xeluc

Alright.... Wel thanks. Your results differ vastly from mine. I will look for a more pure source of Iron.... AS a side note. I have a large quantity of CuCl2 solution evaporating and I see some beautiful crystals. I could never get the CuCl2 pure, it was always very green. I found out that the reason it was so impure was I was boiling it. I let my CuCl crystalize over the course of a few weeks and its hardly green at all. It looks about the same color as your reagent grade. Like an aquamarine color maybe... well bluer than that. Just an update...

also. I told you guys no precipitate was formed when putting HCl in the NaOH mystery liquid solution.. well i cam home today to find white crystals... Some are cubic and some are stick like... lol I may never know what that is...

What WOULD be a pure fom of iron I could buy without spending lots.. Is steel wool decent enough it it is washed thoroughly?

not there`s much anyone CAN do wrong with a 2 ingredient experiment

ha.. Yeah I replicated your dark yellow solution with the steel wool but the scratch pad thing is being weird. I set some of THAT green solution out and it's yet to turn yellow. Oh yeah.. Let me add some HCl... lol Well, We'll see what happens but I still dont know what the white floaty prec. is! I dried some out and it dried off white. It disolves in NaOH but not HCl, but looking at some solubility rules; There isnt a compound that wont disolve in HCL but will in NaOH except I believe Thallium. This leads me to believe that Adding this white compound to NaOH does not change the composition of the mystery compound, rather that it only disolves under alkaline conditions. Either that or not , heh. Also I might add, adding HCl tothe disolved material does not yeild a precipitate. Therefor the beginning white compound was not a chloride.. What other anion could it be?

Of course but YT said there was no visible reaction, where there should STILL be bubbles of H2... That was my experience anyhow

The results of my experiment corrolated with what I found in the Iron Chloride solution. The NaCl disolved very Poorly in HCl as apposed to H20. So basicly, the HCl is saturated with Fe2+ and Cl- very rapidly, but solubility is aweful because of the amount of chloride ions already in solution as HCl. I noticed ungodly amounts of HCl vapor eminating from my reaction vessle. Much more than would be expected in fuming HCl by it's self. I hypothesize that when the point in time comes when the water is saturated in chloride ions and FeCl2 begins to precipitate, that HCl is also forced out of the solution proportionatly. This is of course an educated guess. The liberated Hydrogen could give me false judgment, but these fumes weren't clear like H2, they were white and smelled horribly of HCl. I'm going by the strenth of odor here... Anyhow, learned something ese new.

Next. I have no idea what you guys are talking about with your Iron not disolving at any appreciable rate. I figured steel wool wasnt tremendously pure, so i went and bought some of those steel scratch pads made from the coiled steel. Pop one of those in HCl and it bubbles very vigorously. (31% HCl).

In my last experiment, I added a little steel wool at a time and alowed it to Oxidize to FeCl4- each time. This time, I threw a large amount of iron (Which STILL all disolved) in HCl and let it react away. Some neat has happened. I see 3 distinct layers. One layer of dark green precipitate (FeCl2?). Next, a saturated solution of (FeCl2?), and on top, a lighter green precipitate (Hydroxide? something else?) The layer of precipitate on top of the water is NOT because of hydrogen sticking to it and causing it to float. I have shaken the mixture many times and within 10 seconds a break occurs between the two precipitates. Also, the neat thing is, this top layer of precipitate is capping the liquid and therefor stopping it form oxidizing. The residue on the bottle however is yellow from oxidation as there was nothing preotecting it. Could someone verify what this top precipitate is? I havn't isolated it yet, it could be white for all i know and the solution is staining it. If this was Fe(OH2), I would expect it to sink in water...

EDIT: Ok guys, I decided it was lazy of me to ask what this was when i had done no tests myself. So I set out to isolate the top prec. I put some in excess water in a test tube to wash out any FeCl2 on it. I shook it up and YES, a snow white precipitate floated up very rapidly. So, this compound is white. But what is it? In a whiel I will see if it was oxidized any, but I don't think it will, as its been sitting around a while already.

