.

Sorry for english mistakes

gdflp - 25-7-2014 at 12:20

DON'T heat it. Mercury metal isn't that toxic but mercury vapor is extremely toxic. Try a syringe with a long end. You could also try to drill a small hole in the other end of the thermometer, I suspect it isn't coming out because there is a way for air to get in.

AlphaDecay - 25-7-2014 at 14:23

Yes, the mercury tiny glass tube broke and the mercury is in contact with air, but it is too thin too use a syringe, and it doesn't fall from the broken thermometer...

gdflp - 25-7-2014 at 17:00

You could try(carefully) breaking the other end so that the mercury will drain out. I assume you want the liquid mercury and not a mercury compound?

[Edited on 26-7-2014 by gdflp]

AlphaDecay - 26-7-2014 at 17:07

I thought of storing it as a metal to use later for making its compounds, but why the question? Imerse the glass with mercury in acid?

TheChemiKid - 26-7-2014 at 17:13

I would attempt carefully breaking the thermometer over a large necked glass container. The mercury and some glass will fall in, and the glass can easily be separated.

AlphaDecay - 26-7-2014 at 17:21

Yes, this seems a safe and easy way to solve the problem!

TheChemiKid - 26-7-2014 at 17:22

Make sure that all of the mercury gets in the glass, and do the entire thing in a bucket, that way if there is a problem, any spills are contained.

gdflp - 26-7-2014 at 17:28

Quote: Originally posted by AlphaDecay

I thought of storing it as a metal to use later for making its compounds, but why the question? Imerse the glass with mercury in acid?

Yes, you could just immerse the glass in nitric acid if you wanted to get mercury nitrate.

AlphaDecay - 28-7-2014 at 18:42

Another question: I've been thinking about doing a small scale Solvay Process. It consists of bubbling together ammonia and carbon dioxide together in a flask with saturated sodium chloride solution. Then Sodium Bicarbonate should precipitate leaving ammonium chloride in solution, right? Ok, so my doubt is if I bubble CO2 and NH3 together will they form ammonium carbonate?
(Sorry for English mistakes)

Zyklon-A - 29-7-2014 at 05:47

Yes, that's how it works I think. NH₄CO₃ is more soluble than Na₂CO₃, so the sodium salt precipitates.

NH₄CO₃ + NaCl (aq) → Na₂CO₃ ↓ + NH₄Cl (aq)

[Edited on 29-7-2014 by Zyklon-A]

gdflp - 29-7-2014 at 07:34

No, it precipitates sodium bicarbonate before any ammonium carbonate can form. This is due to the ammonium bicarbonate being more soluble than the sodium bicarbonate. Since a saturated solution of sodium chloride is used, the following occurs. NH₃ + CO₂ + H₂O --> NH₄HCO₃. The sodium bicarbonate will immediately precipitate because the solution is saturated with sodium ions. Since there are no bicarbonate ions left in solution, no ammonium carbonate forms. If no sodium chloride is present, then ammonium carbonate will form NH₄HCO₃ + NH₃ --> (NH₄)₂CO₃.

[Edited on 29-7-2014 by gdflp]

bbartlog - 3-8-2014 at 21:51

Are there any stable cuprous salts where the anion has a charge other than -1? For example, cuprous chloride and cuprous acetate exist (though the latter is not trivial to synthesize). On the other hand, cuprous sulfate is nonexistent - it immediately disproportionates in to CuSO4 and Cu(0), and this seems to be kind of typical of such salts. Are all the cuprous salts Cu(X), or does there exist some cuprous Cu₂(X)??

forgottenpassword - 4-8-2014 at 01:04

Cuprous oxide, Cu2O
Cuprous acetylide, Cu2C2
Cuprous sulphide, Cu2S
Cuprous mercuric iodide, Cu2HgI4
Cuprous selenide, Cu2Se
Cuprous sulphite, Cu2SO3
Cuprous telluride, Cu2Te

