I’m synthesizing nickel salts for electroplating and ran into an interesting reaction with which I’d like to get more clarity.
When nickel metal is dissolved by HCl we can simply write it as:
Ni(s) + 2 HCl (aq) --> H2(g) + NiCl2(aq).
I’m using nickel metal anode crowns and 31% HCl. Because the crowns are large chunks of metal with such little surface area the reaction takes a
painfully long time, so I’ve been adding some heat and small amounts of 30% hydrogen peroxide to speed things along. I believe the reaction goes as
follows:
1. Hydrogen peroxide is catalyzed by Ni metal to H2O and O2
2H2O2 --> O2(g) + 2H2O
2. The oxygen gas oxidizes the Ni metal
2Ni(s) + O2(g) --> 2NiO(s)
3. The nickel oxide is easier to oxidize than the nickel metal so the reaction speed increases
NiO(s) + 2HCl(aq) --> NiCl2(aq) + H2O
This is where things get complicated though, according to what’s discussed in this thread, 30% H2O2 and 30% HCl react together in an oscillating reaction creating both O2(g) and Cl2(g) and possibly HClO as follows:
H2O2 + 2HCl(aq) --> Cl2(g) + 2H2O
H2O2 + Cl2(g) --> O2(g) + 2HCl(aq)
Then,
H2O2 + Cl2(g) --> 2HClO
HClO + H2O2 --> HCl + H2O + O2(g)
HClO + HCl(aq) --> H2O + Cl2(g)
So as the reaction oscillates, O2 and Cl2 exit the solution while the H2O2 is decomposed. The H2O2 becomes less concentrated within the solution and
the reaction slows again.
*It seems like the reaction is producing Cl2 and/or HClO as I’ve definitely smelled something that reminds me of an indoor pool and is distinctly
different from HCl fumes. I’m keeping the solution covered and outside as I’m more worried about the nickel salts (carcinogenic) and don’t know
if my makeshift fume hood can truly extract all the vapors.
So if all these reactions are present in the solution to some extent, am I getting NiCl2(aq) from ALL of the following pathways?:
Ni(s) + 2 HCl (aq) --> H2(g) + NiCl2(aq)
NiO(s) + 2HCl(aq) --> NiCl2(aq) + H2O
Ni(s) + Cl2(g) --> NiCl2(aq)
2NiO(s) + 2Cl2(g) --> 2NiCl2(aq) + O2(g)
NiO(s) + 2HClO --> NiCl2(aq) + H2O + O2(g)
Unless I’ve missed some glaring detail and I’m completely wrong, this reaction is incredibly deceptive. On the surface it seems so simple and has
few unique products, but it’s really quite complex as there are so many unique routes to get the same products. That’s just so fun and beautiful.
Please let me know if I’ve made any obvious or not so obvious mistakes in understanding this synthesis. I appreciate any insight!bnull - 26-3-2024 at 05:18
You forgot nickel(III).
Welcome to the forum.
[Edited on 26-3-2024 by bnull]SuperNova153 - 26-3-2024 at 20:11
Oh wow, are you referring to nickel(III) oxide, nickel(III) chloride, nickel oxide hydroxide or all of the above? Honestly up until now I didn't know
these compounds existed. It seems nickel(III) oxide has been synthesized (https://www.sciencedirect.com/science/article/abs/pii/S00406...) but remains a bit elusive. I can't find much information about nickel(III)
chloride. Do you think nickel(III) compounds would realistically be involved in this reaction? I'd imagine if they are it'd just be trace amounts. I
doubt nickel oxide hydroxide would be produced in these acidic conditions.
It seems the reactions for nickel(III) oxide would go as follows:
Nickel(III) oxide formation:
3H2O2 + 2Ni(s) --> Ni2O3(s) + 3H2O
or
4Ni + 3O2 --> 2Ni2O3
Pathways in which it can produce NiCl2
Ni2O3(s) + 6HCl(aq) --> 2NiCl2(aq) + 3H2O + Cl2(g)
2Ni2O3(s) + 4Cl2(g) --> 4NiCl2(aq) + 3O2(g)
2Ni2O3(s) + 8HCl(aq) --> NiCl2(aq) + 4H2O + O2(g)
I'm also unaware if any of these reactions are thermodynamically favorable.
Thanks for your feedback! This reaction (and similar metallic dissolutions) just keeps getting more and more interesting. bnull - 27-3-2024 at 04:35
Do you think nickel(III) compounds would realistically be involved in this reaction?
No. I can't affirm it but I can't discard it tooNote. Maybe trace amounts of nickel(III) close to the surface of the metal are involved in
the dissolution. Like this:
$$2Ni^{2+}(aq)+H_2O_2(aq)+2H^+(aq) => 2Ni^{3+}(aq)+2H_2O(l),$$
$$2Ni^{3+}(aq)+Ni(s) => 3Ni^{2+}(aq).$$
I haven't seen any paper about it, though.
Note: Because things always get weird with transition metals. Who knows.clearly_not_atara - 17-4-2024 at 18:31
Quote:
Hydrogen peroxide is catalyzed by Ni metal to H2O and O2
Zero points.
You have to consider how this happens. It's probably more like this:
Ni(OH)2 + H2O2 <> Ni(OH)OOH + H2O
Ni(OH)OOH <> Ni(O)OH + HO•
HO• + H2O2 >> H2O + HOO•
HOO• + Ni(O)OH >> O2 + Ni(OH)2
But hydrogen peroxide is a perfectly good oxidant on its own. It's probably just oxidizing the nickel directly. In the absence of HCl, this would lead
to passivation. So together they erode the nickel. bnull - 18-4-2024 at 02:38
The paper was published in Corrosion Science (https://www.sciencedirect.com/journal/corrosion-science), which is quite interesting. There are so many journals that we end up searching not only
the wrong keywords but in the wrong place also.
And there are nickel(III) ions, but not in the way I thought.Ghrrum - 23-4-2024 at 09:43
While I don't have anything to add regarding the synthesis question I do have experience making stuff for electroplating and wanted to chime in a
straightforward source for Nickel Oxide:
I've found pottery glazes are pretty high in purity and it worked well for me, but that was also 10 years ago so your milage may vary. SuperNova153 - 15-5-2024 at 08:24
Very interesting. I was unaware that passivation played a role here as well. Thanks for the article bnull, that was very helpful.
Hydrogen peroxide is catalyzed by Ni metal to H2O and O2
Ni(OH)2 + H2O2 <> Ni(OH)OOH + H2O
Ni(OH)OOH <> Ni(O)OH + HO•
HO• + H2O2 >> H2O + HOO•
HOO• + Ni(O)OH >> O2 + Ni(OH)2
Thank you for the mechanism, but I'm still a bit confused as to where the nickel hydroxide came from initially.
Reviewing the paper bnull shared it's likely from the H2O2 in solution. The paper states "hydrogen peroxide undergoes a heterogeneous catalytic
process decomposing to water and oxygen via the generation of a surface silver hydroxide...The dissolution of Ni occurs via peroxide-driven formation
and oxidative dissolution of surface Ni-OH groups"
Ni + H2O2 --> NiOH(ads) + OH•
or
2Ni + H2O2 --> 2NiOH(ads)
Is there something about the NiOH being adsorbed to the metal surface that allows it to exist in the +1 oxidation state? Or is there another reaction
that may take place to convert the adsorbed NiOH to Ni(OH)2?
This reaction seems to get more and more complicated
