Electrolysis of water

[MDJH]

I know I already started another electrochemistry-related thread, but this one is about a slightly separate subject. I was looking to use water electrolysis to yield hydrogen and oxygen in separate containers. For the most part, I think I'd know what I'm doing, but there's some things I'd like to check on.

 

First off, what would be a good (ie. safe and/or efficient) electrolyte to use for at-home electrolysis? I'm guessing that the electrolyte itself would undergo electrolysis as well... which is why I'm hesitant to use table salt, as that would supposedly yield chlorine gas. (Inside the same container as oxygen, I presume?) So far I'm considering baking soda, or vinegar, except obviously not at the same time... though I guess I could use sodium acetate, the product of the reaction between the two, as an electrolyte istead... in any case, what kind of side-products would baking soda or vinegar or sodium acetate yield?

[insane_alien]

sodium hydroxide (lye) is a good one. it does under go electrolysis but oxygen is released where it should be and the sodium immediately reacts with the water to form more sodium hydroxide and hydrogen.

 

you can use table salt, but you'd need to wait till it stopped producing chlorine (ie, when all that is left is sodiun hydroxide and the chlorine has escaped from the system).

 

it all depends on how pure you need it.

[MDJH]

sodium hydroxide (lye) is a good one. it does under go electrolysis but oxygen is released where it should be and the sodium immediately reacts with the water to form more sodium hydroxide and hydrogen.

 

you can use table salt, but you'd need to wait till it stopped producing chlorine (ie, when all that is left is sodiun hydroxide and the chlorine has escaped from the system).

 

it all depends on how pure you need it.

I'm a little scared to use sodium hydroxide though, as I've heard that can react with the amino acids in my skin. I suppose I could use gloves, but I don't know what kind of gloves would protect my hands from sodium hydroxide.

 

Similar concern with salt. The chlorine should be mostly contained, but if it ends up in the same container as the oxygen that means it would be dangerous to breathe it. (Granted, pure oxygen isn't something you should be breathing anyway, but supposedly doing so for only a few seconds isn't particularily dangerous and instead gives you an energetic feeling or something? Not that I'd necessarily just breathe it, I could also use it to light glowing splints or burn metals instead)

[insane_alien]

sodium hydroxide solution can cause chemical burns but only at high concentrations.

 

you only need a little bit for this and you can always keep the water topped up with distilled water to keep the concentration low.

 

if you get any of the solution on you rinse your hand under the cold tap. if you get some solid sodium hydroxide on you then vinegar will be a better choice to put on yor hand before rinsing. wear gloves and eye protection.

 

if usuing salt you'd want to do it outside and let the chlorine dissipate int othe atmosphere rather than trying to contain it.

[MDJH]

sodium hydroxide solution can cause chemical burns but only at high concentrations.

 

you only need a little bit for this and you can always keep the water topped up with distilled water to keep the concentration low.

 

if you get any of the solution on you rinse your hand under the cold tap. if you get some solid sodium hydroxide on you then vinegar will be a better choice to put on yor hand before rinsing. wear gloves and eye protection.

 

if usuing salt you'd want to do it outside and let the chlorine dissipate int othe atmosphere rather than trying to contain it.

How low, exactly? What concentrations would be considered safe, what would be dangerous? And what kind of gloves should I use? And if I used dilute table salt, how far away would I need to be from the electrolysis?

[Melvin]

I think another common one is very dilute H2SO4...

 

The one that I like is epsom salts. They're cheap, non-corrosive, and don't interfere with the reaction.

 

If you do want to use NaOH, I would think you could definitely use less than a 1M solution, probably you could get away with less than a .5M solution. I'm not sure exactly, but you don't need much. Latex gloves would work.

 

With table salt, you just need to keep it in a well ventilated area. You shouldn't be able to smell any chlorine unless you waft some of the anode gas; if you can smell it any other time, then you don't have enough ventilation. Also, electrolysis of table salt leads to the formation of hypochlorites (i.e. bleach).

[insane_alien]

Melvin, 0.01M would be enough. you don't need something that concentrated.

 

just depends on how quickly you need it to happen. if time isn't an issue then 0 electrolyte is needed.

[MDJH]

I think another common one is very dilute H2SO4...

 

The one that I like is epsom salts. They're cheap, non-corrosive, and don't interfere with the reaction.