[Edited on 4-8-2014 by forgottenpassword]

bbartlog - 4-8-2014 at 07:03

Cuprous oxide, obviously ... should have thought of that one! But the others I was unaware of. Thanks!

gdflp - 7-8-2014 at 04:24

I'm going to be running a workshop in about two weeks about creating a low cost dye-sensitized solar cell using titanium dioxide/anthocyanins. The problem is that due to the expense of the ITO glass, only six pieces were ordered and I have to do a dry run right before I actually run the workshop. My question is, if the dry run works, is there any way to clean off the fused titanium dioxide and iodine solution without harming the indium/tin layer on the glass?

Filter Flask Pressure Rating Question

Texium - 7-8-2014 at 20:12

Ok, for some reason I've never been able to find this information despite a lot of various search terms.
What is the pressure that filter flasks (nice Pyrex ones) can normally handle? I've been very cautious with mine so far, but I just want to check and make sure of the safe pressure range because I don't want to go imploding my fancy new flasks on my first few filtrations. (I already uglied up my formerly pristine Büchner funnel with carbon stains, and I don't want to risk worse stuff)

gdflp - 8-8-2014 at 05:10

From your previous picture, I assume you are using a hand pump. In that case or with an aspirator, I would be very surprised if you could get the flask to implode no matter what you did. As long as the funnel is attached correctly, an electric pump can be connected to it with minimal to no risk of anything occurring. The only possible chance you might have is if you stoppered the flask really tightly with the correct size stopper and connected a strong vacuum pump to it and left it for a while.

[Edited on 8-8-2014 by gdflp]

Texium - 8-8-2014 at 07:46

Ok, thanks! I figured that the hand pump probably wouldn't cause any problems, I just wanted to make sure, particularly since I will probably get an electric one some time in the future.

arkoma - 14-8-2014 at 20:49

Electrolysis with Cu anode, lead cathode in a split cell. NaCl electrolyte. Anodic liquid has a fine light green powdered compound in it insoluble in water, vinegar, methanol, and ethyl acetate. Soluble in HCl. What the hell is it? I've got the chloride crystallizing ATM, but stumped on the powder by-product.

Edit--add picture

from left to right Copper Acetate, Copper Acetylsalicylate, Copper Sulfate, and my "unknown"

[Edited on 8-15-2014 by arkoma]

Texium - 14-8-2014 at 21:44

It looks and sounds kinda like copper carbonate, but I'm not sure how that would happen, other than reaction with atmospheric CO₂ at the anode. Did it give off a lot of gas when added to HCl?

arkoma - 14-8-2014 at 22:06

No, not that I noticed, but used VERY small amount of powder and HCl. The chemistry came from one of Ledgard's books--he sez should be copper hydroxide, but he tells fairy tales also. Does LOOK like carbonate.

From Ledgard's "King's Chemistry Survival Guide":

Quote:

The bluish precipitate will be very finely divided cupric hydroxide, which is
very difficult to filter-off.

ganger631 - 15-8-2014 at 13:17

Will neoprene coated latex gloves handle 95-98 conc sulfuric and conc nitric for a few minutes?

Pyro - 15-8-2014 at 13:26

perhaps,

it depends how thick they are.
also, do you mean immersion or just a few spatters?

Zyklon-A - 15-8-2014 at 13:31

Try food-grade gloves. In a different thread I tried tested them, amazingly, they withstand even fuming nitric acid for over 20 minutes of contact.
Of course try it yourself before putting them on to be sure they are the same, but they should work fine.

ganger631 - 15-8-2014 at 17:23

Few spatters

ganger631 - 15-8-2014 at 17:36

Can you give me an example of a food grade gloves?

Oscilllator - 15-8-2014 at 18:31

Quote: Originally posted by ganger631

Can you give me an example of a food grade gloves?

The ones you buy at the supermarket for preparing food...

Pyro - 15-8-2014 at 19:23

they will be fine with a few spatters.

