Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Mystery metal?
I'm sure that there is probably a stupidly simple solution to this that I'm missing, but I figured I'd post it just in case.
I have a metal, which I originally assumed to be chromium as it was plated on some kind of broken farm implement handle that I found. It's silvery in
color, and very lustrous.
I tested it by putting it in some HCl. I left it in for about 10 minutes, and it didn't react at all, when chromium would have.
At this point, I thought it might possibly be silver, despite that that wouldn't really make any sense based on its use. I then prepared some
H2S and left the metal in the H2S atmosphere with moisture. The metal did not change in appearance after another 10 minutes,
when silver would have tarnished.
I seriously doubt that they make platinum plated farm implements, so does anybody have any idea what I might have done wrong, or any other ideas as to
what it might be?
|
|
|
DraconicAcid
International Hazard
Posts: 4355
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I'm not sure how readily chromium reacts with hydrochloric acid (Cotton and Wilkinson say the metal is resistant to most corrosive agents, but then
adds that it does dissolve in hydrochloric acid). Maybe leave it longer? Or add some copper(II) chloride- that seems to catalyze the reaction
between HCl and zinc, aluminum, etc.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
I've now left a couple flakes of it in hydrochloric acid for over an hour, still nothing.
I'll try that with the copper chloride though. I happen to have some on hand.
Also, there appears to be a thinner layer of copper plating underneath this metal.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Chromium coatings should react quite vigorously with HCl, going by my experience with lump Cr metal.
Is it possible that the implement is nickel coated? Nickel does resist most acids (and bases), except HNO3, very well. Nickel coated steel is quite
common in lab equipment for instance. It's easy (by its luster) to mistake it for chromium, although the latter can be polished to a higher shine, I
think...
[Edited on 10-3-2014 by blogfast25]
|
|
|
phlogiston
International Hazard
Posts: 1379
Registered: 26-4-2008
Location: Neon Thorium Erbium Lanthanum Neodymium Sulphur
Member Is Offline
Mood: pyrophoric
|
|
I suspect it might be a nickel plating. This metal is sometimes used as an alternative to chromium plating.
It is lustrous and silvery, and it dissolves only very very slowly in hydrochloric acid, so it fits your observations well.
You could add some hydrogen peroxide to your hydrochloric acid to speed up dissolution.
Both nickel and chromium metal would dissolve to yield green solutions, but nickel gives a much more pale green than chromium.
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Alright thanks, I'll try it again with hydrogen peroxide. Would the store bought 3% concentration be ok for this?
Oh, and the CuCl2 catalyst appeared to have no effect on the reaction.
[Edited on 3-10-2014 by zts16]
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Ok, I made up a fresh solution of hydrochloric acid and placed the metal in it. I added the H2O2, a fairly decent amount since
mine is not very concentrated. Because of this, I added a bit more concentrated HCl since it got very watered down.
It didn't appear to do anything, so either my peroxide isn't good enough or it isn't nickel.
Edit: Well, I went to check on the reaction again and the solution has become slightly green. However, it appeared that it was the copper
plating on the underside of the metal flake, not the mystery metal that reacted with the HCl/H2O2. This seems to make sense,
since both copper(II)chloride and nickel(II)chloride are greenish. I think I'll check by precipitating the metal with aluminum.
[Edited on 3-10-2014 by zts16]
|
|
|
Artemus Gordon
Hazard to Others
Posts: 178
Registered: 1-8-2013
Member Is Offline
Mood: No Mood
|
|
Could it possibly be Tin? Sn is lower on the reactivity list than either Cr or Ni. It's not normally very lustrous, but it has historically been a
common anti-rust coating on steel.
[Edited on 11-3-2014 by Artemus Gordon]
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
Melting it, would show you if it's tin.
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Good point. I'll check that out. The one thing that makes me doubt that it is tin, is that it is strong and difficult to bend whereas tin is extremely
malleable. Also, it appears to be shinier than tin.
