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t_Pyro
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Copper hydroxide is blue. What you have might be copper chloride.
[Edited on 28-3-2004 by t_Pyro]
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Esplosivo
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It cannot be copper chloride too, since it is soluble and cannot precipitate, though it does have a bright green colour. This salt is also
hydgroscopic. Most chlorides are soluble (some exceptions like AgCl and PbCl2, the latter of which is also soluble in warm water). Copper Carbonate is
light blue, and it is insoluble forming precipitates (colour differences might be due to impurities).
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t_Pyro
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Yeah, you're right about the solubility of the copper chloride- drying a sample of it takes ages for me! I can't think of any green
inslouble compounds... Dark/dirty green could be ferric, dark green could also be a chromium salt. However, I've never heard of a green ppt on
adding NaOH to CuSO<sub>4</sub>.
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Organikum
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Dont judge a copper salt by its color.....
I just added hot concentrated NaOH solution to a concentrated CuSO4 solution and got some dark-green precipitate which dissolved under the addition of
more water to give a almost black solution.
Then I added some CuSO4 to an cold NaOH solution and got a deep-blue precipitate.
Temperature and concentration matters 
[Edited on 28-3-2004 by Organikum]
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Organikum
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And after my quick and dirty experiments the hard truth from the MERCK-index:
- copper sulfate dibasic: Cu3H4O8S, blue-green, rhombic, bipyramidal crystals. Practically insoluble in water.
- copper sulfate tribasic: Cu4H6O10S, very finely divided, light-blue, gelatinous particles. Practically insoluble in water.
- Copper hydrate: Cu(OH)2, blue to blue-green gel or light blue crystalline powder. Stability is dependent on the method of preparation, may decompose
to black CuO on standing a few days or upon heating. Practically insoluble in water. Sol. in concentrated alkali when freshly precipitated.
So I got CuO mostly in my first experiment and a mixture of everything in my second experiment I guess (no CuO up to now though)
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hodges
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| Quote: | Originally posted by Saerynide
Here's my new pvc protected electrodes. They're gonna be used for the HCl/NaOH cell Im planning on making this summer too
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Your electrodes look good. You can get NaOH by electrolysis of NaCl solution. But at the anode you will get Cl2 instead of HCl. The Cl2 will escape
into the air.
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Saerynide
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Yeah I know. Im making bubblers for the Cl2
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t_Pyro
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Bubbling the chlorine through water will result in HOCl and some dissolved chlorine. Check with some blue litmus- it'll turn red, and immediately
white. You might not even be able to detect the red.
Maybe you could make a jet of the H<sub>2</sub> gas to burn in an atmosphere of Cl<sub>2</sub> to get HCl, and then dissolve
it in water (use an inverted funnel for this). The beauty of this method would be that the H<sub>2</sub> gas could be taken from the
cathode, Cl<sub>2</sub> from the anode, so in effect, you'd actually be splitting up a salt of a strong acid/ strong base into the
constituents, NaOH and HCl!
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Saerynide
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I wouldnt dare mix H2 with Cl2. It becomes explosive
Also, the HClO formed would break down into HCl and O2. Ive tried shaking Cl2 into water before, and it does yeild HCl. I saw the O2 bubbles coming
out of the water
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darkflame89
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Haha, that's what i intend to do anyway. I want to make the 3 most needed acids mostly by this method. HCl acid, H2SO4 acid, (dunno its possible
to produce HNO3 this way).
Ignis ubique latet, naturam amplectitur omnem.
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Saerynide
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There was some thread awhile ago saying that NO3- can be reduced at the cathode. Which makes making HNO3 this way would be very hard and tedious, but
if there's a will, theres a way I forgot which thread it was though...
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Esplosivo
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Saerynide, as long as I remeber the decomposition of HOCl to HCl takes a long time to occur. Also Cl2 dissolves in water to give the so-called
'chlorine water'. The Cl2 dissolved in water will decompose to HOCl in the presence of light.
Organikum, sorry for my ignorance, but how can copper sulphate be di/tri basic?! I've never heard of these salts before. How are they made?
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darkflame89
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Never mind i can always use sulphuric acid on nitrate salts to produce NO2 and dissolve them in water to get HNO3.
Or i can always try to biuld a Birkeland-Eyde reactor..
Ignis ubique latet, naturam amplectitur omnem.
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Saerynide
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Im gonna put the HCl/NaOH maker outside anyways, so there'll be lotsa light for it. Im so not getting Cl2 poisoning again, so it wont be
anywhere near indoors
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darkflame89
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Liked to clarify somthing.. In this setup of the MgSO4, won't the OH ions move over to the anode side instead of producing Mg(OH)2?
Ignis ubique latet, naturam amplectitur omnem.
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Saerynide
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Mg(OH)2 is insoluble and it precipitates out of the solution You get a
slightly milky solution if its dilute, but you get a layer of white amorphous stuff if its concentrated.
Edit: Nooooooo!!!! Contamination again!!!! 
How far do you hafta go to have no fscking contamination?!?! I already washed ALL the equipment with dH2O TWICE already  Maybe the silicone doesnt make as much of a water proof seal as I hoped
The contamination's not as bad this time though. I tested the pH of the two solutions. It was 3 and 11 only
[Edited on 29-3-2004 by Saerynide]
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Esplosivo
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Come on Saerynide. Look at it on the good side, now you can carry out an analysis on the stuff  . Its fascinating though. I cannot immagine the source of impurity. What water did you use? What was the source of the
chemicals electrolyesed?
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Saerynide
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Lol, I guess I could have some fun with it now But no H2SO4 for making sugar
turn to carbon or making oil of wintergreen
I used bottled distilled water from the supermarket and MgSO4.7H2O from the drug store, so it *should* be pretty pure. I also covered the cell with
Glad wrap to keep dust from getting in.