 

If you do want to use NaOH, I would think you could definitely use less than a 1M solution, probably you could get away with less than a .5M solution. I'm not sure exactly, but you don't need much. Latex gloves would work.

 

With table salt, you just need to keep it in a well ventilated area. You shouldn't be able to smell any chlorine unless you waft some of the anode gas; if you can smell it any other time, then you don't have enough ventilation. Also, electrolysis of table salt leads to the formation of hypochlorites (i.e. bleach).

I have access to those too! Ironically, I happen to have a piece of wood soaking in epsom salt solution outside, I mentioned the same thing in my salt fire thread... it seems it takes a while to dissolve though.

 

On Wikipedia's page about epsom salt it says the solubility of the heptahydrate is "71 g/100 mL (20 °C)"; I assume that ratio applies regardless of the volume? (ie. 710g potentially dissolving in a litre at 20 °C, etc...)

 

In any case, to get it to dissolve to that ratio, would it spontaneously dissolve to that much if you left it for long enough, or would it be more effective to boil a lot of water, pour the boiling water into a bucket (obviously one that wouldn't melt or break by filling it with boiling water) add a little more than the solubility ratio for it, then leave it to cool and if it precipitates, assume saturation? Also would concentrated epsom salt at room temperature be safe for immersing your hands in? (I'm presuming it would as it's used for bath therapy but I'm not sure if concentration would affect that)

 

And what voltage would be good for electrolysis with concentrated (or close to concentrated) epsom salt as the electrolyte?

[Melvin]

Melvin, 0.01M would be enough. you don't need something that concentrated.

 

just depends on how quickly you need it to happen. if time isn't an issue then 0 electrolyte is needed.

 

Ok, I wasn't sure about that. Thanks for clarifying.

 

On Wikipedia's page about epsom salt it says the solubility of the heptahydrate is "71 g/100 mL (20 °C)"; I assume that ratio applies regardless of the volume? (ie. 710g potentially dissolving in a litre at 20 °C, etc...)...And what voltage would be good for electrolysis with concentrated (or close to concentrated) epsom salt as the electrolyte?

 

You don't need a saturated solution, just enough to make it conductive. Water decomposes at around 1.3v, but there is resistance in the system. 3v should be enough; the higher the amperage, the better.

[Melvin]

As for the bit about hydrogen and oxygen, I saw your post in my water electrolysis thread, and I'm guessing I could expect the hydrogen and oxygen yielded by water electrolysis, when epsom salt is the electrolyte, to be pure?

 

Should be pure...as long as you use an inert anode (best is Pt, graphite is okay)

[MDJH]

Should be pure...as long as you use an inert anode (best is Pt, graphite is okay)

Putting insulated wire in the solution shouldn't affect the purity though, should it? I'll probably use a graphite rod for the uninsulated electrodes. (Ie. the part that's supposed to bubble the gases; would the graphite left behind by a burned pencil suffice?)

[MDJH]

You don't need a saturated solution, just enough to make it conductive. Water decomposes at around 1.3v, but there is resistance in the system. 3v should be enough; the higher the amperage, the better.

What qualifies as "conductive" though? Would trying to get the solution close to saturation make it more efficient, or would it have negligible benefit?

[insane_alien]

negligible benfit.

[MDJH]

negligible benfit.

So, would half that concentrated require only one volt more or something?

[insane_alien]

it wouldn't require a different voltage at all. it will only affect the rate of electrolysis.

[MDJH]

it wouldn't require a different voltage at all. it will only affect the rate of electrolysis.

Ah ok... so would the electrolysis occur at half the rate, or more than half the rate, with half the concentration?

[insane_alien]

it'll be proportional to the current you have going through it the relationship with concentration tends to be non-linear iirc.

[MDJH]

it'll be proportional to the current you have going through it the relationship with concentration tends to be non-linear iirc.

I was under the impression that "proportional" implied "linear" unless it was square proportional, cubed proportional, etc. o.o

 

But anyway, would it presumably occur at more than half the rate than would be involved for saturated MgSO4, if it's half the saturation concentration?

[insane_alien]

proportionaly doesn't mean linear. just means for an increase in the former you get an increase in the latter. it makes no distinction on how this relationship works other than that.

 

i cannot remember which happens as i've not done any electrolysis in years. try it and see.

[Melvin]

A saturated solution would give the highest conductivity, but as soon as the amount of water drops (is converted to H2/O2) the MgSO4 will start precipitating.