Edit: I checked on the beaker in which the metal was stewing with HCl and hydrogen peroxide. It appears to have mostly dissolved and the
solution has become even more green. I'm feeling fairly convinced now that it is nickel. To be sure, I'll remove the water and do a flame test of the
salt tomorrow.
[Edited on 3-11-2014 by zts16]
|
|
|
Electra
Hazard to Others
Posts: 179
Registered: 11-12-2013
Member Is Offline
Mood: No Mood
|
|
Find out.
Put it in a graduated cylinder with water, or some other liquid volume measuring device, and use that technique to calculate the density. You should
know how to do this if you ever had a high school chemistry class. From the density you should be able to determine what metal you have, even if it is
slightly plated.
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Haha, yes, I happen to currently be in my first high school chemistry class this year. Thanks for the advice about the densities, I hadn't thought of
that. I see no need for the condescending tone though...
|
|
|
Electra
Hazard to Others
Posts: 179
Registered: 11-12-2013
Member Is Offline
Mood: No Mood
|
|
Apologies. I didn't intend for that to sound condescending. I was more so being sarcastic because based on your original post I immediately got the
impression you were well beyond high school chemistry.....(this is a complement). I didn't imagine you were still in the course. Keep studying! From
what I can tell you are far beyond any one I ever knew in a HS chemistry course.
|
|
|
phlogiston
International Hazard
Posts: 1379
Registered: 26-4-2008
Location: Neon Thorium Erbium Lanthanum Neodymium Sulphur
Member Is Offline
Mood: pyrophoric
|
|
It would be difficult to determine the density with any accuracy if it is only a plating. You would need a sizeable piece for that, ideally with a
known geometry that you can measure the dimensions of (cylinder, cube, etc).
Although both nickel and chrome are shiny and lustrous, there is a subtle difference in color. Chrome platings appear slightly bluish whereas nickel
has a subtly yellow tinge to it.
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
You might want to try and mechanically remove a small part of the coating, so the coating metal can be tested without affecting the substrate.
Nickel dissolves in HNO3 quite easily.
Tin is attacked by HNO3 but forms the insoluble, white SnO<sub>2</sub>. Tin does dissolve easily in HNO3/HCl.
[Edited on 11-3-2014 by blogfast25]
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Thanks for all the advice!
Based on the color difference, if any is there, it's definitely leaning towards yellow.
Since I don't have any nitric acid, I think I will flame test the chloride salt that I produced yesterday. It is lime green in solution, so it seems
like it may be nickel(II)chloride. From what I could find, nickel compounds should produce a light green flame.
Also, I realized that I made one silly blunder earlier, when I thought it might be silver, since the metal is magnetic...
|
|
|
DraconicAcid
International Hazard
Posts: 4355
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I didn't think nickel gave a colour to flame. Vogel doesn't list it with the flame tests.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Alright, so I flame tested it, and there was no color. I am definitely convinced that it is nickel.
-Magnetic
-Hard
-Lime green chloride
-Very shiny, possibly slightly yellow
-No flame test color
|
|
|
DraconicAcid
International Hazard
Posts: 4355
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
If you take your lime green chloride and heat it to about 80 oC, it should turn to the orange-brown anhydrous NiCl2.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
The classic test for nickel (II) is dimethylglyoxime, with which it forms a bright red complex. Unfortunately that reagent is hard to obtain.
The Ni<sup>2+</sup> ammonia complex is blue, so you might want to add an excess ammonia to your sample: http://en.wikipedia.org/wiki/Nickel#Nickel.28II.29
I didn't think nickel would yield anything in flame testing.
[Edited on 11-3-2014 by blogfast25]
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
I guess it's too late now, but nickel is magnetic, so if you still have any, test it with a magnet, that would narrow it down almost completely.
[Edited on 12-3-2014 by Zyklonb]
|
|
|
Texium
Administrator
Posts: 4618
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
Yeah, it's definitely magnetic.
Previously, I was under the impression that chromium, too, was magnetic, although I realize now that it isn't. As I said earlier today, it's most
definitely nickel.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