Also, I believe that Yt's iron was actualy Fe3O4, and that caused it to react slower. I've presented this before but he insists on "Metal" being deposited on the magnet, but if I remember correctly, Fe3O4 Is attracted to magnets also and I know for a fact it is abundant in beaches....

Just reread Yt's first post. I also had a hark time with Screws, but I did see bubbling. If the scres were coated in soemthing that may retard (heh) the reaction.

Uh oh.. I just re read the whoel thread and Woelen says Fe2+ is almost colorless, but I have a VERY dark green solution. I can't even get light through it very well.. Maybe my Iron was alloyed and the white Prec. as well as the dark green color was form a contaminent (I hope not.)

WOW! I filtered my FeCl4- Solution jsutn ow. I'm left with a very clear dark yellow liquid. To my surprise, I saw that the precipitate I had seen included some crystals of FeCl2. At first I didnt understand how YT had FeCl2 in his solution of FeCl4- solution. I'm not positive how this happened though. Well. Maybe the solution was saturated in FeCl4-, But there was so much HCl left that it continued to Oxideze Iron, resulting in a precipitate of FeCl2. Could someone second this or tell me im wrong?

well, looks like we figured that out

Yeah.. I know this much so far.. I found myself that the situation was very similar to my last one. my question was, how to crystalize the Iron Chloride? Also, you said that the precipitate I have in there was formed because of the Iron III correct? If so, then reducing the solution and decanting followed by your perging of oxygen in the container should get rid of the prec. right?

Hey, sorry to resurrect this board, but it made more sense to add on here then make a new thread. I just decided that it'd be neat to make some Iron Chloride. I took steel wool (I don't have so many reagent grade chems) and disolved it in hydrochloric acid. Came back and it was yellow. Added more steel wool, it turned light green with white prec. Just like YT. I'm just trying to figure out how to actually get FeCl2 when my aolutions oxidize on me. Woelen said that the Oxidation takes weeks but within 4 hours the steel wool dissapeared and the solution was bright yellow. (Also @ YT: If you take care to clean off any oils put on the steel wool to keep it from oxidizing, it reacts very decently (meaning you can actually see hydrogen bubbles. At first nothing happens as a layer of oil,Fe2O3, or w/e is disolved)) So my question is, with my solution oxidizing all the time, how do I create FeCl2 instead of FeCl3?

EDIT: Also YT, did you say that you got FeCl2 from the Ocean? I beleive you made a typo but your message says that at some point. I think you got Fe3O4...

If you get two cups, fill each with KNO3 solution, get some yarn (maybe 3-4 strings wide) and wet it in the solution and drape it between the cups so that the liquid make a "connection" between the cups. This is basicly a Salt bridge. What this does is keeps contaminates form the Cathode (In this case, the Tetramine Copper complex solution) from getting into your Anode solution, which will be your CuO (I Think. Woelen said the Cl-'s from the NaCl solution favor Cu+ Ions, so with KNO3, you should get CuO?)

lol. Nah I dont mean heating I mean letting it warm to room temp' date='

Not that it matters now,

Yeah I had the copper probes in the water all the time.

I did get to a point with one sample of black crap plating out, lol.

Thats whne I knew the soln was full. [/quote']

 

Even letting it go to room temp wont change anything to Copper Hydroxide...

 

 

@Woelen: Actually, it's not easily undertood by me at least why Copper (I) Compounds ar/ efavored over Copper (II) in a NaCl solution. I understand that the Cu+ is stabailized by teh Chlorine Ligands but Cu2+ (obviously) does the same. Does voltage have anything to do with the preference? Maybe you could elaborate on this.. Thanks

I know your gunk contains no nitrated because all nitrates are water soluble.

With the purity of the gunk, it depends. I have gottne pure Iron Oxide from electrolyzing Iron in NaCl and I have gotten pure Cu2O from electrolyzing Cu in NaCl, but when other combinations of solution and electrode are used, different chemicals ar ecreated. For instance. Electrolysis of Cu in NaCl produces (At least for me) Cu2O. Differing voltages may have differing effects. But when you electrolyze in a solution of Copper Chloride, Cu(OH)2 is formed. So there, Hope I helped. If you have any more questions, go for it. And welcome to SFN!

