December 13, 2019, 08:26:47 AM
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Topic: Excess H2SO4 required to make effective electrolyte solution? (electrochemistry)  (Read 622 times)

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Offline rwooduk1

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I want to make an acidic electrolyte solution for an electrochemical experiment. I've considered some possibilities and decided on ##H_{2}SO_{4}##.

When ##H_{2}SO_{4}## is added to water it disassociates and the following reactions occur:

##H_{2}SO_{4} + H_{2}O \rightarrow HSO_{4 (aq)}^{-}+ H_{(aq)}^{+}##

##HSO_{4 (aq)}^{-} + H_{2}O \leftrightharpoons SO_{4(aq)}^{2-} + H_{(aq)}^{+}##

##H^{+}+ H_{2}O\rightarrow H_{3}O^{+}##

For electrochemistry (aq) I've read that the electrolyte needs to be in large excess; "at concentrations ranging from 0.1 to 1 M".

The volume of water I will be using is around 8 L. I've read that the limiting reactant is the one with the least number of moles. For 8 L of water I calculate ~440 moles. So I need a greater number of moles of ##H_{2}SO_{4}##. Which I can calculate, but I am unsure of how much excess ##H_{2}SO_{4}## I would need. What are the factors that determine whether 0.1 M up to 1 M would be suitable?

I'm a little lost, please could someone give me a point in the right direction.

Offline AWK

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The electrolysis of water for experimental purposes is performed on a very small scale and then the concentration of acid or base is not particularly important. Large-scale electrolysis is already very dangerous but technologically mastered. I advise against playing on such a scale unless someone wants to commit spectacular suicide.
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Offline rwooduk1

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A genuine thanks AWK, unfortunately I need to try get this thing working. With your words of warning I think I will try without acid first to practice operating the system. Then, if I'm not dead, will come back to ask again about the acid (needed to lower the pH of solution). Beforehand, I will see if there are any electrochemists in other departments of the university, I hadn't realised it was so dangerous.

Online Borek

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Limiting reactant is not something that you should worry about when talking about dissociation (and no, in general it is not just the one with the lower number of moles, as it depends on the ratio in which things react).

What is the reaction that you are interested in? Just water electrolysis? Then the acid is there only to lower the resistance (it increases conductivity of the solution). In other situations it may play some additional roles.

As to dangers - produced hydrogen easily mixes with oxygen (either the atmospheric one or produced in the anode reaction) and the mixture is highly explosive (and easy to ignite - it won't explode on its own, but just a spark or flame and it goes kaboom). The less hydrogen produced, the safer you are.
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Offline rwooduk1

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Limiting reactant is not something that you should worry about when talking about dissociation (and no, in general it is not just the one with the lower number of moles, as it depends on the ratio in which things react).

What is the reaction that you are interested in? Just water electrolysis? Then the acid is there only to lower the resistance (it increases conductivity of the solution). In other situations it may play some additional roles.

As to dangers - produced hydrogen easily mixes with oxygen (either the atmospheric one or produced in the anode reaction) and the mixture is highly explosive (and easy to ignite - it won't explode on its own, but just a spark or flame and it goes kaboom). The less hydrogen produced, the safer you are.

Hello again Borek, this relates to a question I have asked here and previously on another forum. I am taking the development of the system step-by-step.

The end goal is to cause precipitation of calcium carbonate, from calcium carbonate solution, on a metal tube in the solution (via electrochemical generation of OH- ions at the surface).

I have spent some of the past week researching the effects of bubbling ##CO_{2}## into solution and, as per my previous posts, have come to the conclusion that this would significantly contribute. However the limited change in pH with ##CO_{2}## wont have a massive influence on the secondary effect of calcium carbonate dissolution.

Therefore, I decided to use a supporting acidic electrolyte solution to decrease the pH further than the ##CO_{2}## alone. Allowing more calcium carbonate to be dissolved before electrolysis and, like you say, increase the electrochemical process.

I am currently looking at electrodes that wont dissolve in acid, such as glassy carbon, as opposed to my original idea of copper foil.

Thank you for the warning on the hydrogen, I'll be careful no bunsen burners are running nearby in the lab. Also when i come to the acid will always add acid to water and not water to acid.

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Flames are easy, beware of electric sparks.

If you get pH too low you might have problems raising it back up to precipitate. My bet is if you use sulfuric acid as an electrolyte pH will be so low you won't get any deposits.

Depending on voltages used copper can be safe as an electrode. Other common candidates are graphite and stainless steel. Platinum is often best, but the cost is prohibitive.
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Offline rwooduk1

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Indeed. The hindrance of low pH to precipitation had also crossed my mind. I think this thing will require several experiments to determine optimum pH for the process. The online forums have been very helpful, there aren't many (if any) electrochemists where I am, and hopefully I can get the system to work. Now I just need to order equipment and start tests. Fingers crossed. Thank you.

Offline Enthalpy

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[...] The end goal is to cause precipitation of calcium carbonate [...]

Calcium carbonate in a solution of sulphuric acid? I doubt.

Offline rwooduk1

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[...] The end goal is to cause precipitation of calcium carbonate [...]

Calcium carbonate in a solution of sulphuric acid? I doubt.

pH too low? The electrochemistry will significantly increase the pH near the electrode, but will require more energy to neutralise the acid first. I have been told that a pH rise of upto pH9 can occur very rapidily in that region.

Offline rwooduk1

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[...] The end goal is to cause precipitation of calcium carbonate [...]

Calcium carbonate in a solution of sulphuric acid? I doubt.

Now I see what you mean, I can't get it to dissolve! With the H2SO4 I reduce the pH to ~1.4. But when I add the CaCO3 it just fizzes and the solution goes milky. I will try adding calcium carbonate to CO2 saturated water tomorrow and see if I have more sucess  ???

Edit i think the CO2 method will work if I keep bubbling it in while adding the calcium carbonate:

CO2 (g) + H2O (l) --> H2CO3 (aq)
CaCO3 (s) + H2CO3 (aq) --> Ca2+ (aq) + 2HCO3- (aq)
« Last Edit: December 02, 2019, 12:22:32 PM by rwooduk1 »

Offline AWK

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You got calcium sulfate poorly soluble in water.
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Offline rwooduk1

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You got calcium sulfate poorly soluble in water.

lol, with the acid yes I did :( the idea behind the acid was to use it as a supporting electrolyte, but now I realise that there will be calcium ions in there anyway, if I can get the calcium carbonate to dissolve. Tried the ##CO_{2}## alone today and again just got milky solution and bits of ##CaCO_{3}## at the bottom. The the reaction I posted above clearly didn't work as expected.

I am now going to try a different approach that I found in a paper (image below) by mixing calcium chloride and sodium bicarbonate, to make the calcium carbonate solution:

##CaCl_{2(aq)} + 2NaHCO_{3(aq)} \rightarrow CaCO_{3(s)} + CO_{2(g)} + H_{2}O_{(l)} + 2NaCl##

Then apply CO2 to increase the following reaction...

##CO_{2(g)}+ H_{2}O_{(l)} \rightarrow H_{2}CO_{3(aq)}##

To hopefully get...

##CaCO_{3(s)} + H_{2}CO_{3(aq)} \rightarrow Ca_{(aq)}^{2+} + 2HCO_{3(aq)}^{-}##

The paper also suggests adding 0.1 N NaOH to the ##NaHCO+{3}## to adjust the pH to 9.1. I'm assuming to aid in the precipitaion of the calcium carbonate(?). And also assuming normality for NaOH here is the same as molarity, i.e. N = M = 0.1 M because the valence of NaOH is 1.

Hopefully this time it will work.








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