borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Multiple Complex Synthesis (3rd attempt)
So after a couple of experiments where I successfully made copper(II) carbonate but not much else I decided to scrap the experiments and start again.
I had 4 vials:
Vial 1 = 2g copper wire
Vial 2 = 4g copper wire
Vial 3 = 7g copper wire
Vial 4 = 10g iron screws
They were immersed for 40 hours in 4% w/w sodium hypochlorite solution.
After the immersion time, all vials showed colour change. Vials 2 and 3 were dark grey (I assume CuO because hypochlorite can oxidise metal but this
may not be the case). Vial 4 showed a brown colour so I assume FeO but I know iron oxide is red so maybe this is a low amount in solution or maybe not
even FeO. Vial 1 showed no change.
In previous experiments I added sodium bicarbonate (very fine powder used for baking) and immediately witnessed a colour change to aqua blue, likely
copper(II) carbonate. However my frustration is deepening when I added in sodium bicarbonate granules (product for cleaning but still sodium
bicarbonate) to the copper solutions to no effect. They are not soluble I think because they are larger granules but they are the same bloody
compound. So, no joy there. I made an aqueous solution in a separate vial and added it into the copper solution and still no reaction.
I am starting to wonder if sodium hypochlorite is simply a bad choice for an oxidising agent. Sure, it oxidises the metal, but I can't help wondering
if it is interacting with the counter ions I add into solution when I want the extracted Cu(II) ions in the solution to interact with the new reagent.
Is hydrogen peroxide generally the best oxidising agent for metal objects in the formation of metal complexes?
[Edited on 12-7-2017 by borontrichloride]
|
|
|
woelen
Super Administrator
Posts: 8014
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
For clean oxidation of copper (and many other metals) you need a strong acid, besides the oxygen-containing oxidizer.
Many common oxidizers are oxo-compounds (either ionic or molecular, including oxygen from air). In the oxidation process either oxide is produced, or
hydrogen ion is consumed. In the latter case, the oxidation proceeds much smoother and faster and more cleanly.
If you have 10% hydrochloric acid, then add some copper to this acid and then add some hydrogen peroxide, or a little bleach. You will see that the
solution quickly turns nice green and remains clear. You get a clean solution with copper(II) ions in solution, coordinated to chloride ions.
If you want really clean oxidation, then use dilute sulphuric acid, and a non-chlorine-based oxidizer (e.g. H2O2, or Na2S2O8). You then get a bright
blue solution, containing hydrated copper(II) ions.
Without the acid, the oxidation is slow, and you get dirty-looking goop, in the form of oxides, or basic salts, depending on what oxidizer is used.
A word of warning: If you use bleach, combined with acid, be careful to do the experiment only on a small scale (a few ml in test tubes). You will get
chlorine gas. A few ml is not a real issue besides the smell, but large quantities can be dangerous in a small confined space.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Thanks woelen.
After looking up oxidising acids (acids that produce strong conjugate bases usually with oxygen in the structure) I realised the importance of acids.
Your advice is great. I have 4 new experiments now:
VIAL 1: copper metal immersed in conc. lactic acid,
VIAL 2: copper metal immersed in 5% acetic acid (vinegar),
VIAL 3: copper metal immersed in sodium hydroxide,
VIAL 4: iron screws immersed in sodium bicarbonate.
All oxidising agents are oxygen-containing. This process is certainly slower than with the hypochlorite so I am eager to see if the respective
compounds form as oxidation occurs. When all of them were immersed in hypochlorite, they all simply went black.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
I'm now actually tempted to throw in a small amount of the thick bleach (4% w/w sodium hypochlorite) into the vials; attempt to speed up the
oxidation.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
OK so I discarded the reactions that had metals simply sitting in base. Given the advice of the great guys on this forum I have realised that the acid
dissolves the metal whilst the slight touch of base simply assists the oxidation process in conjunction with the conjugate base of the acid. In the
case of HCl of course some oxidising agent will be needed.
