Author Topic: Electrochemical hydrogenation (Read 325 times)

2bfrank · « **on:** May 25, 2009, 09:37:08 AM »

I think I got this via SM, I think, but it don't matter. I haven't had the time to run this experiment, but it appears so dam convenient - by the look of this report. The nickel, could easily be plated on another OTC metal, the nickel salt purchased from a pottery supply place. I have plated nickel on aluminum before and it is somewhat easy. This would serve as the sacrificial anode, and purely a rod of iron as the cathode. It looks very easy and could reduce a variety of compounds. I know the article is small scale, but I am sure if warranted, it could easily be scaled up. I thought it was at arms reach to the home chemist hence this inclusion.

Vesp · « **Reply #1 on:** May 25, 2009, 08:43:22 PM »

That is very interesting and it does look useful. It seems like the Ni electrode could be a bit annoying to obtain, especially if you have to make it. I wonder if Nickels would have any use? They are composed of 25% Nickel, and the rest is copper. If the copper didn't ruin the wanted properties of the nickel, maybe they could find some use in this - though I doubt it.
Welding rods are another possibility, as I think they are sometimes nickel.

2bfrank · « **Reply #2 on:** May 26, 2009, 12:01:18 AM »

Hi Vesp,
One can buy nickel foil, but it is expensive. (AUS) via a ripoff distributor, 30-40 or so,(cannot quite remember) bucks for 25g
I think the go would be to buy elcheapo nickel salt from a pottery supply place, convert that to Nickel chloride, and use that in a simple electrolysis set up to plate the nickel on say aluminum, or some other metal. If there is interest, or a need for more detail, and I get through the next few weeks of exams(currently studying) I will share what I have learnt on this subject. As far as using a nickel alloy, I am not that up with electrolysis, but it might work fine. Nickel is a good metal to have, so I reckon getting some via a pottery supply place is a bonus, as one can do a shitload of things with it.

2b

Douchermann · « **Reply #3 on:** May 26, 2009, 04:52:59 PM »

You can buy high nickel welding rods as well. They're not the cheapest thing in the world, but certainly not as bad as the nickel foil mentioned in this thread earlier. The only problem is you can't go to a typical hardware store to find them, you have to go to a specialty hardware store, or a welding store.

Vesp · « **Reply #4 on:** June 08, 2009, 10:44:46 PM »

the electronegativity of copper is 1.9 while the elctronegativity of Nickel is 1.91.

is a 0.01 difference to little to separate them via electrolysis? They are to damn hard to separate chemically, though when I put in Iron it gets a copper colored coating on it, instead of the nickel (currency) color I'd assume it would be if it were alloying.

Perhaps adding the correct amount of Fe powder to the 75/25 mixture would be do-able getting out 75% of the copper first, leaving you with a iron nickel solution.

Then you'd be able to later add more Fe powder to extract the Ni, leaving you with a Iron solution of whatever acid you used. As well as some Copper powder.

Not to sure with that small of electronegative difference if the metals would just plate on depending on who was the closest.

I'd like to try some hydrogenations but I've got to get the nickel first.

Sedit · « **Reply #5 on:** June 08, 2009, 11:59:33 PM »

This is simular to what I was just talking about in the other threed.

I got an idea now though.

Nickle is magnetic... Copper is not. Precipitate the metal powders as the oxides reduce them with carbon or whatever means and then mix them with a magnetic stirrer. I would think that the Ni should coat the stir bar if you have a magnetic one and could just be pulled right out of the copper.
It sucks that I just dumped all my precipitate to make room because I didn't think of any use for it before. I did save some of the oxide mix converted it to the chloride and soaked some crushed unglazed ceramic in it. Let it dry and put it in the tube furnace to convert back to the oxide in a very fine form. I should reduce it now just to see what color it forms. I want to know if it will be silvery or copper color. Im going to use this as the catalyst for my tube furnace so next time I fire it up ill just pass some methanol thru it to reduce the oxides.

Vesp · « **Reply #6 on:** June 09, 2009, 12:28:01 AM »

The Ni powder is very fine and intimately mixed with Cu powder which is also very fine. I tried this with various magnets. Some very weak and some very strong. I could only collect a little bit of Ni regardless and it wasn't pure. It caputured Cu powder, and was a bitch. The weak magnet seemed to get more pure Ni, but gathered much less. The stronger magnet trapped more Cu, but got more powder in total.

I didn't have much luck with using a magnet, but I tried.

I think maybe if I got one of my recently acquired neodymium magnets that is much stronger then the ones I used previously, and shook a water metal powder mixture, it might work a lot better then what I was doing before. I'll consider trying this with my messed up Cu/Ni mixture.