Brain&Force - 16-8-2014 at 13:42

Do you think my Samsung Galaxy S3 can handle being placed in a vacuum dessicator, so I can demonstrate air pressure with the phone's barometer app? The screen can't crack and the phone should be able to remain powered on for the duration of the demo.

forgottenpassword - 21-8-2014 at 02:56

Is there some machine that can find the triple point(s) of an unknown pure substance automatically? Or is there a standard methodology for finding triple points which is independent of the substance's identity?

Acetylene tetrabromide ?

franklyn - 21-8-2014 at 15:30

I'm sure this is a misprint. Ethylene terabromide or perbromide perhaps , or else Acetylene dibromide or perbromide.
wwwapps.tc.gc.ca/saf-sec-sur/3/erg-gmu/erg/guidepage.aspx/guide159/id1668/mnid1807

.

PHILOU Zrealone - 22-8-2014 at 03:09

Quote: Originally posted by franklyn

I'm sure this is a misprint. Ethylene terabromide or perbromide perhaps , or else Acetylene dibromide or perbromide.
wwwapps.tc.gc.ca/saf-sec-sur/3/erg-gmu/erg/guidepage.aspx/guide159/id1668/mnid1807

.

Acethylen tetrabromide...

Old chemical names/technical names looks a bit like alchemistry names and are sometimes misleading.
-You can have the addition product
CH2=CH2 + Br2 --> Br-CH2-CH2-Br (ethylene dibromide = 1,2-dibrom-ethane)
H-C#C-H + 2 Br2 --> Br2CH-CHBr2 (acethylene tetrabromide = 1,1,2,2-tetrabrom-ethane) (*)
-You can have the substitution product
Br2C=CBr2 (ethylene tetrabromide / perbrom-ethylene)
Br-C#C-Br (acethylene dibromide)

(*) I have some litters of that superheavy solvent (almost feels like mercury

), I got those for 1€/L

HeYBrO - 22-8-2014 at 20:46

Just a quick question... I was cleaning glassware with manganese dioxide stains on glassware with dilute hydrogen peroxide and 3-4 drops of conc. sulphuric acid. One of the pipettes i was cleaning had a small amount (like a grain of sand or less) of potassium permanganate which formed some bubbles which i presume was oxygen when it came into contact with the cleaning mix. However, there was a slight sulphurous odour, almost eggy (almost like Sulphur dioxide), but not rancid and did not smell like rotten eggs. I guess it was kind of like the smell of sulphuric acid (sort of sulphurous). Is there any chance i accidentally produced hydrogen sulphide? If so, should i be concerned? I had less than five minutes exposure, only smelt it once or twice and i used small amounts of chemicals (~20-50 ml of 3% hydrogen peroxide , ~50 ml of distilled water and >0.5 ml of conc. Sulphuric acid)
I don't think so but i just wanted to be sure.

EDIT: I don't know why I'm such a hypochondriac once I've been in the lab... I looked up the equation and forgot thats how you titrate hydrogen peroxide.
2 KMnO4 + 3 H2SO4 + 5H2O2 = 2 MnSO4 + K2SO4 + 5O2 + 8H2O

[Edited on 23-8-2014 by HeYBrO]

[Edited on 23-8-2014 by HeYBrO]

[Edited on 23-8-2014 by HeYBrO]

Screen Shot 2014-08-23 at 2.53.37 pm.png - 34kB

Screen Shot 2014-08-23 at 2.53.37 pm.png - 34kB

Texium - 22-8-2014 at 20:54

From the description, it doesn't sound like hydrogen sulfide. If you had produced it, it would definitely smell of rotten eggs. I have produced it in small amounts a few times, once on purpose and a couple times on accident, and the smell is quite potent even at low concentrations. As long as you get away from it quickly and breath some fresh air, you're just fine. Even if you did produce hydrogen sulfide, if you didn't stick your nose right in it, there shouldn't be a problem.