VIAL 1: copper metal immersed in conc. lactic acid - solution is light brown and has been since last night when I added the 4% sodium hypochlorite of
which more was added today and still no reaction progression. I can't find any mention online of the colour of copper(II) lactate but a lot of metal
lactates are white powders so I will assume this to be an unsuccessful experiment.
VIAL 2: copper metal immersed in 5% acetic acid (vinegar) - following the addition of some sodium hypochlorite, the solution turned green and then
into blue. It is currently light blue and allegedly copper(II) acetate is a dark blue so either it's [Cu(H₂O)₆](acetate)₂ or a heavily diluted
solution of copper(II) acetate.
Vials 3 and 4 were discarded as previously mentioned.
I then set up two new experiments. I will name these vials 3 and 4.
VIAL 3: copper metal immersed in hydrochloric acid - the metal immediately reacted with the acid and further more when some 2% sodium hypochlorite
was added. The solution is now green and thus I propose this to be CuCl₂.
VIAL 4: iron metal immersed in hydrochloric acid - the metal immediately reacted with the acid and further more when some 2% sodium hypochlorite was
added. The solution is now orange-brown and thus I propose this to be FeCl₃.
The iron reaction also produced a white precipitate which is likely NaCl.
[Edited on 13-7-2017 by borontrichloride]
|
|
|
DraconicAcid
International Hazard
Posts: 4333
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I've made copper lactate. It's blue.
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.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Thanks for that. I guess it is unsuccessful then. Although I wonder if it's copper(I) lactate? I think I remember copper(I) iodide being brown.
[Edited on 13-7-2017 by borontrichloride]
|
|
|
DraconicAcid
International Hazard
Posts: 4333
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Copper(I) halides are white when they are pure.
If you want copper lactate, react the copper with bleach in the presence of an acid that's not going to react directly with the bleach- hydrochloric,
acetic, sulphuric, but not lactic. Add sodium carbonate to your blue solution to get a precipitate of "copper carbonate", collect the precipitate,
and add lactic acid to that.
Attached: Copper(II) lactate trihydrate
[Edited on 13-7-2017 by DraconicAcid]
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.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
OK so a little update.
VIAL 1: copper metal immersed in conc. lactic acid - this solution has been poured out. It didn't come anywhere near blue. I will need to re-think my
method. I appreciate the suggestion from DraconicAcid. I will likely try that.
VIAL 2: copper metal immersed in 5% acetic acid (vinegar) - the solution was turquoise/light blue. I poured approximately half into a separate jar and
into this jar I added 5% ammonia. It took very little before turning dark blue. I therefore conclude that this was enough of an excess to form the
octahedral ammonia complex and thus I propose this to be [Cu(NH₃)₆](acetate)₂.
VIAL 3: copper metal immersed in hydrochloric acid - solution turned green very quickly and I left it to react to see if it would turn more green thus
producing more CuCl₂. It was almost turning black however and so I extracted it and left it as it is.
VIAL 4: iron metal immersed in hydrochloric acid - here is the one that confused me and I would welcome any interpretation from you guys. I produced a
dark yellow complex which I conclude to be FeCl₃. However, upon adding ammonia to a portion of this solution in a separate jar, it initially turned
a lighter yellow. I considered this too similar to the FeCl₃ and so I added more ammonia. It needed quite a lot of ammonia but here is the confusing
bit; the solution turned pitch black and small brown specks appeared at the top.
Have I reduced the Fe²⁺ back into the ground state metal?
[Edited on 14-7-2017 by borontrichloride]
|
|
|
DraconicAcid
International Hazard
Posts: 4333
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Quote: Originally posted by borontrichloride  |
I therefore conclude that this was enough of an excess to form the octahedral ammonia complex and thus I propose this to be
[Cu(NH₃)₆](acetate)₂. |
Copper only forms a tetrammine comnplex, not a hexammine.
| Quote: | | VIAL 3: copper metal immersed in hydrochloric acid - solution turned green very quickly and I left it to react to see if it would turn more green thus
producing more CuCl₂. It was almost turning black however and so I extracted it and left it as it is. |
If it's turning black, that just means you have more copper in solution. At high concentrations of chloride, you get CuCl4(2-) ions, which are
yellow. There are also some complex mixed oxidation-state ions that are dark brown.
| Quote: | VIAL 4: iron metal immersed in hydrochloric acid - here is the one that confused me and I would welcome any interpretation from you guys. I produced a
dark yellow complex which I conclude to be FeCl₃. However, upon adding ammonia to a portion of this solution in a separate jar, it initially turned
a lighter yellow. I considered this too similar to the FeCl₃ and so I added more ammonia. It needed quite a lot of ammonia but here is the confusing
bit; the solution turned pitch black and small brown specks appeared at the top.
