BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
Sulfuric acid from copper sulfate
I am going to make some sulfuric acid through the electrolysis of a solution of copper sulfate (CuSO4 + 2H2O --electricity-->
H2SO4 + Cu + O2 + 2H2) and by taking advantage of how cheap it is to buy it and how copper does not really
react with sulfuric acid, (unless under special conditions) I am hopping it will work. But because of the strong oxidizing conditions, I need to have
an anode that will not corrode or react during the process. Any ideas (besides: platinum, lead dioxide(?), and maybe even gold...)?
"The last man in the world was in his room. There was a knock on the door"
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
What voltage current do you intend to use?
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
Somewhere around 5 or 6 volts... maybe more... it depends i guess.
[Edited on 20-6-2012 by BiOCl:BoatsInOutCrashLand]
"The last man in the world was in his room. There was a knock on the door"
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
Why not use thick carbon electrodes and filter out your solution when you're done?? 5 or 6V at a reasonable current will not erode them too
much....how much acid do you want to produce??
Are you doing this for the experience of doing electrochemistry or do you actually need sulfuric acid? If the latter, then NurdRage has a great video
on doing it using methods much better than here, or you can just buy a suitable drain opener from the hardware store. Look at the labels; some of them
contain up to 98%, 18M H2SO4.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
Carbon rod will corrode But contamination will be easy to remove
I never asked for this.
|
|
|
MR AZIDE
Hazard to Self
Posts: 64
Registered: 21-5-2012
Location: UNITED KINGDOM
Member Is Offline
Mood: Fizzing
|
|
Have you actually managed to purchase Conc Sulphuric, as drain cleaner in the UK, Hexavalent.??
If so, what were the products.....???
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
If I try to filter the solution, can't it absorb the water molecules from the filter (unless it's very dilute)? I am doing this for the experience and
i plan to use it for a few experiments that i'm planning.
"The last man in the world was in his room. There was a knock on the door"
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
filter the acid before boiling it. T
I never asked for this.
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
Mr Azide
This stuff I bought from a local building supplies company, but I imagine it can be purchased at hardware stores;
http://www.amazon.co.uk/Knock-Out-Drain-Cleaner-Ltr/dp/97921...
The label on my bottle specifies 97% H2SO4.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
Yea, as plante1999 said you can just filter off the any small carbon bits that break off the electrode.
Filtering it will not dilute the acid in any way....if you feel confident, you can boil down the acid to concentrate it....but I would not advise it,
certainly not to someone with little experience such as yourself. If you want more concentrated acid, I would recommend simply purchasing it as I
mentioned above.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
My plann is to have a higher concentration of the copper sulfate (not too much though) to have a higher concentration of the acid (that way i do not
need to boil it down).
"The last man in the world was in his room. There was a knock on the door"
|
|
|
Hexavalent
International Hazard
Posts: 1564
Registered: 29-12-2011
Location: Wales, UK
Member Is Offline
Mood: Pericyclic
|
|
At 5-6V, conversion time for concentrated copper sulfate solution to sulfuric acid is going to be quite long, and even then you may need to boil it
down.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
Well i do not need it to be too concentrated for my needs and if the process is going too slowly, i will use a higher voltage...maybe 9.
"The last man in the world was in his room. There was a knock on the door"
|
|
|
bob800
Hazard to Others
Posts: 240
Registered: 28-7-2010
Member Is Offline
Mood: No Mood
|
|
<s>I just calculated that this method is capable of generating acid of no higher than 1 M concentration, but this seems a bit low. Could someone
please double-check my work?</s> -- [EDIT -- OK, found my mistake. 9% seems more reasonable]
CuSO<sub>4</sub> becomes saturated in water at 320 g/ 1 L (anhydrous, @STP)—that's a 2 M solution. Assuming that this reaction proceeds
with 100% efficiency and to completion, 2 mol of CuSO<sub>4</sub> will form 1 mol of H<sub>2</sub>SO<sub>4</sub>:
2H<sub>2</sub>O ---> O<sub>2</sub> + 4H<sup>+</sup> + 4e<sup>-</sup>
<s>2 SO<sup>-2</sup><sub>4</sub> + 4H<sup>+</sup> ---> 1
H<sub>2</sub>SO<sub>4</sub></s>
So the 2 M solution of copper sulfate is now a <s>1</s>2 M solution of sulfuric acid in water—<s>98 g/
L</s>196 g/ L, or roughly <s>9%</s> 20% concentration.
[EDIT: 2nd equation was completely unbalanced. See next post for correction]
[Edited on 21-6-2012 by bob800]
|
|
|
vmelkon
National Hazard
Posts: 669
Registered: 25-11-2011
Location: Canada
Member Is Offline
Mood: autoerotic asphyxiation
|
|
9%? That's pretty weak.