Edit: Ah, I see you answered your own question before I could reply.

[Edited on 8-23-2014 by zts16]

HeYBrO - 22-8-2014 at 22:14

Thanks.

confused - 23-8-2014 at 05:17

well, better to err on the side of caution than to totally disregard something potentially hazzardous
(within reason of course)

solo - 23-8-2014 at 08:55

...when using Oxone in many of its applications, I've often wondered, when they say x amount of oxone, do they mean the oxone as it comes with the two other salts, since the active ingredient is only 48% of the whole, or do they mean of the active ingredient ....it's confusing. some clarification please....solo

DraconicAcid - 23-8-2014 at 14:31

Quote: Originally posted by arkoma

Electrolysis with Cu anode, lead cathode in a split cell. NaCl electrolyte. Anodic liquid has a fine light green powdered compound in it insoluble in water, vinegar, methanol, and ethyl acetate. Soluble in HCl. What the hell is it? I've got the chloride crystallizing ATM, but stumped on the powder by-product.

Edit--add picture

from left to right Copper Acetate, Copper Acetylsalicylate, Copper Sulfate, and my "unknown"

[Edited on 8-15-2014 by arkoma]

That's impure copper(I) chloride.

Amos - 24-8-2014 at 11:45

Any cheap method of separating sodium nitrate and sodium chloride, particularly by solvent extraction?

PHILOU Zrealone - 25-8-2014 at 12:21

Quote: Originally posted by DraconicAcid

Quote: Originally posted by arkoma

That's impure copper(I) chloride.

Or copper hydroxychloride Cu(OH)Cl... I have some 200g in a glass jar and it has that colour, almost unsoluble in water.
It is also refered as copper oxychloride.

[Edited on 25-8-2014 by PHILOU Zrealone]

DraconicAcid - 25-8-2014 at 13:14

Quote: Originally posted by PHILOU Zrealone

Or copper hydroxychloride Cu(OH)Cl... I have some 200g in a glass jar and it has that colour, almost unsoluble in water.

Wouldn't that react with vinegar, though?

PHILOU Zrealone - 26-8-2014 at 11:18

Quote: Originally posted by DraconicAcid

Quote: Originally posted by PHILOU Zrealone

Or copper hydroxychloride Cu(OH)Cl... I have some 200g in a glass jar and it has that colour, almost unsoluble in water.

Wouldn't that react with vinegar, though?

Slowly especially with the 5-10% variant...much better with 80% acetic acid. Household vinegar dissolves metallic copper but you need air (O2) contact with the copper and agitation for the acid to contact those parts and it takes weeks...actually acetic acid dissolves the hydroxy-oxy-carbonate layer of the copper ...once naked it is oxydized again and reforms an hydrated-oxyd layer ...so atomic layer by atomic layer it is dissolved.

HCl conc (30%) on Cu(OH)Cl does the job very wel...you get CuCl2.

I use Cu(OH)Cl as colourizer for green-blue flame in some pyro mixes. It is not very stable with NaClO3 based pyro compositions owing to hygroscopicity and bad interactions between Cu(2+) and ClO3(-)...so not storage advisable for long.

PHILOU Zrealone - 26-8-2014 at 11:23

Quote: Originally posted by No Tears Only Dreams Now

Any cheap method of separating sodium nitrate and sodium chloride, particularly by solvent extraction?

You can play on the difference of solubility of both salts in water as a function of the temperature.
NaCl remains relatively constant while NaNO3 increases exponentially with T.

Haber - 27-8-2014 at 01:06

Quote: Originally posted by Haber

text...

All of the product have now turned into a black dusty powder. Vanillin smell almost completely gone.
Even if it was 5-bromovanillic acid it shouldn't be black.
I'm very curious what has happened. Nobody who have any guess?