Have I reduced the Fe²⁺ back into the ground state metal? |
Iron doesn't form a complex with ammonia. You've probably just formed hydroxides.
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.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Thanks DraconicAcid. Would this seem a fair interpretation (below)?
2 FeCl₃ + 6 NH₃ + 14 H₂O → 2 Fe(OH)₂ + 6 NH₄Cl + 9 H₂ + 5 O₂
I propose iron(II) hydroxide instead of iron(III) because Fe(OH)₂ is black and Fe(OH)₃ is yellow I think.
There was a gaseous release and it was an exothermic reaction because the jar became quite warm.
|
|
|
DraconicAcid
International Hazard
Posts: 4333
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
No. FeCl2 + 2 H2O + 2 NH3 --> Fe(OH)2(s) + 2 NH4Cl(aq) The exothermicity is from the reaction of ammonia with excess hydrochloric acid (as is,
probably, the "gaseous release" you mention.
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.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Oh right. I started with FeCl₃ though. I know this due to the dark yellow colour. FeCl₂ is off-white.
Regardless, thanks for your advice. Has been helpful. I am now researching where I can obtain some chromium salts. I would like to make CrCl₃
because it is a nice purple. Chromium metal is expensive on Amazon.
[Edited on 14-7-2017 by borontrichloride]
|
|
|
DraconicAcid
International Hazard
Posts: 4333
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
You probably want to leave chromium until you've had more experience with transition metal chemistry. Due to its low lability, it can do some weird
and confusing things.
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.
|
|
|
borontrichloride
Harmless
Posts: 23
Registered: 8-7-2017
Member Is Offline
Mood: No Mood
|
|
Fair enough. I do want to experiment with different transition metals. During my degree we did a lot of work with Nickel(II) but I would rather not
use that at home as it is apparently carcinogenic.
I have done quite a bit of work with transition metals but we always started from the salts already made, ordered from Sigma I think. This is the
first time I have been starting from the actual ground state metals and for that reason there is an element of learning as I go.
I'll do a bit more reading around for interesting experiments. Hopefully I can find some sulfuric acid and make copper(II) sulfate as I haven't done
that yet.
|
|
|
woelen
Super Administrator
Posts: 8014
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I siggest you first keep on experimenting with copper metal. It gives a lot of experience and it has quite rich chemistry.
Chromium is interesting and only marginally toxic (in the +3 oxidation state), so it seems a good candidate, but it also has very complicated
chemistry, which is not easy to start with.
With your iron, the black precipitate is perfectly understandable. It must be Fe3O4, a mixed oxidation state oxide. If you add hydroxide (or ammonia)
to a pure iron(III) solution, then you get a brown precipitate. If you add hydroxide to a pure iron(II) solution, you get a nearly white precipitate.
The latter, however, quickly turns grey or even black, due to oxidation of iron(II) to iron(III) by oxygen from air.
If you add hydroxide to a mix of iron(II) and iron(III) then you get a dark grey precipitate or even black precipitate. This precipitate is funny, it
is attracted by a magnet:
http://woelen.homescience.net/science/chem/exps/magnetite/in...
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
Copper is fun to play with and not that toxic. It also does cool color changes and has a lot of uses.
Calcium is quite useful, although it can be hard to work with at times and calcium salts are usually just white powders.
I like chromium a lot. There is a lot of chromium chemistry to explore. Chromium can be turned to any color of the rainbow and can also be used in
making a wide variety of oxidizers. It can be pretty toxic in higher oxidation states but is far less toxic than mercury.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