 

Since the solubility of hydrated MgSO4 is around 71g/100mL, I would say use around 15g/100mL. This should give you plenty of conductivity without using too much.

[Melvin]

Putting insulated wire in the solution shouldn't affect the purity though, should it? I'll probably use a graphite rod for the uninsulated electrodes. (Ie. the part that's supposed to bubble the gases; would the graphite left behind by a burned pencil suffice?)

 

The issue is where the wire (presumably Cu) contacts the graphite. If any copper is touching the solution, it will dissolve away pretty quickly to give insoluble Cu(OH)2. The pencil graphite is mixed with clay so will usually fall apart, so I use electrodes from D-size zinc-carbon (aka "Heavy Duty") batteries. These are nice, big graphite electrodes great for electrolysis. Just make sure the battery does not have the words "alkaline" or "lithium" on it. There are videos on youtube about it so you can search around. (You should still be careful with it and wear gloves when working with those batteries.)

[MDJH]

The issue is where the wire (presumably Cu) contacts the graphite. If any copper is touching the solution, it will dissolve away pretty quickly to give insoluble Cu(OH)2. The pencil graphite is mixed with clay so will usually fall apart, so I use electrodes from D-size zinc-carbon (aka "Heavy Duty") batteries. These are nice, big graphite electrodes great for electrolysis. Just make sure the battery does not have the words "alkaline" or "lithium" on it. There are videos on youtube about it so you can search around. (You should still be careful with it and wear gloves when working with those batteries.)

I wouldn't want to waste a battery though. I'll probably look in some hardware store for graphite electrodes. As for the bit about copper reacting, I didn't know that. Now that I do, I'll probably use graphite covered with electrical tape for the "insulated component" of the electrodes.

 

EDIT: That said, I was talking to someone who works in the chemistry department of the university I go to, and she said MgSO4 wasn't a very good electrolyte for water electrolysis; she was recommending cobalt something chloride or whatever, but I mentioned that I'd be doing this inside and she recommended using some iodine compound instead. What was she referring to as not being good for an electrolyte?

[Melvin]

I wouldn't want to waste a battery though. I'll probably look in some hardware store for graphite electrodes. As for the bit about copper reacting, I didn't know that. Now that I do, I'll probably use graphite covered with electrical tape for the "insulated component" of the electrodes.

 

EDIT: That said, I was talking to someone who works in the chemistry department of the university I go to, and she said MgSO4 wasn't a very good electrolyte for water electrolysis; she was recommending cobalt something chloride or whatever, but I mentioned that I'd be doing this inside and she recommended using some iodine compound instead. What was she referring to as not being good for an electrolyte?

 

Just make sure you seal the connection well.

 

The "something" was probably II (as in cobalt(II) chloride). That would give off chlorine at the anode (just like table salt would) instead of oxygen. Not sure where the person was going with that one. As long as you use good inert anodes (which you are), MgSO4 shouldn't give you any problems.

 

The only time you would get different results is if you used a divided cell with MgSO4. In that case, you would get insoluble magnesium hydroxide AND hydrogen at the cathode, and sulfuric acid AND oxygen at the anode.

[MDJH]

Just make sure you seal the connection well.

 

The "something" was probably II (as in cobalt(II) chloride). That would give off chlorine at the anode (just like table salt would) instead of oxygen. Not sure where the person was going with that one. As long as you use good inert anodes (which you are), MgSO4 shouldn't give you any problems.

 

The only time you would get different results is if you used a divided cell with MgSO4. In that case, you would get insoluble magnesium hydroxide AND hydrogen at the cathode, and sulfuric acid AND oxygen at the anode.

Divided cell? What do you mean? My idea was to put water in a bucket, add some epsom salt and stir, then take two plastic bottles (ideally one of them being twice the volume of the other) and fill both with the epsom salt solution, then invert them (ie. where the cap would be is toward the bottom of it) onto the water, then connect them to my voltage supply with partly-insulated graphite electrodes... as in, everything but the segment that's actually in the bottle would be covered with electrical tape. This way the anode's gas products would be separate from the cathode's gas products... would that qualify as a "divided cell"?

 

Also, I'm intuitively guessing that what said person was referring to was probably the efficiency. It didn't occur to me to re-ask it today though.

[insane_alien]

no, a divided cell is where the electrodes would be in different vessels or buckets in your case with an ion bridge between them.