I've discovered that this occured because of some kind of reaction with the metal on my thermometer probe.
Would be interesting to know what kind of reaction it was, if anyone have an guess, let me know

gdflp - 4-9-2014 at 18:02

Quick question, white phosphorus won't burn in carbon dioxide right? So if one wanted to convert white phosphorus(I don't have any currently but am thinking of trying to make some using the current method in the phosphorus sticky thread) to red phosphorus, a flask could be charged with some white phosphorus, flushed with CO₂, and heated to convert to red phosphorus without ignition of the phosphorus occurring? If not, what is the best way to convert white phosphorus to red phosphorus?

elementcollector1 - 4-9-2014 at 18:12

Quote: Originally posted by gdflp

Quick question, white phosphorus won't burn in carbon dioxide right? So if one wanted to convert white phosphorus(I don't have any currently but am thinking of trying to make some using the current method in the phosphorus sticky thread) to red phosphorus, a flask could be charged with some white phosphorus, flushed with CO₂, and heated to convert to red phosphorus without ignition of the phosphorus occurring? If not, what is the best way to convert white phosphorus to red phosphorus?

Place in flask with 0.5% wt. I₂, flush with CO₂ or other inert gas, and heat for 24 hours.
At least, that was what was told to me - I've never had the chance to try it, never having had any phosphorus to speak of.

learningChem - 6-9-2014 at 14:58

Is it possible to reduce 3 nitrobenzoic acid to 3 aminobenzoic acid using the Fe/HCl method?

Metacelsus - 6-9-2014 at 16:30

It should work fine.

Texium - 6-9-2014 at 16:38

One thing I've always wondered about, what causes HCl to turn bright yellow when it reacts with concrete? I always refill my small bottle over concrete, and if I spill any it always bubbles up and turns bright yellow.

DraconicAcid - 6-9-2014 at 17:32

Iron impurities, probably.

learningChem - 7-9-2014 at 14:53

Thanks Cheddite Cheese!

Brain&Force - 8-9-2014 at 09:55

What is an ideal container for storing lithium metal? I have some xylene and one of those 8 oz flasks that you can find at a Wal-Mart, but I fear it'll be too difficult to remove the lithium, as the neck is narrow.

Also, would it be a good idea to add a few drops of methanol to clean the metal surface, or is it too reactive for this?

[Edited on 8.9.2014 by Brain&Force]

AlphaDecay - 8-9-2014 at 16:32

I've already asked this before but I really wanna make sure of it. Does mixing ammonium sulfate and potassium nitrate really make ammonium nitrate? Has anyone really tried this reaction? I dont want to any waste chemicals.
Here is my post: http://www.sciencemadness.org/talk/viewthread.php?tid=29592#...
Sorry if I'm asking this again, but I really want to be sure of it.

Metacelsus - 8-9-2014 at 17:36

Yes, but the ammonium nitrate will still have some potassium and sulfate in it.

Oscilllator - 8-9-2014 at 19:12

If you mix them in a solution you wont have either of those things, you will have a mixture of ammonium, potassium, sulfate and nitrate ions in water. If you want to see which one will crystallise out first, I suggest you check the rather comprehensive list on wikipedia, it's quite useful.

Metacelsus - 9-9-2014 at 06:46

Potassium sulfate will crystallize first, but it is still pretty soluble, so you won't get all of it out.

AlphaDecay - 9-9-2014 at 14:02

Ok, thanks. Another question, sorry if it will sound stupid, but, it is better to ask first than regreting after making a mistake. Consider a stand with clamps attached to it. If I'm going to hold an erlenmeyer in the clamp, for example, can I break it by forcing the clamp too much? Or this harldy happens? I am confused, if I don't force it enough, it will fall, if I force too much, will it break? Imagine if it falls during an chemical reaction...

bismuthate - 9-9-2014 at 15:24

Well I doubt it will fall because of the lip on the flask so i don't put it on too tightly. Also what does SWIM mean? I keep on seeing it here.

elementcollector1 - 9-9-2014 at 15:53

SWIM means "someone who isn't me", and is common slang (and a massive red flag) for drug dealers.

