Producing Sulphuric Acid From Copper Sulphate
Posted 07 September 2019 - 08:22 AM
Posted 07 September 2019 - 08:58 AM
I'm thinking about starting a new business and part of the business requires me anodise aluminium. Anodising usually requires a 15% solution of dilute sulphuric acid, however purchasing sulphuric acid in the UK in quantity is now difficult and requires a licence. So I've thought about making it myself. I understand the process of making sulphuric acid from copper sulphate through electrolysis is straight forward and copper sulphate is readily available. My question really is what weight of copper sulphate would I need to add to a given quantity of distilled water, say 100ml, to produce a 15% solution of sulphuric acid? Would it be as simple as 15g copper sulphate per 100ml water?
This does not ring true. Any UK business can get a licence for buying sulphuric acid. Also, as I understand it, you can use other acids for the anodising process, besides sulphuric acid. And I doubt it will make economic sense to make sulphuric cid by electrolysing CuSO4 yourself.
To be honest, it sounds to me as if you either want to make sulphuric acid for another, perhaps less innocent, reason, or else this is a homework question.
If it is a homework question, I'll be glad to help but not to do it for you. What electrodes would you use to make H2SO4 from CuSO4, and what do you think would be the products of electrolysis at each electrode, i.e. what would be the overall reaction?
Edited by exchemist, 07 September 2019 - 09:47 AM.
Posted 07 September 2019 - 11:56 AM
- exchemist likes this
Posted 07 September 2019 - 12:59 PM
My business idea is only in its preliminary stages. It's actually been spawned from me wanting to build a coffee table which has built in DC motors that slide a projector out the table at the touch of a button. Part of the table decoration is a laser engraved design on aluminium inlaid into the the top, however the laser will only engrave onto anodised aluminium. Some basic research revealed that to anodise aluminium I need a weakish solution of sulphuric acid to do this, and by submerging the workpiece in the acid and having the workpiece act as the cathode and platinum coated titanium as the anode then running a current through, the anodising reaction would begin. My plan was to dissolve some CUSO4 in distilled water and again using platinum as the anode but copper as the cathode, when a current is passed through the solution, copper metal should be deposited on the cathode eventually turning the solution clear, this clear solution should be dilute H2SO4. I understand that I could just buy anodised aluminium, but I just thought the it would be fun to anodise it myself
Oh that's different, then. .
Apologies for being suspicious but, as you may know if you frequent forums like this, we do get a lot of people who are not what they pretend to be - including students trying to cheat with their homework!
I'm not any kind of expert on anodising aluminium but yes I think if you have a copper cathode, the Cu 2+ will migrate there and be discharged to copper metal while sulphate will migrate to the anode, where it will get neutralised but will then pinch electrons from water molecules, leading to a rising concentration of H+. There should be no net change in sulphate content as a result so one mole of CuSO4 should in theory give you one mole of H2SO4, I think.
However, when anodising, you need to make the Al the anode (hence the name), as you want to build up an oxide layer on it. So you want oxygen to be released at the workpiece rather than hydrogen.
But I think you can still get sulphuric acid from hardware stores in the UK as a drain cleaner - I was offered some the other day, as an alternative to caustic soda. The government regs seem to cut in when you reach a solution of 15% w/w. Presumably the drain cleaner is less than that. I don't know what strength acid you need to anodise Al but it may not need to be very strong. All the acid does is condition the oxide layer as it forms, I understand, by dissolving it (paradoxically), so that there is a balance between competing oxidation and dissolving.