Also, if you want only H2 and O2 given off, Try a solution of NaOH. Using copper Electrodes is fine then, the solution stays clear with no gunk forming.

Ok, lets disect this message and see what we can do. I love Copper, you can do so much stuff with it and it's really cheap. Anhow...

 

You know the story' date='

You get (blue) copper hydroxide and (brown) copper oxide.

The copper hydroxide will decomposses if over a certain temperature.

[/quote']

 

For SOME reason, I get only Cu2O, maybe Woelen can save the day and explain our differing outcomes. I use a high potential (12 Volts) so maybe that makes a difference. Anyhow, this is irrelevant. You are correct, Copper Hydroxide will dehydrate to form CuO. But watch, here it looks like you understand that the Copper Hydroxide is a precipitate and therefor insoluble. Keep this in mind, we will go back to it later.

 

So I thought lets cool it and try and make just copper hydroxide.

So I Let the salty water sit in the freezer until nice and cool and then placed in a ice bath just to make sure.

 

 

Although you are correct in saying that heat will decompose the Cu(OH)2' date=' The Copper oxide formed during electrolysis may or may not be from heat. Voltage may also contribute. This is just something to think about, I go in work in an hour and can't test this, but if noone has by tomorrow I will.

But I didnt get copper hydroxide.

I got a green/brown precipitate.

 

No idea here. Are you sure only copper metal is touching the solution?

 

I'm sure I made copper chloride.

I tested this by taking a sample of the sludge and adding Sodium hydroxide.

I got copper oxide precipitate and copper hydroxide soln

 

WOAH' date=' stop right there. remember that thought I told you to hold onto on the first quote? Well right here, you forget that Copper Hydroxide is insoluble in water. It is possible however that you had some Cu2+ ions in solution which would crystalize CuCl2. I have done this myself actually. So after your test, you should have blue Copper hydroxide and copper Oxide maybe. If you SOLUTION is blue, then You have something else. Woelen told me that you would get a royal blue solution of the CuO2(2-) complex.

I did the same with my sample of copper oxide as a control test, No reaction.

I also notice if I dont descent a sample it will (as it warms) turn into copper hydroxide(which is should anyway).

 

The first part of this is absolutly right. Next you say that warming your mixture turns it into Copper Hydroxide? This obviously can't be true since you said already that heating copper hydroxide yeilds Copper oxide. This isn't a reversable reaction with heat or lack of alone.

 

 

 

 

So anyway, Electro chemistry is a finicky thing. In truth, when you first start out, you spend more time figuring what went wrong than conducting something that you can successfully predict. I hope I've shed some light on this, but I don't have a lot of time to experiemtn right now. So, I'll do that tonight or tomorrow. Hope I helped a little

i remember someone in my chemistry class spilt some phenolphthalien on his hand and didn't notice. he then wiped his mouth with his hand and ingested some of the phenolphthalien and he didn't even make it to the toilet. bad for him funny as hell for the rest of us. poor guy. i think he works in a mcdonalds now.

I work at Mcdonalds. RRRR I mean I did yesterday I work at Dairy queen now. HA. Anyhow, your saying that he licked his hand and spontaniously pooed?....

Hello.

I think you will like this site. I have it on my favorites:

Solubility Rules

if your going with a peroxide, use H2O2. So much cheaper... I don't even know what benzoyl peroxide would do, but H2O2 will work for you plus it's unlikely to do irreversable damage to your shoes. If you look at a sporting goods store or some large shoe stores in malls, they sell little applicators specifically for whitening shoes. I can vouch for their effectivness also. Of course it is somewhat more satisfying when you use rough chemistry to acheive things ;-)

I never could understand why YT went through all the trouble to make HCl... He used Sulphric acid to make it and where I live at least, sulphuric acid is over 2X as expensive as HCl. Maybe he just doesn't have any hardware stores around...