Crowfjord - 9-9-2014 at 16:00

My advice would be to clamp tight enough to hold, but without forcing. I've stripped the threads on the screw to one of my good three-fingered clamps this way.

SWIM means "someone who isn't me," and is a long-since abandoned way of avoiding self-incrimination. I think it got started at the Hive? Anyway, it has no place here, but wannabe drug cooks and kewls and whatnot still seem to like to use it.

Whoops, haha, beat to it.

[Edited on 10-9-2014 by Crowfjord]

AlphaDecay - 9-9-2014 at 17:08

Sorry the ignorance but, is this "strip the threads on the screw" an expression?
What does it mean?

Crowfjord - 9-9-2014 at 17:25

Nope, it was meant to be taken literally. The threads are the helical (spiral) part that make a screw what it is. To strip them means to remove or break them off, smoothing the shaft of the screw, thereby making it non-functional as a screw.

AlphaDecay - 9-9-2014 at 17:38

Oh, I see... Also I bought a 1L distillation flask, is it suitable to distill nitric acid from H2SO4 and KNO3? Or it needs to be a complete aparatus?

[Edited on 10-9-2014 by AlphaDecay]

Crowfjord - 9-9-2014 at 18:55

Assuming the flask is really just a flask (not a retort or something similar) then a complete apparatus will be needed. This includes a distillation head, condenser (either a Liebig or West type will do), a take-off adapter (optional, I suppose), and of course some sort of vessel with which to collect distillate.

AlphaDecay - 10-9-2014 at 09:04

Damnit, it is one like this http://m.ebay.com/itm/170984880403?nav=SEARCH ...

Crowfjord - 10-9-2014 at 10:05

Hmm, yeah to use that you will need a non reactive stopper of some sort to plug the top with, and probably a condenser. It depends on how long the side arm is, but if it is as small as the one in that photo, a separate condenser would probably be the safest and most effective way to go.

AlphaDecay - 10-9-2014 at 10:13

In fact the flask is about to arrive. So when it arrives I will post here the lenght of the side arm, then you tell me if its ok to distill without a condenser.

Texium - 10-9-2014 at 14:06

Yeah, you'll definitely need a condenser, but you won't need the distillation head since your flask already has the side arm.
Basically, something like in this picture:

AlphaDecay - 10-9-2014 at 14:34

That makes things much easier, just need to get some money to buy a condenser

...

[Edited on 10-9-2014 by AlphaDecay]

alexleyenda - 10-9-2014 at 18:53

Quote: Originally posted by Brain&Force

What is an ideal container for storing lithium metal? I have some xylene and one of those 8 oz flasks that you can find at a Wal-Mart, but I fear it'll be too difficult to remove the lithium, as the neck is narrow.

Also, would it be a good idea to add a few drops of methanol to clean the metal surface, or is it too reactive for this?

[Edited on 8.9.2014 by Brain&Force]

I used to store my lithium in a mason jar (cheap, airtight, wide neck) and use a relatively rigid wire twisted in the shape of a disk with one end coming out of the oil (to lift it without a mess) to keep the lithium pushed under the oil.

Texium - 13-9-2014 at 09:42

I heated some copper(II) carbonate to make some copper(II) oxide for use in a later experiment, and it turned out unusually brown... Normally I'm used to seeing CuO be solid black, but this particular stuff is very dark brown, like coffee grounds.
Any ideas about what is causing this?
If it helps, I used copper carbonate that I made a long time ago that was cleaned very thoroughly, and I heated it strongly with my Bunsen burner while stirring it until it had all changed color.

Edit: I had 2.00g of CuCO₃ to start with, and now I have 1.70g of CuO, which would be about .005mol of CuO more than there is supposed to be (theoretical yield of CuO should be 1.29g), so I'm thinking that the discoloration is caused by carbonate that didn't decompose. That surprises me, because the first batch of copper carbonate that I made would go black from the heat of a space heater, while this batch can seemingly survive a few minutes of roasting from a Bunsen burner...
Later I'll put it back in a crucible and heat it more, and see if it will look more black and weigh the amount that it's supposed to.

