# How to change pH of CuSO4?

I'm currently very desperate to know this. I'm doing an experiment about adsorption/chelation of Cu ions in CuSO4 using banana peels using the UV-Visible spectroscopy. How do I change the pH of 0.1M or 0.05M of CuSO4? My main problem is with making the CuSO4 more alkaline/basic or whatever u call it. My base pH of the CuSO4 is about 4.0 or 4.1. I've tried making a pH 5 and pH 6 buffer solution then adding it to the CuSO4 solution, but so far it's only made it more acidic. These are the calculations:

pH 5

Weigh 5.10575g of potassium hydrogen phthalate(C8H5KO4)

Dissolve it in a 100cm3 beaker

Transfer the solution into a 250cm3 volumetric flask and fill it with distilled water

Measure 5cm3 of 2M NaOH and add it to a 100cm3 volumetric flask and fill it with distilled water

Add 100cm3 of 0.1M potassium hydrogen phthalate and 45.2cm3 of 0.1M NaOH into a 250cm3 beaker

pH 6

Weigh 3.40225g of KH2PO4

Dissolve it in a 100cm3 beaker

Transfer the solution into a 250cm3 volumetric flask and fill it with distilled water

Measure 5cm3 of 2M NaOH and add it to a 100cm3 volumetric flask and fill it with distilled water

Add 100cm3 of 0.1M KH2PO4 and 11.2cm3 of 0.1M NaOH

But I realised the phosphate buffers will cause a precipitate. So I tried using a citric acid + sodium citrate buffer. I think the calculation was something like 88.5cm^3 of citric acid + 11.5 cm^3 of sodium citrate (or it could be the other way around). But I don't think that was a buffer solution, cus when I added the slightest amount of deionized water, the pH increased. I got it from https://www.sigmaaldrich.com/life-science/core-bioreagents/biological-buffers/learning-center/buffer-reference-center.html#citric2.

edit: I made the copper sulfate using CuSO4.5H2O if that helps

• Copper(II) phosphate is insoluble in water. You could try acetic acid/acetate (pKa = 4.76). pH 6 is very close to neutral i would rather avoid it. – Andrew Kovács Aug 29 '20 at 15:06
• Also this could be useful for you. microscopy.berkeley.edu/Resources/instruction/buffers.html – Andrew Kovács Aug 29 '20 at 15:17
• Your major problem is you are using way too high concentration of your $\ce{Cu^2+}$ solution. The maximum concentration of solutions used in the given article is $\pu{0.00787 M}$. – Mathew Mahindaratne Aug 30 '20 at 21:19

Alas, your problem has no solution, because of the risk of producing $$\ce{Cu(OH)2}$$ at pH $$5$$ and $$6$$. The solubility product of $$\ce{Cu(OH)2}$$ is $$2.2·10^{-20}$$. At pH $$5$$ and $$6$$, the concentrations of the ion $$\ce{OH-}$$ are respectively $$10^{-9}$$ and $$10^{-8}$$ M. This means that, at pH $$5$$ and $$6$$, the maximum molar concentrations of the ion $$\ce{Cu^{2+}}$$ are $$2.2·10^{-2}$$ M and $$2.2·10^{-4}$$ M respectively. So you will never be able to obtain a $$0.1$$ M solution of $$\ce{CuSO4}$$ at pH $$5$$ or $$6$$.