I just thought of something. There would be an equilibrium in the cell because the electrodes can only draw so many anions and cations to them at any
moment, so some ions would be crossing the salt bridge and reacting back with the acid and base to make the salt again, right? That would mean, the
pH can only go so low, and would level off. So, if after each run, I filter the catholyte out to remove the Mg(OH)2, then run it again, and keep
repeating this, I would eventually remove all the Mg2+ from the cell? That
would leave me with a salt free acid?
Considering the fact that I am extremely tired/sleepy right now, it might be all gibberish
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Esplosivo
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I think you're right. Well, Mg2+ ions would move to the cathode, therefore there is not need to filter after each run. One must calculate the
point at which the salt will be all used up (and I have no idea how this can be done, probably when Mg(OH)2 stops forming). Btw, I do not think that
you will be able to get an acid totally free of the salt. Some would still remain, though in very small conc.
Just call in when carrying out the qualitative analysis. (I love analysis lol)
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Saerynide
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| Quote: | Originally posted by Organikum
And after my quick and dirty experiments the hard truth from the MERCK-index:
- copper sulfate dibasic: Cu3H4O8S, blue-green, rhombic, bipyramidal crystals. Practically insoluble in water.
- copper sulfate tribasic: Cu4H6O10S, very finely divided, light-blue, gelatinous particles. Practically insoluble in water.
- Copper hydrate: Cu(OH)2, blue to blue-green gel or light blue crystalline powder. Stability is dependent on the method of preparation, may decompose
to black CuO on standing a few days or upon heating. Practically insoluble in water. Sol. in concentrated alkali when freshly precipitated.
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Ohh!!!! The time when I tried to make CuSO4 using a paper towel salt bridge, MgSO4 electrolyte ,and a copper penny anode, I *did* get CuSO4 but I
thought I got Cu(OH)2. So that light blue precipitate was the tribasic stuff??
Btw, I looked up the definition of di/tribasic, and I dont understand how CuSO4 could fit the definition. Also, I dont understand why there are all
those extra O's, H's and Cu's
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Esplosivo
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Same problem here. I've searched again and again and I cannot find any basic copper sulphate, neither di- not tri- basic. Are these complex salts
or basic salts. Copper ions may form complexes having 'strange' colours. Well if they exist I would like a simple experiment for making one
plz.
Organikum, not to doubt your knowledge, but are the chemical formulas stated previously correct. They do not seem correct to me, though I might be
wrong.
[Edited on 29-3-2004 by Esplosivo]
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t_Pyro
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When you're talking of a dibasic/tribasic <i>acid</i>, it refers to the number of replaceable hydrogen atoms. For sulfuric acid,
it's 2. I've never heard of any dibasic or tribasic salt, though what he might have meant might have been the oxidation state of the element
in the compound. The "Cu3H4O8S" might be Cu<sub>3</sub>(OH)<sub>4</sub>SO<sub>4</sub>, where the copper
is in the +2 state. This compound is new to me, though. I would have thought the hydroxyl ions would have been easily neytralised in the solution...
The "Cu4H6O10S" might be Cu<sub>4</sub>(OH)<sub>6</sub>SO<sub>4</sub>, but I have the same misgivings
against it...
Copper ions do form complexes, but I wouldn't call them "strange" coloured! A simple way to prepare tetramine copper sulfate is to add
excess of ammonium hydroxide to a solution of copper sulfate. The dark blue solution is tetramine copper sulfate, a simple coordination compound.
Nickel and cobalt salts respond to ammonium hydroxide in a similar manner, forming coordination compounds.
Coming back to chlorine and hydrogen: A <i>mixture</i> of H<sub>2</sub> and Cl<sub>2</sub> is explosive, as is a
mixture of H<sub>2</sub> and O<sub>2</sub>. However, if pure H<sub>2</sub> is burnt as a jet in an atmosphere of
O<sub>2</sub> or Cl<sub>2</sub>, the reaction is quite smooth.
[Edited on 29-3-2004 by t_Pyro]
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Esplosivo
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Ok lol. Sorry for my terminology. By 'strange coloured' I wanted to mean that it is not common to see copper (II) sulphate which is green,
at least for me. I still don't know how those basic compounds named may be formed.
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Saerynide
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What should I try tomorrow? I could 1) filter out Mg(OH)2 and electrolyze again or 2) Try to plate out the metal ion contaminant.
Id like to try plating out the metal, because it will either remove the contaminant (best case scenario), or end up concentrating the acid to the
point that the bit of metal is irrelevant (still good)
But, how am I gonna keep the freshly made metal from reacting with the H2SO4? More over, what if the H2SO4 reacts with my cathode??
Analysis would have to wait til Friday. Ive got a ton of tests to study for
Btw, how do you guys fill your salt bridges? I had to shove a smaller tube into my bridge, dunk that under the solution, and use a syringe to suck up
the liquid. This way is *very* painstaking cause I have to do it many times. Bubbles always get in and liquid goes into the syringe! I was so desperate once I even tried syphoning up the solution. That left me
choking on a mouthful of nasty MgSO4
Does anyone pity poor me?? *sniffle*
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darkflame89
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You could try and plate out the metal. You will have to keep an eye on it. When the solution turns clear( i hope this does no take too long), you can
just remove the cathode and anode. I dunno about this, hope it works.
About the filling of salt bridge, i am worrying over it too.
Another thing, i need to know how long i can run the setup if i have a 9 V battery or a 18 V series of batteries.
Ignis ubique latet, naturam amplectitur omnem.
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