[Edited on 9-13-2014 by zts16]

bbartlog - 13-9-2014 at 17:25

Any possibility that some of your copper is still copper(I)? A bit of red Cu₂O might give you a brown oxide rather than a solid black one...

Texium - 13-9-2014 at 17:41

Quote: Originally posted by bbartlog

Any possibility that some of your copper is still copper(I)? A bit of red Cu₂O might give you a brown oxide rather than a solid black one...

I highly doubt it, as it would have had to have been reduced from copper(II) since I started with CuCO₃, and that would have made it weigh less than it was supposed to since it would contain less oxygen.

Amos - 13-9-2014 at 18:08

@zts16: It has been my experience that large amounts of radiant heat at lower temperatures are more efficient than the smaller and more concentrated flame of a bunsen burner. The best way for me to make completely charcoal black cupric oxide is in a pan on the stove. If you can't use a pan, make a rigid container out of aluminum foil and stick it on there. You could also use the oven, it doesn't require phenomenal temperatures. Bunsen burners for me have typically been terrible at properly heating anything other than the layer of material in direct contact with the walls of my crucible.

[Edited on 9-14-2014 by No Tears Only Dreams Now]

Texium - 13-9-2014 at 18:16

Alright, thanks for the advice, I'll try using my hotplate. It worked (on accident) in the past, but this time I wanted to make sure that it fully decomposed, and my Bunsen burner was already set up for something else. Interesting how that works... you'd think that the burner would work fine for an amount as small as what I was using.

Amos - 13-9-2014 at 18:29

Quote: Originally posted by zts16

Alright, thanks for the advice, I'll try using my hotplate. It worked (on accident) in the past, but this time I wanted to make sure that it fully decomposed, and my Bunsen burner was already set up for something else. Interesting how that works... you'd think that the burner would work fine for an amount as small as what I was using.

I know, it's infuriating for me that it doesn't. Don't forget to stir the cupric oxide often! And maybe wash and filter it when you're done if that doesn't work.

bbartlog - 13-9-2014 at 18:43

Quote: Originally posted by No Tears Only Dreams Now

... cupric oxide is in a pan on the stove. If you can't use a pan, make a rigid container out of aluminum foil and stick it on there.
[Edited on 9-14-2014 by No Tears Only Dreams Now]

This *probably* won't cause a catastrophe, but nonetheless this is bad advice. Don't put copper compounds in an uncoated aluminum container. Odds are good you will reduce the copper and end up with aluminum compounds where once you had a container.

Amos - 13-9-2014 at 18:46

Quote: Originally posted by bbartlog

Quote: Originally posted by No Tears Only Dreams Now

... cupric oxide is in a pan on the stove. If you can't use a pan, make a rigid container out of aluminum foil and stick it on there.
[Edited on 9-14-2014 by No Tears Only Dreams Now]

At 200 degrees C or so, and in the absence of moisture, with an insoluble compound? Aluminothermic reactions need temperatures in the thousands to ignite.

[Edited on 9-14-2014 by No Tears Only Dreams Now]

AlphaDecay - 15-9-2014 at 14:45

About making nitric acid: This guy distilled nitric acid (KNO3 + H2SO4) with a retort http://www.youtube.com/watch?v=TpWwBxsyJok&feature=youtu...
Perhaps a 25cm long side arm distillation flask could do the same?