Personally, I get my H2SO4 from old car batteries. Some people abandon them or you can visit certain garages (they leave the old ones outside). Just
connect them to a battery charger for 1 day just to make sure it makes as much H2SO4 as possible. The conc never reaches the same as a new battery.
I'm in Canada and there aren't hardware stores that sell H2SO4. They are all NaOH + sodium metasilicate.
|
|
|
barley81
Hazard to Others
Posts: 481
Registered: 9-5-2011
Member Is Offline
Mood: No Mood
|
|
In a 2M solution of copper sulfate, there are 2 moles of sulfate ions per liter. If only one mole of sulfuric acid were formed per two moles of copper
sulfate, where would the extra sulfate go? It's a 1:1 ratio between copper sulfate and sulfuric acid, and you would get 196g sulfuric acid per liter
of solution (roughly 20% concentration)
|
|
|
bob800
Hazard to Others
Posts: 240
Registered: 28-7-2010
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by barley81  | | In a 2M solution of copper sulfate, there are 2 moles of sulfate ions per liter. If only one mole of sulfuric acid were formed per two moles of copper
sulfate, where would the extra sulfate go? |
Ah, you are correct! My second equation is only balanced electrically (there's more sulfate on the left side than the right, as you pointed out). It
is a 1:1 ratio:
2 SO4 + 4 H ---> 2 H2SO4
in which case you are correct about the 2 M acid. I guess this method isn't quite as bad as I originally thought!
|
|
|
electrokinetic
Hazard to Self
Posts: 52
Registered: 6-6-2012
Location: Rhuidean
Member Is Offline
Mood: No Mood
|
|
Funny you posted this, I'm doing the same thing except my goal isn't to get sulfuric acid, it's just to do some copperplating. Here's a couple of
things I found out.
First, you can always dissolve more copper sulfate. Those calculations are based on starting with a saturated CuSO4 solution but there's no reason
you can't keep it saturated. I kept adding CuO or copper carbonate because I didn't want the pH to get too low.
Second, when the pH gets low enough, hydrogen will start evolving at the cathode even with copper ions present. I know there's some equation that
explains that, but I'm too lazy to look it up right now. All I know is, I came in after having it going all weekend and no matter how low I turned
the voltage I couldn't keep hydrogen from evolving.
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
@barley81-Yes, I began to wonder that also. @electrokinetic- I was planning to supper saturate my solution originally...but now that i think of it,
the copper sulfate will most likely recrystalize during the procces...keeping it saturated may be a better idea... the hydrogen gas may be the
electrolysis of the acid i am guessing...
Okay, so now I have finished a few calculations on how long the electrlysis proces may take... it wil take three days (at 9 volts) to make one mole of
the acid (but instead of using a molar solution, I will be using just enough water to make a saturated solution of one mole of copper sulfate). But of
course the the proces is not 100% efficient so there will be liess than the expected yield...but as long as i add a bit more than a mole it still may
be able to reach the expected yield... I think I will try what electrokinetic did another time...
[Edited on 21-6-2012 by BiOCl:BoatsInOutCrashLand]
"The last man in the world was in his room. There was a knock on the door"
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Compare the product results above with mixing Oxalic acid (usually available as a dihydrate) and CuSO4 and filtering:
H2C2O4.2H2O + CuSO4 --> H2SO4 + CuC2O4 (s) + 2 H2O
Obviously, dissolving the CuSO4 creates even further dilution requiring the final Sulfuric acid product to be concentrated. See the thread "Using
Oxalic Acid to Make H2SO4,..." at http://www.sciencemadness.org/talk/viewthread.php?tid=18963
Copper oxalate is an interesting salt as upon heating it in an inert atmosphere liberates Cu, but in air stream for some 2 hours, produces a very fine
CuO with a high surface area (useful for catalytic purposes). See: "Copper Oxalate synthesis and decomposition" at http://info.tuwien.ac.at/struchem/files/poster_cuox.pdf
|
|
|
BiOCl:BoatsInOutCrashLand
Harmless
Posts: 8
Registered: 19-6-2012
Location: In front of computer
Member Is Offline
Mood: Happily skipping through a field of unnecessary flowers
|
|
Thanks for the idea AJKOER! I will try that some day...
If the pentahydrate is used... CuSO4.5H2O + 2H2O --electrcity--> H2SO4 + 5H2O +
Cu + H2 + 2O2... meaning that there will be 5 moles of extra water in the final products (not including the water that did not
come from the pentahydrate). On that note, is it possible to heat the pentahydrate to drive of the water molecules without it decomposing? If so, how
should it be heated?
[Edited on 22-6-2012 by BiOCl:BoatsInOutCrashLand]
"The last man in the world was in his room. There was a knock on the door"
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