2bfrank · « **Reply #7 on:** July 24, 2009, 08:55:19 AM »

what about in solution, where the nickel may collect on a stir bar. One would have to stop, collect and repeat perhaps numerous times. I ve not tested this, as I have not ever used a stir bar with metals that could be magnetized-etc. If its dry the metal will drag another metal with it, which sounds like has been distribed. .

Sedit · « **Reply #8 on:** July 24, 2009, 02:42:17 PM »

I managed to reduce the metals with aluminum to a black powder that sticks to magnet but a flame test made obvious that there was large amounts of Cu also contained in it.

Here is a quote from Not_important

Quote

Check the properties of NiSO4 and CuSO4, I believe copper sulfate decomposes around 650 C while nickel lasts until 840 C or thereabouts. Heating the mixture to 700 - 730 C should decompose the copper salt to the oxide, leaving the nickel salt. Having not tried this, I can't say that the mixture might behave differently that the isolated salts, or that you might need to leech out the NiSO4, dry, and heat once again.

If you overheat the NiSO4 also decomposes into the oxide. The mixed oxides can be extracted with hot dilute H@SO4, with dissolves the copper oxide. I don't know what 'dilute' is exactly, consult a text on nickel mining.

Other wet methods include precipitation of CuS from an acid solution, precipitation of cuprous iodide and iodine by adding a soluble iodide (only for countries where the element iodine and its compounds have not been outlawed), using dimethylgloxime to precipitate the nickel (only for people with deep pockets).

I would have tried starting with the metal alloy, and electro-refine it by plating copper from the alloy to a copper electrode in an acid solution, using a low voltage. The nickel would go into solution as the copper dissolves to be plated out on the other electrode.

Another method, when you have a mix that is mostly nickel, the finely divided metals (from reducing the oxides) can be treated by the Mond process to extract the nickel. As this involves carbon monoxide (toxic) and nickel carbonyl (more toxic) it is not a process for the amateur or to do at home, unless your home is very isolated and your will is up to date. But it is a elegant method.

Don't even think about the last method as you may poison yourself just thinking about it.

Also take a look at this for some ideas: http://www.sciencemadness.org/talk/viewthread.php?tid=534&page=2#pid129332

Douchermann · « **Reply #9 on:** July 24, 2009, 04:28:50 PM »

Lol yeah, I told Vesp the magnet trick wouldn't work about a year ago when he first wanted to separate the two metals. Produce the nitrate of both metals. Evap all acid and water, then dissolve in a dilute mineral acid solution. Bubble H2S into this solution, and the copper sulfide will ppt, however the nickel will either not form the sulfide, or not ppt. There is a special name for using the transitional metals valence electrons to your advantage in scenarios like this.

Vesp · « **Reply #10 on:** July 24, 2009, 09:18:37 PM »

Quote from: Douchermann on July 24, 2009, 04:28:50 PM

Lol yeah, I told Vesp the magnet trick wouldn't work about a year ago when he first wanted to separate the two metals. Produce the nitrate of both metals. Evap all acid and water, then dissolve in a dilute mineral acid solution. Bubble H2S into this solution, and the copper sulfide will ppt, however the nickel will either not form the sulfide, or not ppt. There is a special name for using the transitional metals valence electrons to your advantage in scenarios like this.

I tried this and it failed as well. The magnet method works only sorta-ish
It concentrates the Ni, but it doesn't make it completely free of Cu.

How much would Cu even interfere if you are using it for a catalyst?

Sydenhams chorea · « **Reply #11 on:** May 28, 2010, 12:13:59 PM »

Quote from: Vesp on May 25, 2009, 08:43:22 PM

That is very interesting and it does look useful. It seems like the Ni electrode could be a bit annoying to obtain, especially if you have to make it. I wonder if Nickels would have any use? They are composed of 25% Nickel, and the rest is copper. If the copper didn't ruin the wanted properties of the nickel, maybe they could find some use in this - though I doubt it.
Welding rods are another possibility, as I think they are sometimes nickel.

The Canadian nickel coins from the following periods were composed of 99,9% nickel :
1922-1942, 1946-1951, 1955-1981

They each individually weigh around 4.5 grams.

But I think nickelplating an iron rod would be a good idea as well.

Vesp · « **Reply #12 on:** May 28, 2010, 09:51:42 PM »

Rumour has it that Ni and Cu are easily separated by the addition of an ammonia solution to some of their soluble salts.. This makes U.S nickels useful.

Allowing for the Ni(OH)2 or whatever it maybe to precipitate, then would allow for it to be redissolved easily, now being a pure solution of Ni+2 and electroplated on something such as graphite ought to do the trick.