AlphaDecay - 18-9-2014 at 10:52

Does anyone know the reactions involved in Benedict's test? http://en.wikipedia.org/wiki/Benedict%27s_reagent

gdflp - 18-9-2014 at 11:39

Quoting that wikipedia article, "the cupric ion (complexed with citrate ions) is reduced to cuprous ion by the aldehyde group (which is oxidized), and precipitates as cuprous oxide, Cu2O".

arkoma - 18-9-2014 at 20:06

quick question--aromatic aldehydes generally soluble in MeOH? specifically http://en.wikipedia.org/wiki/3-Acetyl-6-methoxybenzaldehyde

DraconicAcid - 18-9-2014 at 20:39

Quote: Originally posted by arkoma

quick question--aromatic aldehydes generally soluble in MeOH? specifically http://en.wikipedia.org/wiki/3-Acetyl-6-methoxybenzaldehyde

Should be.

TheChemiKid - 21-9-2014 at 14:19

Are there alternatives to ethanol for extracting capsaicin?

gdflp - 21-9-2014 at 14:29

Acetone or methanol should work as well.

arkoma - 21-9-2014 at 14:33

This paper shows decent results with acetone

bismuthate - 22-9-2014 at 12:38

Can compounds, say KNO3, still decompose while in solution? If so it must be harder to do so right?

Brain&Force - 22-9-2014 at 13:02

It depends on the compound, redox potential, and pH, but yes. Are you looking at a specific compound?

bismuthate - 22-9-2014 at 13:39

Ferrous oxalate complexes mostly.

Brain&Force - 23-9-2014 at 11:19

What do you call a solvent that can dissolve polar and nonpolar compounds, such as methanol and ethanol? Is there a special word for them?

DraconicAcid - 23-9-2014 at 11:49

I'd call them polar organic solvents.

Crowfjord - 23-9-2014 at 14:28

@Brain&Force: maybe you're thinking of amphiphilic/amphipathc? These words are synonymous and usually refer to things like surfactants or phospholipids, but they mean "likes both." I'm not sure if they would apply to methanol or ethanol, though; those are usually just considered polar, protic solvents.

AlphaDecay - 24-9-2014 at 13:02

I have found a thread talking about the synthesis of ZnCO3, but I still have doubts about using Na2CO3 or NaHCO3... Has anyone tried this synthesis? http://www.sciencemadness.org/talk/viewthread.php?tid=18167#...

Brain&Force - 25-9-2014 at 11:23

Quote: Originally posted by Crowfjord

@Brain&Force: maybe you're thinking of amphiphilic/amphipathc? These words are synonymous and usually refer to things like surfactants or phospholipids, but they mean "likes both." I'm not sure if they would apply to methanol or ethanol, though; those are usually just considered polar, protic solvents.

Yes, that's what I was talking about. The thing is, on the Sciencemadness Wiki we have a few solvents classed as both polar and nonpolar solvents, which seems contradictory.

gdflp - 26-9-2014 at 05:47

Can you use 10% ammonia when using the hypochlorite ketazine process to make hydrazine sulfate, and just increase the volume proportionately, or do you need to use 20% or higher concentrations of ammonia?

Amos - 26-9-2014 at 06:26

Quote: Originally posted by AlphaDecay

I have found a thread talking about the synthesis of ZnCO3, but I still have doubts about using Na2CO3 or NaHCO3... Has anyone tried this synthesis? http://www.sciencemadness.org/talk/viewthread.php?tid=18167#...

Yeah I have quite a bit of nice flour-fine zinc carbonate(I think it is, anyway) made with just simple baking soda and zinc chloride. It decomposes in a torch flame to the oxide, turning yellow with the heat, so if your end-goal is the oxide, then whatever I have works for it.

[Edited on 9-26-2014 by No Tears Only Dreams Now]

Amos - 26-9-2014 at 06:30

Quote: Originally posted by gdflp

Can you use 10% ammonia when using the hypochlorite ketazine process to make hydrazine sulfate, and just increase the volume proportionately, or do you need to use 20% or higher concentrations of ammonia?

According to NurdRage this works, so if you trust him then go ahead. You'll probably get a lower yield, though.

gdflp - 26-9-2014 at 06:45

When did he say that? I think by "lower concentrations" he meant 20%, though I could be wrong.

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