(wee post #100 )

ha, I don't mess around. Full strength. Ohwell, she was obsessive and naggy anyhow

CuCl2 jsut became a little less scarce instead of opening my 20 ounce bottle and letting HCl fumes cloud up my room, I filled a prescription canister half full with my stuf and threw in excess copper wire. in a day or so I should have a very concentrated CuCl2 solution.. yay.. CuCl2 is very special to me, it is the very first inorganic compound I made myself not in Chem class. Also, of course you were the one to help me Woelen, on chemical forums. I need to buy some copper sulphate. I heard it's pretty easy to find in stores but i havn't had any luck..

Ok guys. Once apon a time My girlfriend (ex now, ha) bought me a hoodie. Now her beign a girl for some reason bought me a medium.. I like large, im not gonna feel bad and think im fat if sheda got me large.. sigh.. That's a womens thinking for ya. Anyhow. I got pen on it. For some reason nothing would take this pen out, so i dabbed on some bleach. Instantly the pen was gone, but the blue still stayed on the hoodie. I was very happy. So I quickly washed out hoodie. Days later, there was a hole replacingthe pen mark... lol. In short. Yes, bleach deteriorates organics.. But slowly. If you bleach your shoes.. and soak them in water and scrub afterwards, your cool. But if mister ty-shirt dude didnt wash off his bleach good enough, sure, the shirt could have been ruined.. Just my 2 cents..

Woelen, you never mentioned what happened to the Na+ in that reaction, or is it unknown or did i just miss it?

No, I'm not Woelen, but basicly the Na+ is a spectator Ion. Woelen mixes NaCl and CuSO4 together because he needs Cu2+ Ions and Cl- Ions. The Na and SO4 Ions don't do much of anything. This is talking about his experiment though. Click on his link to view. When he said to mix CuCl2 and NaCl, I am pretty sure that was jsut a mistake, as CuCl2 already possesses the neede Chloride Ions, so adding NaCl will have no effect on the outcome of the reaction.

I still have my 20 Ounce bottle cookin up Cu Ions. So I'll have some soon.. or not. I have been doign this for like 4 days. I know that I have more Cu ions in solution, as I have to keep putting more wire in, but when I take off hte cap, the acid STILL fumes, indicating that its still very strong. I just can't find a good way to get the Cu+ Ions to oxidize. Putting new oxygen into the bottle works, but not very fast. H2O2 would work, but the lowsy 3% I have will dilute the acid pretty fast. So I don't know. Anyhow, I will try this experiemtn when I get a chance. I also found out that NaOH in solution will make H2 indefinatly as long as you have water and aluminum, but maybe the CuCl2/NaCl (Do I need both when I'm using CuCl2? (The CuCl2 has chlorine already...)) solution will work faster, although I'm not sure if I need that. The reaction with the NaOH solution is so exothermic that I had to put the bottle down because it got too hot. Soon after, the bottle actually shrunk, which was a cool thing to see.

if you MUST use copper wire to attatch to the nail' date=' then dip that end in wax, it`ll provide some protection against splashes when the soln fizzes, bubble gum works just as well

 

if you use Steel nails and do it properly you should get a tiny amount of a Black powdery PPT also, that`s actualy the Carbon coming out of the steel! (perfectly harmless to the reaction though)[/quote']

 

Maybe it should be stressed that even when this is done, you should still try your best to keep the copper out of solution. Even if it is protected; If everything is in the water, the metal underneath the wax could corrode, breaking your seal. YT hinted at this and Woelen went and said it, heh..

It's not wasting if you get something done with it that nothing else could achieve. Besides. HCl is like 4-5 dollars a gallon at the hardware store. and all you'd need is like, a capful or two.

The oxide layer protects the pure aluminum beneath from anything except for things powerful enough to disolve the oxide, lol. The oxide layer will protect from things such as Oxygen and water, but other things like HCl and NaOH are too pwerful for it.

Conc. HCl will get most salts off glassware...