Sedit · « **Reply #13 on:** May 29, 2010, 03:06:27 AM »

This is no rumor I just havent provided a step by step for yall folks yet. You have seen the pictures of the Ni oxides(hydroxides?) over at Science madness and just trust me the Chlorides are pretty damn pure just a bitch to crystalize because they appear hygroscopic. But the presence of Cu has not been noted yet and will be neglectable if present at all.

Its just matter of working the quantities of Ammonia hydroxide solution to selectvly precpitate the Ni oxides and make sure I get the most out of it. I have about 80 gram of Nickles dissolving in H₂SO₄ right now but its a slow process to say the lest. Need some heat I think to speed the process along is all but when its done you will see pure NiCl₂ from US nickles.

[Edit]

Ill add some pictures to show. I feel some data on a nickle ammonia complex spoken of in literature according to my research is FALSE look. The blue is Cu+Ni in NH₃ solution royal blue from the fomation of copperamine complex. The clear one with green precipitate is also Ammonia solution with NiO in it but notice even with all the nickle present its still not blue as literature suggest it should be.

lugh · « **Reply #14 on:** May 29, 2010, 04:49:22 PM »

The Chemical Reviews 62.19-40 (1962) article that was scanned back in the days of the Hive:

Sydenhams chorea · « **Reply #15 on:** May 29, 2010, 07:54:32 PM »

Just to clear up matters for further discussion and thus to avoid confusion, I'd like to state here again the difference between electroreduction and electrochemical hydrogenation.

Electroreduction takes advantage of the high overvoltage of some metals, like Hg, Pb, and Cd, favoring the addition of monoatomic hydrogen to the reducible substrate instead of diatomic hydrogen gas production. It is similar in that way to the mechanism of the Al/Hg reduction.

Electrochemical hydrogenation is just catalytic hydrogenation, using metals like Pd, Pt and Ni who have low overvoltage (easy H2 evolution at the cathode) but the capacity to adsorb monoatomic hydrogen on the surface (or form intermediate hydrides).

Sydenhams chorea · « **Reply #16 on:** May 29, 2010, 08:01:33 PM »

Quote

I think the go would be to buy elcheapo nickel salt from a pottery supply place, convert that to Nickel chloride, and use that in a simple electrolysis set up to plate the nickel on say aluminum

and have chlorine evolution at the anode? No thanks! Use nickel sulfate or acetate.

Speaking of the acetate... anyone ever tried dissolving Ni in acetic acid to which a little 30% H2O2 has been added? This trick works well for producing acetate salts from various other metals. Just beware, the reaction is very vigorous.

and why electroplate on aluminium? Al will already displace Ni from its solution, you will have massive Ni particle disposition... Better you use a steel rod I think?

2bfrank · « **Reply #17 on:** May 29, 2010, 08:18:38 PM »

Thanks, and quite right, why gas yourself when you dont have too,... and yes he has used H2O2 to kick things along, and somewhat cautiously.

Sydenhams chorea · « **Reply #18 on:** May 29, 2010, 08:22:07 PM »

Quote from: ziggy on May 29, 2010, 08:18:38 PM

Thanks, and quite right, why gas yourself when you dont have too,... and yes he has used H2O2 to kick things along, and somewhat cautiously.

You have made the acetate succesfully from Ni and acetic acid/H2O2?

2bfrank · « **Reply #19 on:** May 29, 2010, 08:26:00 PM »

No I dont recall specifically Nickel acetate, I think Ive added it to Iron salts.

Author Topic: Electrochemical hydrogenation (Read 325 times)

2bfrank

Electrochemical hydrogenation

Vesp

Re: Electrochemical hydrogenation

2bfrank

Re: Electrochemical hydrogenation

Douchermann

Re: Electrochemical hydrogenation

Vesp

Re: Electrochemical hydrogenation

Sedit

Re: Electrochemical hydrogenation

Vesp

Re: Electrochemical hydrogenation

2bfrank

Re: Electrochemical hydrogenation

Sedit

Re: Electrochemical hydrogenation

Douchermann

Re: Electrochemical hydrogenation

Vesp

Re: Electrochemical hydrogenation

Sydenhams chorea

Re: Electrochemical hydrogenation

Vesp

Re: Electrochemical hydrogenation

Sedit

Re: Electrochemical hydrogenation

lugh

Re: Electrochemical hydrogenation

Sydenhams chorea

Re: Electrochemical hydrogenation

Sydenhams chorea

Re: Electrochemical hydrogenation

2bfrank

Re: Electrochemical hydrogenation

Sydenhams chorea

Re: Electrochemical hydrogenation

2bfrank

Re: Electrochemical hydrogenation